Adsorption of arsenic from aqueous solution on

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沸石咪唑骨架材料（ZIF-8）的结构生长过程

沸石咪唑骨架材料（ZIF-8）的结构生长过程陈英波;赵林飞;王彪;胡晓宇;刘冬青;周凤潇;张宇峰【摘要】Zeolitic imidazolate framework-8 (ZIF-8) was synthesized at room temperature by using 2-methylimidazole as an organic ligand, Zn(NO3)2 as metal ions source and methanol as solvent. The structure and formation process of ZIF-8 were characterized by using FTIR, XRD, TEM and TG. The results showed that ZIF-8 crystals are formed quickly at the initial stage, and they gradually grow up to form hexagon and dodecahedron in 3D. A fundamental understanding of ZIF-8 structural evolution could facilitate the preparation of functional metal-organic frame-work phases with controlled crystal size and morphology.%以2-甲基咪唑为有机配体，由硝酸锌提供金属离子，在室温下以甲醇为溶剂合成了沸石咪唑骨架（ZIF-8）晶体。

通过红外光谱、X射线衍射、透射电子显微镜、热重分析等手段对晶体结构及其生长过程进行了表征。

结果表明：随着反应时间延长ZIF-8晶体迅速形成，然后逐渐堆积成六面体二维结构，进而形成三维十二面体结构。

对ZIF-8结构的生长过程的理解有助于通过控制ZIF-8的大小和形态来制备出功能化的金属有机骨架材料。

玉米秸秆黄原酸盐的制备及其去除污染土壤淋洗废液中重金属的效果研究

玉米秸秆黄原酸盐的制备及其去除污染土壤淋洗废液中重金属的效果研究李武楠;廖晓勇;裴亮;王凌青;黄占斌【摘要】以玉米秸秆为原料制备黄原酸盐,用于吸附去除淋洗废液中的重金属,研究玉米秸秆黄原酸盐制备条件对去除污染土壤淋洗液中重金属离子效果的影响.结果表明,玉米秸秆(10 g)制备黄原酸盐的最优条件为:20％NaOH 200 mL,CS2加入量5 mL,黄化时间1 h,5％MgSO4加入量3 mL.在此条件下,玉米秸秆黄原酸盐可使废液中重金属Pb2+、Cd2+、Zn2+、Cr6+和Cu2+的离子去除率达83.3％～99.2％.处理重金属废液后的玉米秸秆黄原酸盐残渣浸出率低,污泥稳定.【期刊名称】《应用化工》【年(卷),期】2018(047)011【总页数】4页(P2385-2387,2393)【关键词】玉米秸秆黄原酸盐;重金属;铅;淋洗废液【作者】李武楠;廖晓勇;裴亮;王凌青;黄占斌【作者单位】中国矿业大学(北京) 化学与环境工程学院,北京 100083;中国科学院地理科学与资源研究所场地污染评估与修复中关村开放实验室,北京 100101;中国科学院地理科学与资源研究所场地污染评估与修复中关村开放实验室,北京100101;中国科学院地理科学与资源研究所场地污染评估与修复中关村开放实验室,北京 100101;中国科学院地理科学与资源研究所场地污染评估与修复中关村开放实验室,北京 100101;中国矿业大学(北京) 化学与环境工程学院,北京 100083【正文语种】中文【中图分类】TQ352;X53;X52;X71生物质黄原酸盐是一种优良的重金属吸附剂，具有吸附效率高、出水效果好、可回收利用、环保的特点。

生物质黄原酸盐可由多种生物质原材料制备，如椰枣木屑[1]、麦秆[2]、羟丙基纤维素[3]、甘蔗渣纤维素[4]、交联淀粉[5-6]等。

我国是一个农业大国，作物秸秆资源丰富。

2014年我国作物秸秆资源总量8.97亿t[7]，但利用率不高，大量秸秆被丢弃或就地焚烧[8-10]，为此我国国家发展改革委员会同农业部颁布了秸秆综合利用技术目录，推进秸秆综合利用。

吸附材料去除水体中砷的研究进展

㊀收稿日期:２０２２－０８－０８基金项目:国家自然科学基金(４１９７７２０５ꎬ５１８３２００７)ꎻ辽宁大学研究生优质在线课程建设与教学模式综合改革研究项目(ＹＪＧ２０２２０１０４５)ꎻ辽宁省科学技术计划项目(２０２１－ＭＳ－１５２)作者简介:孙丛婷(１９８３－)ꎬ女ꎬ辽宁辽阳人ꎬ博士ꎬ教授ꎬ研究方向:环境功能材料.㊀∗通讯作者:宋有涛ꎬＥ￣ｍａｉｌ:ｙｓｏｎｇ＠ｌｎｕ.ｅｄｕ.ｃｎ.㊀㊀辽宁大学学报㊀㊀㊀自然科学版第５０卷㊀第４期㊀２０２３年ＪＯＵＲＮＡＬＯＦＬＩＡＯＮＩＮＧＵＮＩＶＥＲＳＩＴＹＮａｔｕｒａｌＳｃｉｅｎｃｅｓＥｄｉｔｉｏｎＶｏｌ.５０㊀Ｎｏ.４㊀２０２３吸附材料去除水体中砷的研究进展孙丛婷ꎬ刘相正ꎬ刘㊀丹ꎬ宋有涛∗(辽宁大学环境学院ꎬ辽宁沈阳１１００３６)摘㊀要:近年来ꎬ砷污染由于其高毒性和致癌性成为了一个紧迫且具有挑战性的问题.无机砷离子在水环境中主要以Ａｓ(Ⅲ)和Ａｓ(Ⅴ)的形式存在ꎬ当其浓度超过允许标准时ꎬ将对人体健康产生严重威胁.由于矿业开采㊁金属冶炼和各类砷化物的广泛使用ꎬ全球多达２.２亿人正在遭受砷污染的威胁.因此ꎬ开发绿色㊁经济和有效的除砷技术对保护人类健康至关重要.现有的除砷技术包括沉淀法㊁离子交换法㊁膜分离法㊁生物法㊁电凝聚法和吸附法等.然而ꎬ上述大多数技术存在初始成本昂贵㊁维修成本高和二次污染等缺点.吸附法由于成本低㊁操作简单和材料来源范围广等优点ꎬ成为最常用的除砷技术之一.因此ꎬ本文综述了各种吸附剂去除砷污染的研究进展ꎬ以期为去除水中不同形态砷提供参考.关键词:砷ꎻ吸附ꎻ材料ꎻ除砷技术中图分类号:Ｘ５２㊀㊀㊀文献标志码:Ａ㊀㊀㊀文章编号:１０００－５８４６(２０２３)０４－０３２５－１５ＲｅｓｅａｒｃｈＰｒｏｇｒｅｓｓｏｆＡｒｓｅｎｉｃＲｅｍｏｖａｌｆｒｏｍＷａｔｅｒｂｙＡｄｓｏｒｂｅｎｔｓＳＵＮＣｏｎｇ￣ｔｉｎｇꎬＬＩＵＸｉａｎｇ￣ｚｈｅｎｇꎬＬＩＵＤａｎꎬＳＯＮＧＹｏｕ￣ｔａｏ∗(ＳｃｈｏｏｌｏｆＥｎｖｉｒｏｎｍｅｎｔꎬＬｉａｏｎｉｎｇＵｎｉｖｅｒｓｉｔｙꎬＳｈｅｎｙａｎｇ１１００３６ꎬＣｈｉｎａ)Ａｂｓｔｒａｃｔ:㊀Ｉｎｒｅｃｅｎｔｙｅａｒｓꎬａｒｓｅｎｉｃｐｏｌｌｕｔｉｏｎｈａｓｂｅｃｏｍｅａｎｕｒｇｅｎｔａｎｄｃｈａｌｌｅｎｇｉｎｇｐｒｏｂｌｅｍｄｕｅｔｏｉｔｓｈｉｇｈｔｏｘｉｃｉｔｙａｎｄｃａｒｃｉｎｏｇｅｎｉｃｉｔｙ.ＩｎｔｈｅｗａｔｅｒｅｎｖｉｒｏｎｍｅｎｔｉｎｏｒｇａｎｉｃａｒｓｅｎｉｃｉｏｎｓｍａｉｎｌｙｏｃｃｕｒｉｎｔｗｏｓｔａｔｅｓꎬＡｓ(Ⅲ)ａｎｄＡｓ(Ⅴ).Ｉｔｗｉｌｌｂｅｃｏｍｅａｓｅｒｉｏｕｓｔｈｒｅａｔｔｏｈｕｍａｎｈｅａｌｔｈｗｈｅｎｔｈｅｃｏｎｃｅｎｔｒａｔｉｏｎｅｘｃｅｅｄｓｔｈｅａｌｌｏｗａｂｌｅｌｉｍｉｔｓ.Ｈｕｍａｎｉｓａｔｒｉｓｋｏｆａｒｓｅｎｉｃｃｏｎｔａｍｉｎａｔｉｏｎｄｕｅｔｏｍｉｎｉｎｇꎬｍｅｔａｌｓｍｅｌｔｉｎｇꎬａｎｄｔｈｅｗｉｄｅｓｐｒｅａｄｕｓｅｏｆｖａｒｉｏｕｓａｒｓｅｎｉｃｃｏｍｐｏｕｎｄｓ.Ｅｘｉｓｔｉｎｇｔｅｃｈｎｏｌｏｇｉｅｓｉｎｃｌｕｄｅｐｒｅｃｉｐｉｔａｔｉｏｎｍｅｔｈｏｄꎬｉｏｎｅｘｃｈａｎｇｅｍｅｔｈｏｄꎬｍｅｍｂｒａｎｅｓｅｐａｒａｔｉｏｎｍｅｔｈｏｄꎬｂｉｏｌｏｇｉｃａｌｍｅｔｈｏｄꎬｅｌｅｃｔｒｏｃｏａｇｕｌａｔｉｏｎｍｅｔｈｏｄａｎｄａｄｓｏｒｐｔｉｏｎｍｅｔｈｏｄ.Ｈｏｗｅｖｅｒꎬｍｏｓｔｏｆｔｈｅａｂｏｖｅｔｅｃｈｎｏｌｏｇｉｅｓｈａｖｅｄｉｓａｄｖａｎｔａｇｅｓꎬｓｕｃｈａｓｈｉｇｈｉｎｉｔｉａｌｃｏｓｔꎬｈｉｇｈｍａｉｎｔｅｎａｎｃｅｃｏｓｔａｎｄｓｅｃｏｎｄａｒｙｐｏｌｌｕｔｉｏｎ.Ａｄｓｏｒｐｔｉｏｎｍｅｔｈｏｄｈａｓｂｅｃｏｍｅｏｎｅｏｆｔｈｅｍｏｓｔｃｏｍｍｏｎｌｙｕｓｅｄａｒｓｅｎｉｃｒｅｍｏｖａｌｔｅｃｈｎｉｑｕｅｓｄｕｅｔｏｉｔｓｌｏｗｃｏｓｔꎬｓｉｍｐｌｅｏｐｅｒａｔｉｏｎａｎｄｗｉｄｅｒａｎｇｅｏｆｍａｔｅｒｉａｌｓｏｕｒｃｅｓ.Ｔｈｅｒｅｆｏｒｅꎬｔｈｉｓｐａｐｅｒｒｅｖｉｅｗｓｔｈｅｒｅｓｅａｒｃｈｐｒｏｇｒｅｓｓｏｆｖａｒｉｏｕｓａｄｓｏｒｂｅｎｔｓｉｎｔｈｅｒｅｍｏｖａｌｏｆａｒｓｅｎｉｃｐｏｌｌｕｔｉｏｎꎬｉｎ㊀㊀ｏｒｄｅｒｔｏｐｒｏｖｉｄｅａｒｅｆｅｒｅｎｃｅｆｏｒｔｈｅｒｅｍｏｖａｌｏｆａｒｓｅｎｉｃｉｎｗａｔｅｒ.Ｋｅｙｗｏｒｄｓ:㊀ａｒｓｅｎｉｃꎻａｄｓｏｒｐｔｉｏｎꎻｍａｔｅｒｉａｌｓꎻａｒｓｅｎｉｃｒｅｍｏｖａｌｔｅｃｈｎｏｌｏｇｙ０㊀引言砷(Ａｓ)作为一种剧毒的元素ꎬ普遍存在于自然界之中[１].由于自然过程和人类活动导致其大量进入水体环境中ꎬ对人类健康产生严重威胁.伴随着全球工业化进程的飞速发展ꎬ水污染形势严峻ꎬ受到砷污染的水源在许多国家和地区被用作主要饮用水源[２－３].世界卫生组织(ＷＨＯ)建议饮用水中砷的最高允许质量浓度为１０μｇ/Ｌ[４].然而ꎬ在孟加拉国㊁印度㊁越南㊁中国和智利等许多发展中国家ꎬ数百万人正遭受着严重的砷中毒ꎬ其地下水中砷的质量浓度在１００~２０００μｇ/Ｌ之间ꎬ远超人体健康所允许的质量浓度ꎬ人体长期摄入被砷污染的水源会导致肝脏和肾脏衰竭ꎬ损害人体的免疫系统ꎬ并提高死亡风险(膀胱癌㊁肾癌㊁胃癌和皮肤癌等)[５－６].水体环境中砷的来源可以分为自然过程和人类活动ꎬ而造成水体环境中砷质量浓度超标的主要原因则是人类活动.在自然界中ꎬ砷主要以其他金属(铁㊁铜㊁锌等)或非金属(硫等)伴生矿物的形式存在.因此ꎬ矿物开采和金属冶炼过程中会产生大量含砷废水ꎬ并导致砷进入水体环境中[７－８].同时ꎬ砷化物在电子工业㊁陶瓷制造业㊁食品工业㊁农业和畜牧业生产中有着广泛应用ꎬ各类砷化物的使用也在一定程度上加重了砷污染[９－１０].图１　常见的除砷技术在水环境中ꎬ砷主要以无机阴离子砷酸根Ａｓ(Ⅲ)(ＡｓＯ３－４)和亚砷酸根Ａｓ(Ⅴ)(ＡｓＯ３－３)的形式存在[１１－１２].砷在水中的价态主要由ｐＨ和氧化还原条件决定ꎬ在还原或厌氧条件下(例如地下水中)ꎬ砷(Ⅲ)为主要的无机砷物种.相对于砷(Ⅴ)ꎬ砷(Ⅲ)的流动性更高ꎬ且毒性是砷(Ⅴ)的６０多倍[１３－１４].传统含砷废水的处理方法是先将废水中的砷(Ⅲ)预氧化为砷(Ⅴ)ꎬ然后再使用其他技术处理ꎬ这样的工艺过程繁琐且费用高.近年来ꎬ随着砷污染的不断加剧ꎬ越来越多的学者开始寻找治理砷污染的新方法ꎬ他们将含砷废水的氧化和去除相结合ꎬ简化了工艺流程并减少处理费用.含砷废水的处理方法可以大致分为３类:物理法㊁化学法㊁生物法.其中ꎬ主要包括沉淀法[１５－１６]㊁生物法[１７－１８]㊁离子交换法[１９－２０]㊁膜分离法[２１－２２]㊁电凝聚法[２３－２４]和吸附法[２５－２７](图１).上述大多数方法存在着初始成本和维修成本高㊁工艺复杂等缺点.近年来ꎬ吸附法成为最常用的除砷方法之一ꎬ尤其在发展中国家ꎬ因为其成本低㊁操作简单㊁原材料来源范围广和再生潜力大而被广泛应用.６２３㊀㊀㊀辽宁大学学报㊀㊀自然科学版２０２３年㊀㊀㊀㊀１㊀水体中砷的去除技术１.１㊀沉淀法沉淀法的原理是利用外加的化学试剂与废水中的砷形成难以溶解的砷化合物ꎬ然后将其过滤去除.砷酸根能够和许多金属离子形成难溶的沉淀ꎬ砷(Ⅴ)比砷(Ⅲ)更容易形成稳定的沉淀化合物ꎬ所以先向含砷废水中投加氧化剂ꎬ将砷(Ⅲ)氧化为砷(Ⅴ)再进行沉淀分离.常用的沉淀剂有钙盐㊁铁盐㊁铝盐和硫化物等.沉淀法具有成本低㊁操作简单㊁处理范围广等优点ꎬ但存在沉淀剂投加量大㊁废渣多等缺点ꎬ且沉淀物中含砷ꎬ如果得不到有效处理ꎬ容易造成二次污染.因此ꎬ该方法在许多地区的应用受到限制[２８－２９].１.２㊀生物法生物法的原理是利用植物或微生物的富集作用ꎬ将砷离子在生物体内浓缩和富集ꎬ或借助生物体本身的生命活动将砷离子转化为低毒性或无毒性的物质.砷(Ⅴ)离子的结构与磷酸盐相似ꎬ因此生物自身可以通过具有高亲和力的磷酸盐转运蛋白将砷(Ⅴ)在生物体内富集.此外ꎬ一些植物的根系分泌物可以和砷离子形成稳定的络合物ꎬ可降低砷在土壤中的移动性ꎬ起到固定和钝化作用.生物法具有高效㊁处理费用低廉和低浓度条件下处理效果好等优点ꎬ但生物体受外界环境影响较大ꎬ且不适合处理高浓度废水ꎬ对需要处理的水体要求较高ꎬ因此ꎬ生物法的实际应用也受到了一定限制[３０－３１].１.３㊀离子交换法离子交换法的原理是借助离子交换剂上的离子ꎬ选择性地去除水中砷离子.利用离子交换树脂的大比表面积和高反应活性ꎬ可以实现对砷离子的有效去除ꎬ而对砷离子的吸附能力主要取决于离子交换树脂中相邻电荷的空间距离和官能团的流动性.目前ꎬ纳米钼酸锆杂化离子交换树脂等多种离子交换树脂均可将砷(Ⅴ)的质量浓度从０.１ｍｇ/Ｌ降低至０.０１ｍｇ/Ｌ(符合ＷＨＯ标准).离子交换法具有能耗小㊁选择性强和出水水质好等优点ꎬ适用于小规模的贵金属工业废水ꎬ但因其维护成本高㊁设备价格贵㊁树脂易磨损等缺点限制了其应用[３２－３３].１.４㊀膜分离法膜分离法的原理是利用膜的选择透过性ꎬ以外界能量或化学位差为推动力ꎬ对双组分或多组分的溶质和溶剂进行分离㊁分级㊁提纯和浓缩.目前常用的技术有微滤(ＭＦ)㊁纳滤(ＮＦ)㊁超滤(ＵＦ)和反渗透(ＲＯ)等.膜分离技术可根据驱动压力大小分为高压驱动膜技术和低压驱动膜技术ꎬ高压驱动膜技术主要通过化学扩散除砷ꎬ而低压驱动膜技术主要通过物理筛分除砷.聚醚酰亚胺基超滤(ＵＦ)膜㊁聚偏氟乙烯基微滤(ＭＦ)膜等多种聚合膜的除砷效率均在９５％以上.膜分离法优点在于操作简单㊁无二次污染㊁污泥量低ꎬ但存在成本相对较高㊁能耗大等缺点ꎬ因此ꎬ膜分离法主要用于水净化[３４－３５].１.５㊀电凝聚法电凝聚法的原理是利用阳极板上电离出的金属(铁㊁铝等)离子与废水中的砷酸根发生絮凝反应ꎬ从而将砷酸根去除.电解生成的二价铁㊁三价铁㊁三价铝等金属离子经过一系列水解㊁聚合生成的羟基络合物可吸附砷离子形成共絮体ꎬ从而完成对砷的吸附和去除.与化学混凝相比ꎬ电凝聚法生成的污泥量少且絮团结构更加稳定.但电凝聚法需要专业的操作设备ꎬ运行成本和技术要求高ꎬ除砷效果受阳极材料和反应器设计的影响大.因此ꎬ在实际除砷应用中受到一定限制[３６－３７].７２３㊀第４期㊀㊀㊀㊀㊀㊀孙丛婷ꎬ等:吸附材料去除水体中砷的研究进展㊀㊀１.６㊀吸附法吸附法的原理是通过范德华力(物理吸附)或强共价键力(化学吸附)将待处理溶液中的离子吸附到吸附剂上ꎬ从而将待处理溶液中的目标物去除.根据作用力不同可将其分为物理吸附和化学吸附.物理吸附主要是依靠吸附材料与砷离子之间的范德华力ꎬ将砷离子吸附到材料表面ꎻ化学吸附是利用吸附材料上的离子与砷离子发生电子转移或交换ꎬ生成化学键ꎬ将砷离子固定到吸附材料上.常用的吸附材料有金属有机框架纳米材料㊁沸石㊁金属氧化物和生物炭等.吸附法具有操作简单㊁成本低㊁去除效率高等优点ꎬ同时吸附剂原材料来源广㊁种类多ꎬ是目前污水处理中最常用的方法之一[３８－３９].２㊀水体除砷的吸附剂类型２.１㊀沸石沸石是一种多孔硅酸盐的总称ꎬ其化学通式为[Ｍ(ⅠꎬⅡ)]ＯＡｌ２Ｏ３ｎＳｉＯ２ｍＨ２Ｏꎬ其中[Ｍ(ⅠꎬⅡ)]代表着一价和二价金属(通常为钾㊁钙㊁钠等).硅铝酸盐和这些一价㊁二价金属离子结合较弱ꎬ易与溶液中的其他金属离子发生交换作用ꎬ且交换后沸石的结构不会被破坏[４０].因此ꎬ沸石自身的性质和结构决定着其是一种优良的吸附剂ꎬ其主要吸附机理详见图２.但天然沸石自身杂质较多ꎬ对阴离子污染物的亲和力较低ꎬ纯度较低ꎬ其除砷效果有限.针对其缺点ꎬ对天然沸石进行改性处理ꎬ可以提高沸石的吸附性能和去除效率.常用的改性方式包括金属改性㊁表面活性剂改性等.图２㊀改性沸石除砷主要机理近年来ꎬ铁基材料由于具有丰富的孔隙结构㊁大比表面积和超顺磁性等优点成为国内外学者关注的重点.铁氧化物可通过静电吸引㊁离子交换㊁配位等多种吸附机理除砷ꎬ因此铁改性沸石除砷逐渐成为研究的热点.一些学者发现用铁改性可以提高天然沸石对砷的吸附能力ꎬ研究结果表明ꎬ在沸石改性的过程中ꎬ铁通过在沸石表面形成羟基氧化铁(ＦｅＯＯＨ)和铁氧化物提高对砷的吸附能力[４１－４２].但由于天然沸石的不同㊁铁负载量的不同和初始溶液ｐＨ不同等因素ꎬ铁改性沸石对砷的吸附能力也不同.因此ꎬ国内外学者对铁改性沸石开展了研究工作.Ｎｅｋｈｕｎｇｕｎｉ等[４３]利用硝酸铁对天然沸石进行改性ꎬ制得铁改性沸石(ＩＨＯＭＺ)ꎬ通过Ｘ射线衍射(ＸＲＤ)㊁扫描电子显微镜(ＳＥＭ)和能量色散Ｘ射线光谱(ＥＤＸ)等多种方法对改性沸石进行表征ꎬ发现天然斜发沸石的改性没有导致明显的结构变化ꎬ且铁成功负载到天然沸石表面.实验结果表明ꎬ砷(Ⅴ)的最佳吸附参数:初始砷(Ⅴ)质量浓度为１０ｍｇ/Ｌ㊁吸附剂用量为３.０ｇ㊁反应时间为９０ｍｉｎ.ＩＨＯＭＺ对砷(Ⅴ)的吸附量为１.６９ｍｇ/ｇ.溶液的初始ｐＨ对除砷效果无显著影响ꎬ但温度对ＩＨＯＭＺ的吸附能力有显著影响.ＩＨＯＭＺ对砷(Ⅴ)的吸附是通过内球配位的离子交换进行的ꎬ同时热力学结果(吸附能(ＥＤＲ)为１０.４３ｋＪ/ｍｏｌ)也证实了吸附过程为化学吸附过程.Ｌｉ等[４４]也利用三价铁离子对天然沸石进行改性得到铁改性沸石(Ｆｅ－ｅｚ).通过ＸＲＤ㊁ＳＥＭ㊁透射电子显微镜(ＴＥＭ)等８２３㊀㊀㊀辽宁大学学报㊀㊀自然科学版２０２３年㊀㊀㊀㊀表征方法发现ꎬ天然沸石在改性过程中晶体结构没有改变ꎬ且物理性质稳定.实验结果表明ꎬｐＨ对Ｆｅ－ｅｚ的吸附量有较大的影响ꎬ与Ｎｅｋｈｕｎｇｕｎｉ等[４３]研究结果相反ꎬ溶液的ｐＨ对砷的吸附有一定影响.在ｐＨ为３~６时ꎬＦｅ－ｅｚ对砷(Ⅴ)吸附量基本保持不变ꎬ随着ｐＨ升至１０ꎬ其吸附量则显著降低.Ｆｅ－ｅｚ对砷(Ⅲ)的吸附略有不同ꎬ在ｐＨ为６~９时ꎬＦｅ－ｅｚ对砷(Ⅲ)的吸附量最高.Ｆｅ－ｅｚ对砷的去除是通过在沸石表面形成的氧化铁与砷形成络合物.因此ꎬ相对于化学吸附ꎬ其吸附作用力较弱ꎬＦｅ－ｅｚ对砷(Ⅲ)和砷(Ⅴ)的吸附量为０.１ｍｇ/ｇ和０.０５ｍｇ/ｇ.相比于铁氧化物ꎬ纳米零价铁(ＮＺＶＩ)具有较高的阴离子吸附能力和独特的核壳结构[４５]ꎬ因此沸石与ＮＺＶＩ结合可能会提高其吸附能力.但ＮＺＶＩ易发生氧化聚集ꎬ致使表面积减少ꎬ降低吸附能力[４６].因此制备复合材料过程中要选取合适的材料减少ＮＺＶＩ的团聚和氧化.Ｌｉ等[４７]利用三价铁离子和天然沸石为原料ꎬ通过还原法合成了纳米零价铁改性沸石(Ｚ－ＮＺＶＩ).ＳＥＭ表征显示ꎬＮＺＶＩ在沸石表面均匀分散ꎬ同时傅里叶红外光谱(ＦＴＩＲ)显示ꎬＮＺＶＩ在Ｚ－ＮＺＶＩ表面没有被氧化.上述结果证明ꎬＮＺＶＩ成功地负载到天然沸石上.ＸＲＤ和Ｘ射线光电子能谱(ＸＰＳ)证实ꎬＺ－ＮＺＶＩ吸附砷(Ⅲ)后生成了砷酸铁(ＦｅＡｓＯ４).在ｐＨ为６时ꎬＺ－ＮＺＶＩ对砷(Ⅲ)的吸附量为１１.５２ｍｇ/ｇ.实验表明ꎬＺ－ＮＺＶＩ对砷吸附存在静电吸附㊁离子交换㊁氧化还原㊁共沉淀㊁络合等多种吸附机制.相对于铁改性ꎬ表面活性剂也广泛应用于天然沸石改性.壳聚糖因为成本低㊁生物降解性好㊁生物相容性好受到了广泛关注.壳聚糖是通过甲壳素的碱性脱乙酰反应得到的ꎬ其含有丰富的羟基和氨基ꎬ同时还具有一定的吸附能力ꎬ因此可以被用来除砷[４８].然而ꎬ壳聚糖在酸性介质中会发生一定的溶解ꎬ单一的壳聚糖对酸性废水中砷的去除效果有限[４９].因此ꎬ为了提高壳聚糖对砷的吸附ꎬ更多的研究集中在壳聚糖复合材料上.Ｈａｎ等[５０]制备了壳聚糖包覆的Ｎａ－Ｘ改性沸石.通过ＳＥＭ㊁ＸＲＤ㊁ＥＤＸ等表征发现ꎬ通过壳聚糖包覆的改性沸石没有出现壳聚糖的团聚现象.实验表明ꎬ壳聚糖包覆可以显著改善Ｎａ－Ｘ沸石的除砷性能.壳聚糖包覆的Ｎａ－Ｘ沸石去除砷(Ⅴ)最佳ｐＨ为２.１ʃ０.１ꎬ最大吸附量为６３.２３ｍｇ/ｇ.ＦＴＩＲ和ＸＰＳ分析结果表明ꎬ改性沸石对酸性废水中Ａｓ(Ⅴ)的去除是通过改性沸石表面官能团( ＮＨ２等)的吸附作用以及形成的ＡｓＮ键ＡｓＯ键的键合作用完成的.表１对上述改性沸石材料的除砷性能进行比较.表１㊀改性沸石除Ａｓ(Ⅲ)/Ａｓ(Ⅴ)性能比较材料目标最佳ｐＨ吸附等温线吸附量/(ｍｇｇ－１)参考文献ＩＨＯＭＺＡｓ(Ⅴ) Ｆ１.６９[４３]Ｆｅ－ｅｚＡｓ(Ⅲ)Ａｓ(Ⅴ)３９ＦＦ０.１０.０５[４４]Ｚ－ＮＺＶＩＡｓ(Ⅲ)９Ｆ１１.５２[４７]Ｎａ－ＸＡｓ(Ⅴ)２.１Ｌ６３.２３[５０]㊀㊀天然沸石作为一种晶体状的硅铝酸盐ꎬ由于其特殊的结构和离子交换能力而受到许多学者的关注.然而ꎬ天然沸石由于自身的一些缺陷导致其除砷能力有限ꎬ大量学者把目光集中在改性沸石上.在上述沸石除砷研究进展中ꎬ大量学者对天然沸石进行改性ꎬ制得的改性沸石相较于天然沸石除砷能力有较大提升.但由于沸石自身的结构和性质ꎬ改性后的沸石的吸附量相较于新型吸附材料(氧化石墨烯㊁金属有机框架)仍然存在着一定的差距.因此ꎬ近年来改性沸石逐渐淡出了学者们的视野.２.２㊀氧化石墨烯石墨烯(Ｇｒａｐｈｅｎｅ)是碳的一种同素异形体ꎬ也是构成其他石墨材料的基本单元.石墨烯具有许多优点ꎬ如大比表面积㊁高导电性和良好的机械弹性ꎬ但其结构稳定㊁片层间堆积力较大㊁分散性较差９２３㊀第４期㊀㊀㊀㊀㊀㊀孙丛婷ꎬ等:吸附材料去除水体中砷的研究进展㊀㊀这些特性限制了其应用和发展.因此ꎬ具有更好特性的氧化石墨烯(ＧｒａｐｈｅｎｅｏｘｉｄｅꎬＧＯ)获得更多关注.氧化石墨烯是石墨粉经过化学氧化得到的ꎬ其一般是指单层氧化石墨.与石墨烯相比ꎬ氧化石墨烯具有更丰富的含氧官能团ꎬ这些官能团可以嫁接其他化学物质ꎬ从而形成新的氧化石墨烯复合材料[５１－５２].氧化石墨烯表面的羟基㊁羧基等官能团的存在可以促进砷在氧化石墨烯表面的吸附ꎬ因此氧化石墨烯作为一种潜在的除砷吸附剂引起了学者们的兴趣ꎬ其主要吸附机理见图３.然而ꎬ氧化石墨烯的使用存在着一些问题ꎬ如处理后纳米颗粒较难分离㊁易团聚等[５３].因此ꎬ研究人员通常用不同的材料修饰氧化石墨烯以提高其处理性能.常用的修饰材料有金属纳米颗粒和天然聚合物等.图３　氧化石墨烯复合材料除砷主要机理铁纳米颗粒(ＦｅＮＰＳ)是一种常见的金属纳米颗粒.在改性沸石材料研究进展中ꎬ本文提到的ＮＺＶＩ是一种有效的铁基吸附剂ꎬＦｅＮＰＳ和ＮＺＶＩ具有相似的性质ꎬ其也具有较大的比表面积和良好的吸附能力.有研究表明ꎬ将ＦｅＮＰＳ装配在碳质材料表面可以提高其在溶液中的分散性ꎬ而石墨烯是目前最有应用前景的碳质材料之一[５４－５５].此外ꎬ氧化石墨烯中的官能团还可以充当纳米材料的成核位点ꎬ并促进更多的纳米颗粒分散在氧化石墨烯表面[５６].因此ꎬＦｅＮＰＳ－氧化石墨烯复合材料具有一定的除砷价值.Ｄａｓ等[５７]通过溶胶－凝胶法将ＦｅＮＰＳ装配在氧化石墨烯表面ꎬ合成了一种氧化石墨烯－铁复合纳米材料(ＧＦｅＮ).ＴＥＭ表征显示ꎬＦｅＮＰＳ均匀地分布在氧化石墨烯薄片上ꎬ且ＦｅＮＰＳ不仅均匀分布在氧化石墨烯表面ꎬ而且还分散到氧化石墨烯薄片之间.ＸＰＳ结果表明ꎬＧＦｅＮ表面的主要成分为ＦｅＯＯＨ和Ｆｅ２Ｏ３/Ｆｅ３Ｏ４.实验结果表明ꎬ将ＦｅＮＰＳ负载在氧化石墨烯表面ꎬ可以提高砷在纳米复合材料中的吸附率.ＧＦｅＮ对砷(Ⅲ)和砷(Ⅴ)的吸附量分别为３０６和４３１ｍｇ/ｇ.当砷(Ⅲ)和砷(Ⅴ)初始质量浓度为１００μｇ/Ｌ㊁ＧＦｅＮ投加量为２５０ｍｇ/Ｌꎬ砷的去除率可以在１０ｍｉｎ时达到９９％.共存离子实验中ꎬ杂离子和有机物对除砷效率无明显影响.Ｆｅ３Ｏ４也是一种常见的铁纳米颗粒ꎬ且Ｆｅ３Ｏ４本身带有磁性ꎬ可以解决纳米材料分离难的问题.Ｙｏｏｎ等[５８]研究了Ｆｅ３Ｏ４/氧化石墨烯复合材料(Ｍ－ＧＯ)和Ｆｅ３Ｏ４/还原氧化石墨烯复合材料(Ｍ－ｒＧＯ)对砷的去除效果.通过ＸＲＤ㊁ＦＴＩＲ和ＸＰＳ等多种方法对Ｍ－ＧＯ和Ｍ－ｒＧＯ进行表征发现ꎬ在Ｍ－ＧＯ和Ｍ－ｒＧＯ表面可以观察到Ｆｅ３Ｏ４颗粒.实验发现ꎬＭ－ＧＯ含有较多的含氧基团ꎬ有助于在氧化石墨烯表面形成Ｆｅ３Ｏ４ꎬ从而更好地吸附砷.所以ꎬＭ－ＧＯ对砷(Ⅲ)和砷(Ⅴ)都表现出更强的除砷能力.Ｍ－ＧＯ对砷(Ⅲ)和砷(Ⅴ)的吸附量分别为８５和３８ｍｇ/ｇꎬ而Ｍ－ｒＧＯ对砷(Ⅲ)０３３㊀㊀㊀辽宁大学学报㊀㊀自然科学版２０２３年㊀㊀㊀㊀和砷(Ⅴ)的吸附量分别为５７和１２ｍｇ/ｇ.许多稀土金属氧化物对重金属离子表现出较好的吸附能力ꎬ而二氧化铈(ＣｅＯ２)作为最常见的稀土材料之一ꎬ由于其具有优越的性能得到了广泛应用[５９].ＣｅＯ２在一定条件下具有在＋３价和＋４价之间相互转换的能力ꎬ还具有一定氧化能力ꎬ因此ＣｅＯ２在环境治理中有一定的发展潜力[６０].但ＣｅＯ２纳米颗粒尺寸较小ꎬ易团聚ꎬ一般需要以复合材料或嵌入支架的形式使用.氧化石墨烯具有的大比表面积以及稳定的空间架构可以较好地解决ＣｅＯ２纳米颗粒的问题ꎬ因此ＣｅＯ２/氧化石墨烯复合材料(ＣｅＯ２－ＧＯ)可能是一种良好的除砷吸附剂.Ｓａｋｔｈｉｖｅｌ等[６１]通过水热法合成氧化铈/氧化石墨烯(Ｃｅｒｉａ－ＧＯ)复合材料.ＳＥＭ表征显示ꎬＣｅｒｉａ－ＧＯ复合材料的形态结构与原始氧化石墨烯差异不大ꎬ表明合成过程没有改变氧化石墨烯的形态.ＴＥＭ分析结果表明ꎬＣｅＯ２纳米颗粒成功负载在氧化石墨烯表面.所制备的Ｃｅｒｉａ－ＧＯ复合材料在０.１ｍｇ/Ｌ的初始质量浓度范围内ꎬ砷(Ⅲ)和砷(Ⅴ)几乎被完全去除(去除效率超过９９.９９％).Ｃｅｒｉａ－ＧＯ对砷(Ⅲ)和砷(Ⅴ)的吸附量分别为１８５和２１２ｍｇ/ｇ.实验结果表明ꎬ砷通过静电作用吸附到Ｃｅｒｉａ－ＧＯ复合材料ꎬ三价铈离子是一个活性位点ꎬ可以吸附砷ꎬ由于溶液氧化还原环境的影响ꎬ四价铈离子向三价铈离子的转变是同时发生的ꎬ因此ꎬ三价铈离子质量浓度的增加进一步促进了溶液中砷的完全去除.相对于金属纳米颗粒ꎬ壳聚糖作为一种天然聚合物ꎬ其具有丰富的官能团和一定的吸附能力ꎬ但同时也具有一些缺点ꎬ因此需要制备壳聚糖复合材料以提升其性能ꎬ这一点已经在改性沸石除砷研究进展中提到过.Ｓｈｅｒｌａｌａ等[６２]制备了壳聚糖/磁性氧化石墨烯(ＣＭＧＯ)纳米复合材料.通过ＢＥＴ比表面积检测法(ＢＥＴ)㊁ＦＴＩＲ㊁ＳＥＭ㊁ＥＤＸ和振动样品磁强计(ＶＳＭ)对复合材料进行了表征.结果表明ꎬ所制备的ＣＭＧＯ纳米复合材料具有大比表面积(１５２.３８ｍ２/ｇ)和较好的饱和磁化强度(４９.３０ｅｍｕ/ｇ).当初始溶液ｐＨ从酸性到中性时ꎬＣＭＧＯ除砷效率不断提高ꎬ然而在碱性条件下ꎬ除砷效率有所降低.在ｐＨ为７.３时ꎬＣＭＧＯ吸附量最高ꎬ为４５ｍｇ/ｇ.此外ꎬ纳米复合材料的超顺磁特性可以利用外部磁场分离和回收纳米颗粒.表２对上述ＧＯ复合材料的除砷性能进行比较.表２㊀ＧＯ复合材料除砷(Ⅲ)/砷(Ⅴ)性能比较材料目标最佳ｐＨ吸附等温线吸附量/(ｍｇｇ－１)参考文献ＧＦｅＮ砷(Ⅲ)砷(Ⅴ)３５ＬＬ３０６４３１[５７]Ｍ－ＧＯＭ－ｒＧＯ砷(Ⅲ)砷(Ⅴ)砷(Ⅲ)砷(Ⅴ)７３７３ＦＦＦＦ８５３８５７１２[５８]Ｃｅｒｉａ－ＧＯ砷(Ⅲ)砷(Ⅴ)ＬＬ１８５２１２[６１]ＣＭＧＯ砷(Ⅲ)７Ｌ４５[６２]㊀㊀氧化石墨烯由于其自身的优异特性成为一种潜在的除砷吸附剂ꎬ但自身的团聚问题影响着其吸附性能.在上述氧化石墨烯复合材料除砷的研究中发现ꎬ氧化石墨烯复合材料很好地解决了自身团聚的问题ꎬ并提高了其除砷性能.相较于沸石㊁生物炭等传统吸附剂ꎬ氧化石墨烯复合材料的吸附能力有着较大的提升ꎬ成为近年来学者们关注的重点.２.３㊀生物炭生物炭(ＢｉｏｃｈａｒꎬＢＣ)是生物有机材料在缺氧或者无氧条件下高温碳化裂解得到的一种具有丰富微孔结构的固相物质.生物炭的基本组成元素为碳㊁氢㊁氧㊁氮等ꎬ富碳为其重要的特征.生物炭具１３３㊀第４期㊀㊀㊀㊀㊀㊀孙丛婷ꎬ等:吸附材料去除水体中砷的研究进展㊀㊀有丰富的孔隙结构和较大的比表面积ꎬ除此之外ꎬ其表面还含有大量官能团ꎬ如羟基( ＯＨ)㊁羧基( ＣＯＯＨ)㊁醌基( ＣＨＯ)等ꎬ以及一些含氮和含硫的官能团[６３－６４].独特的结构和大量的官能团赋予了生物炭良好的稳定性和强大的吸附能力ꎬ同时这些独特的理化性质使其有着巨大的应用前景.在原材料方面ꎬ其来源广泛ꎬ木屑㊁污泥㊁畜禽粪便㊁厨余垃圾和农田秸秆等经济易得废弃生物质都可以作为其原材料.制备成本低㊁操作简单㊁原材料来源广和吸附能力强等一系列优点使生物炭广泛应用于农业㊁工业㊁能源㊁环境等领域.但未改性生物炭作为重金属吸附剂受到处理后难以分离和再利用等条件的限制[６５]ꎬ且砷离子以氧阴离子的形式存在于水中ꎬ生物炭通常带负电荷ꎬ砷和未改性生物炭之间存在静电排斥ꎬ很难吸附到未改性生物炭上[６６].近年来ꎬ许多研究人员对生物炭进行改性处理以提高性能ꎬ常见的改性方式有金属改性㊁磷改性和氮改性等ꎬ改性生物炭除砷主要机理如图４所示.图４　改性生物炭除砷主要机理㊀㊀在改性沸石和氧化石墨烯复合材料除砷研究进展中ꎬ本文提到铁改性已经广泛应用于沸石和氧化石墨烯中.目前ꎬ铁改性生物炭因易于固液分离而受到越来越多的关注.因为活性炭在处理后难以分离和再利用ꎬ所以在污水处理中受到了限制.铁是一种常见的磁性颗粒ꎬ经铁改性后的生物炭具有磁性ꎬ因此吸附后可通过简单的磁性过程与介质分离[６７].此外ꎬ在生物炭上引入铁可以增加其表面积并形成更加丰富孔隙结构ꎬ从而为去除废水中的污染物提供更多的吸附位点[６８].然而ꎬ在生物炭上负载过多的铁可能会占据吸附位点或堵塞孔结构ꎬ导致吸附量的降低[６９].因此ꎬ铁改性生物炭需要进一步的研究.Ｙａｏ等[７０]利用氯化铁和硫酸铁对生物炭进行改性ꎬ制得氧化铁改性生物炭(ＦｅＯｘ－ＢＣ).ＸＲＤ表明ꎬ氧化铁改性生物炭中氧化铁保持立方尖晶石结构ꎬ这说明氧化铁的磁性基本不变ꎬ这使得氧化铁改性生物炭容易通过磁性过滤器分离.通过考察初始ｐＨ㊁反应时间㊁吸附剂用量和共存阴离子对砷(Ⅴ)的去除影响ꎬ结果表明ꎬ在初始砷质量浓度为１０ｍｇ/Ｌ㊁吸附剂用量为５.０ｇ/Ｌ㊁溶液初始ｐＨ为３.０~８.０㊁反应时间为１ｈ的条件下ꎬ氧化铁改性生物炭对砷(Ⅴ)的吸附量为２０.２４ｍｇ/ｇꎬ砷(Ⅴ)的去除率可超过９５％.在共存离子实验中ꎬ磷酸盐和硅酸盐对砷(Ⅴ)的去除影响较大ꎬ而硫酸盐对砷(Ⅴ)的去除几乎没有影响.Ｎｈａｍ等[７１]以水稻秸秆为原料ꎬ通过氯化铁对生物炭进行改性ꎬ制得铁改性水稻秸秆生物炭(Ｆｅ－ＢＣ).ＦＴＩＲ㊁ＳＥＭ和ＥＤＸ结果表明ꎬ铁成功负载在生物炭表面.与未改性生物炭相比ꎬ铁改性生物炭表现出较强的异质性ꎬ因此ꎬ铁改性生物炭比未改性生物炭的结构更有利于提高溶液中砷(Ⅴ)的吸附.ＦＴＩＲ结果表明ꎬ改性过程增强了生物炭表面官能团的强度ꎬ这些官能团可以成为生物炭吸附的新活性位点.铁改性生物炭在溶液初始ｐＨ为５.０时吸附量最大.根据Ｌａｎｇｍｕｉｒ等温线显示ꎬ铁改性生物炭的最大吸附量为２６.９ｍｇ/ｇ.近年来ꎬ锰氧化物作为吸附材料也广泛应用于砷的去除.锰氧化物其比表面积较大且表面活性能力较强ꎬ对重金属有络合和氧化作用ꎬ从而具有一定吸附固定重金属的能力.为了进一步提升生物炭对重金属的吸附性能ꎬ许多研究人员开始把目光集中在双金属改性材料上.一些研究发现ꎬ铁/锰二元改性吸附剂在吸附重金属离子方面比单金属改性材料更有优势[７２－７３]ꎬ因此铁/锰改性材料具备２３３㊀㊀㊀辽宁大学学报㊀㊀自然科学版２０２３年㊀㊀。

Adsorption of uranium(VI) from aqueous solution

Adsorption of uranium(VI)from aqueous solutionby diethylenetriamine-functionalized magnetic chitosanJinsheng Xu •Mansheng Chen •Chunhua Zhang •Zhengji YiReceived:16February 2013/Published online:12June 2013ÓAkade´miai Kiado ´,Budapest,Hungary 2013Abstract In this paper,the modiﬁed magnetic chitosan resin containing diethylenetriamine functional groups (DETA-MCS)was used for the adsorption of uranium ions from aqueous solutions.The inﬂuence of experimental conditions such as contact time,pH value and initial ura-nium(VI)concentration was studied.The Langmuir,Fre-undlich,Sips and Dubinin–Radushkevich equations were used to check the ﬁtting of adsorption data to the equilib-rium isotherm.The best ﬁt for U(VI)was obtained with the Sips model.Adsorption kinetics data were tested using pseudo-ﬁrst-order and pseudo-second-order models.Kinetic studies showed that the adsorption followed the pseudo-second-order kinetic model,indicating that the chemical adsorption was the rate-limiting step.The present results suggest that DETA-MCS is an adsorbent for the efﬁcient removal of uranium(VI)from aqueous solution.Keywords Uranium ÁAdsorption ÁDiethylenetriamine-functionalized magnetic chitosan (DETA-MCS)ÁIsothermIntroductionUranium is not only a main raw material for nuclear industry,but also a radioactivity element,which is one of metal ions such as cesium and strontium are highly toxic,causes progressive or irreversible renal injury and in acutecases may lead to kidney failure and death [1–5].For this reason,the recovery,accumulation,and removal of ura-nium are of great importance.Nowadays,recovery of uranium(VI)from dilute aqueous solution commonly includes coagulation,chromatographic extraction,chemi-cal precipitation,ion exchange,membrane dialysis,etc.[3,4],but they have several disadvantages,like clogging,high cost and ineffectiveness when uranium(VI)ions are present in the wastewater at low concentrations,especially in the range of 1–100mg/L [6,7].Therefore,the above methods have limitation in application.In recent years,much attention has been focused on various adsorbents with metal-binding capacities and low cost,such as chitosan,zeolites,clay or certain waste products [8].As well known,chitosan and its derivatives have great potential application in the areas of biotech-nology,biomedicine,food ingredients,and cosmetics.Furthermore,chitosans are the most important materials examined for the removal of toxic metal ions due to their inexpensive and effective in natures [9,10].Magnetic chitosan resins (MCR)have been widely used in various applications such as enzyme puriﬁcation,cell separation,and waste treatment [11,12].On the other hand,as we know,one of the promising methods is the use of chelating resins that have suitable functional groups capable of interaction with metal ions.And the number of chelate rings can increase the stability of complex formed by poly amine,furthermore,the diethylenetriamine can exhibit different kinds of coordination modes of interaction of the U(VI)ions by ﬁve-memberd chelating rings in order to increase the adsorption capacity.Based on the above dis-cussion,in this study,the diethylenetriamine-modiﬁed magnetic chitosan resins (DETA-MCS)were synthesized.It was expected that DETA-MCS should be efﬁcient for the removal of U(VI)ions,owing to the strong adsorptionJ.Xu (&)ÁM.Chen ÁC.Zhang ÁZ.YiDepartment of Chemistry and Material Sciences,KeyLaboratory of Functional Organometallic Materials of Hengyang Normal University,College of Hunan Province,Huangbai Road No.165,Hengyang 421008,Hunan,China e-mail:hynuxujs@J Radioanal Nucl Chem (2013)298:1375–1383DOI 10.1007/s10967-013-2571-2capacity for the target ions and quick separation from aqueous by the magnetism.The different factors affecting the uptake behavior such as pH,initial concentration of the U(VI)ions,and contact time were investigated.Moreover, the adsorption isotherms,kinetics and regeneration studies were also identiﬁed.Materials and methodsChitosan with40mesh,90%degree of deacetylation(DD) and molecular weight of1.39105,and diethylenetriamine were purchased from Shanghai Medicine Company.0.45l m polypropylene membrane(PPM)ﬁlter was pur-chased from Sinopharm Chemical Reagent Co.Ltd.(ori-ginal from Kenker Company,US).A stock solution of U(VI)(1,000mg/L)was prepared by dissolving U3O8in a mixture of HCl,H2O2and HNO3.The U3O8was provided by School of Nuclear Resources and Nuclear Fuel Engi-neering,University of South China.All working solutions of different U(VI)concentrations were obtained by diluting the stock solution with distilled and deionized water at room temperature.All other reagents and solvents used in this study were of analytical grade.Uranium adsorption experimentsBatch sorption experiments of U(VI)were conducted in a series of250mL conicalﬂasks.Generally speaking, 100mL U(VI)solution was mixed with a known amount of DETA-MCS powder.The pH of the U(VI)solution was adjusted as desired using1.0mol/L NaOH and1.0mol/L HCl before mixing with the adsorbent(20mg DETA-MCS powder).A sample of solution was collected at suitable time intervals andﬁltered through a0.45l m PPMﬁlter which does not adsorb uranyl cations.Then theﬁltrates were analyzed for U(VI)concentration in the supernatants using a standard method given by Xie et al.[13].The U(VI)removal efﬁciency and adsorption capacity of U(VI) onto the DETA-MCS were calculated using the following equations:removal efficiencyð%Þ¼C0ÀC fC0Â100ð1Þq e¼ðC0ÀC fÞVWð2Þwhere q e denotes the adsorption capacity of U(VI)onto the DETA-MCS(mg/g);C0and C f the concentrations of the U(VI)in the solution before and after adsorption(mg/L), respectively;V the volume of the aqueous solution(L);and W is the mass of dry adsorbent used(g).Results and discussionPreparation and characterizationThe magnetic chitosan microspheres(MCS)were prepared according to the literature[14].The MCS(10.0g)were suspended in120mL isopropyl alcohol to which10mL epichlorohydrine(125mmol)dissolved in200mL ace-tone/water mixture(1:1v/v)was added.The solid which is modiﬁed MCS with epichlorohydrine(MCS-ECH)was ﬁltered and washed by ethanol followed by water for three times.Then the MCS-ECH obtained were suspended in 200mL ethanol/water mixture(1:1v/v),then diethylene-triamine(10mL)was added.The reaction mixture was stirred at60°C for12h,and the solid products DETA-MCS wereﬁltered and successively washed with acetone, demineralized water,and methanol,and dried in a vacuum oven at60°C.The resins studied were synthesized as shown in Fig.1.Figure2shows the FTIR spectra of DETA-MCS and MCS,The peaks at560–660cm-1were assigned to Fe–O bond vibration of Fe3O4.The absorption band around 3,380cm-1,revealing the stretching vibration of N–H group and–OH group in magnetic chitosan,and at 1,588cm-1conﬁrms the N–H scissoring from the primary amine,due to the free amino groups in the cross-linked chitosan.But in DETA-MCS,the peak becomes broad because of the existence of–NH2.The increasing intensity at1,667and1,082cm-1in the spectrum of DETA-MCS indicates that DETA-MCS has more amine groups than the unmodiﬁed magnetic chitosan(MCS).Effect of contact timeSince the contact time between the adsorbate and adsorbent is a key parameter for the adsorption process,the contact time required for the sorption equilibrium experiments was ﬁrst determined.Under the conditions of50mL solution contain20mg adsorbent amount,pH 3.5,298K and 50mg/L U(VI),the adsorption experiments were carried out for contact times ranging from20to180min.The results are shown in Fig.3.The sorption capacity increased with increasing contact time and a larger amount of ura-nium was removed by DETA-MCS in theﬁrst60min of contact time.Then the U(VI)sorption process proceeded slowly and reached saturation levels gradually at about 120min.After120min,the change of adsorption capaci-ties for U(VI)did not show notable effects.In our study,a contact time of120min was selected to guarantee an optimum U(VI)uptake.Effect of initial pH valuesIt is known that the medium pH has an inﬂuence upon the uranium sorption process because it controls the solubility of metals as well as the dissociation state of some func-tional groups,such as carboxyl,hydroxyl and amino on the adsorbent surface [15–17].In order to search for the opti-mum pH for the adsorption process as well as to ﬁnd out whether the DETA-MCS was able to show a good U(VI)uptake at extreme pH values,metal uptake was studied at pH ranging from 1.0to 6.0.The dependence of adsorption percentage of U(VI)ions on the pH of solution was given in Fig.4.The adsorption percentage of the U(VI)ions adsorbed on the DETA-MCS indicated a marked inﬂuencewith increasing pH of solution from 1.0to 3.5then started to decrease slightly with further increase in the pH of solution after reaching a maximum of 96.1%at pH 3.5.In strong acidic solutions (pH \3.5),more protons will be available to protonate amine groups to form groups –NH 3?,reducing the number of binding sites for the adsorption of UO 22?,therefore,the removal efﬁciency of uranium is lower in strong acidic solutions (pH \3.5).However,the availability of free U(VI)ions is maximum at pH 3.5and hence maximum adsorption,when pH value increase beyond 3.5,hydrolysis precipitation starts because of the formation of complexes in aqueous solution [18].The hydrolysis of uranyl ions play signiﬁcant role in deter-mining the equilibrium between U(VI)in solution and on adsorbent.Hydrolysis products occur,includingUO 2(OH)?,(UO 2)2(OH)22?,(UO 2)3(OH)53?,which results in decline of adsorption removal efﬁciency of U(VI),similar results were also observed [19].The hydrolysis equilibria are as follows:UO 2þ2þ2H 2O UO 2ðOH ÞþþH 3O þpK 1¼5:82UO 2þ2þ4H 2O ðUO 2Þ2ðOH Þ2þ2þ2H 3O þpK 2¼5:623UO 2þ2þ10H 2O ðUO 2Þ3ðOH Þþ5þ5H 3OþpK 3¼15:63where pKs are the logarithms of the equilibrium constants.When the pH becomes low enough,the divalent free UO 22?becomes the dominant ion form in the solution.Along with increasing pH,the percentage of UO 22?in the solution decreases,whereas the percentage oftheFig.1Scheme for the synthesis ofDETA-MCSFig.2FT-IR spectra MCS andDETA-MCSFig.3Effect of contact time on the adsorption of uranium (VI)([UO 22?]=50mg/L,DETA-MCS =20mg,pH =3.5,and T =298K)Fig.4Effect of initial pH on the adsorption of uranium (VI)([UO 22?]=50mg/L,DETA-MCS =20mg,and T =298K)monovalent hydrolyzed species,UO 2(OH)?,(UO 2)3(OH)5?,increases.At higher pH [5.5,dissolved solid schoepite (4UO 3Á9H 2O)exist in the solution.In view of the above result,all subsequent experiments were performed at pH 3.5.Effect of initial uranium(VI)concentrationThe percentage removal and adsorption capacity of U(VI)by contacting a ﬁxed mass of DETA-MCS (20mg)at the temperature (298K)and initial pH (3.5)using a range of initial U(VI)concentrations were shown in Fig.5.It was found that the adsorption removal efﬁciency of U(VI)decreased with increasing the initial U(VI)concentration in the aqueous solution.On one hand,this is because more mass of uranium is put into the system with increasing the initial U(VI)concentration in the aqueous solution.On the other hand,because of the higher mobility of uranyl ions (UO 2)2?in the diluted solutions,the interaction of this ion with the adsorbent also increases slowly.All in all,the adsorption capacity of DETA-MCS for uranium increased with increase in the initial uranium concentration.Similar results on the inﬂuence of the U(VI)biosorption has beenreported by Ku¨tahyal ıet al.[20]in their study using acti-vated carbon prepared from charcoal by chemical activation.Adsorption isothermThe adsorption isotherm is the most important information,which indicates how the adsorbent molecules distribute between the liquid and the solid phase when the adsorption process reaches an equilibrium state [21].The parameters of Langmuir,Freundlich,Sips and Dubinin–Radushkevich(D–R)models obtained are given in Table 1.In the research,the sorption data have been subjected to different sorption isotherms,namely the Langmuir,Freundlich,Sips and D–R isotherm models.Figure 6shows the adsorption isotherm of uranium(VI)on the DETA-modiﬁed MCR from the non-linear models.The Langmuir equation assumes that:(i)the solid surface presents a ﬁnite number of identical sites which are energetically uniform;(ii)there is no interactions between adsorbed species,meaning that the amount adsorbed has no inﬂuence on the rate of adsorption and (iii)a monolayer is formed when the solid surface reaches saturation.The Langmuir isotherm con-siders the adsorbent surface as homogeneous with identical sites in terms of energy.Equation (3)represents the Langmuir isotherm:q e ¼q m K L C e 1þK L C eð3Þwhere C e is the concentration of the adsorbate in solution at equilibrium (mg/g),q e is the adsorption capacity at equilibrium (mg/g),q m is the maximum adsorption capacity of the adsorbent (mg/g),and K L is the Langmuir adsorption constant related to the energy of adsorption (L/mg).The empirical Freundlich equation based on adsorption on a heterogeneous surface is given as follows:q e ¼K F C 1=neð4Þwhere q e denotes the equilibrium adsorption capacity (mg/g);C e the residual U(VI)concentration in the solution at equilibrium (mg/g);K F the Freundlich constant related to the adsorption capacity of sorbent (mg/g);n the Freundlich exponent related to adsorption intensity.To resolve the problem of continuing increase in the adsorbed amount with a rising concentration as observed for Freundlich model (Fig.6),an expression was proposed as Sips isotherm model [22,23],which is a combined form of Langmuir and Freundlich expressions deduced for predict-ing the heterogeneous adsorption systems.It is given as:q e ¼q s K s C 1=me1þK s C 1=með5Þwhere q s (mg/g)is the Sips maximum uptake of U(VI)per unit mass of DETA-MCS,K S (L/mg)is Sips constant related to energy of adsorption,and parameter m could be regarded as the Sips parameter characterizing the system heterogeneity.Figure 6shows the equilibrium adsorption of U(VI)ions onto the DETA-MCS and the ﬁtting plot of the three iso-therm models.For the studied system,the Sips isotherm correlates best (R =0.998)with the experimental data from adsorption equilibrium of U(VI)ions by DETA-MCS in these models.The phenomenon also suggeststheFig.5Effect of initial concentration on the adsorption of ura-nium(VI)(DETA-MCS =20mg,pH =3.5,and T =298K)heterogeneity of the adsorption,which may be attributed to the complicated form of U(VI)ions at the acid pH regions and the heterogeneous distribution of the active sites on DETA-MCS surface.The maximum adsorption capacity of DETA-MCS for U(VI)ions obtained by Sips isotherm model is 69.68mg/g.Dubinin–Radushkevich isotherm is more general than the Langmuir isotherm because it does not assume a homogeneous surface or constant sorption potential [24].Therefore,in this paper,the D–R isotherm is also used to analyze the experimental isotherm data.The linearized form of the D–R isotherm may be written as:ln C ads ¼ln q m ÀKE 2ð6Þwhere C ads is the amount of metal ions adsorbed on per unit weight of adsorbent,q m is the maximum sorption capacity and K is the activity coefﬁcient related to the mean adsorption energy and E is the Polanyi potential which is equal to:E ¼RT ln ð1þ1=C e Þð7ÞThe values of q m and K deduced by plotting ln C ads versus E 2(Fig.7),and the mean energy of adsorption (E )was calculated from the equation,according to the D–R isotherm,as:E ¼1=ðÀ2K Þ1=2ð8ÞThe plot of ln C ads versus E 2as shown in Fig.7is a straight line.From the slope and intercept of this plot the values of K =-5.9983910-9mol 2/kJ 2and q m =70.52mg U(VI)/g have been estimated.As we know,the adsorption value of the mean sorption energy is in the range of 1–8kJ/mol and in that of 9–16kJ/mol predicted the physical adsorption and the chemical adsorption,respectively [25].The value of E is calculated to be E =9.13kJ/mol and evaluated in the range of 9–16kJ/mol for composite adsorbent.The value of E is expected for chemical adsorption.It is assumed to be heterogeneous in the structure of composite.The results of linearized equations are shown in Table 1,the Langmuir model effectively described the sorption data with all R values [0.99.The adsorption isotherms of U(VI)ions exhibit Langmuir behavior which indicates a mono-layer paring the four isotherm models described above,Sips isotherm is most suitable to char-acterize the uranium-sorption behavior of DETA-MCS according to the values of R .Adsorption kineticsIn order to investigate the kinetic mechanism,which con-trols the adsorption process,the pseudo-ﬁrst-order andTable 1Isotherm constants and values of R for DETA-MCS ParameterValue R Langmuir isotherm q m (mg/g)65.160.991K L (L/mg)1.24Freundlich isotherm K F (mg/g)36.350.898n3.36Sips isotherm q s (mg/g)69.680.998Ks (L/mg) 1.27m2.74Dubinin–Radushkevich (D–R)isotherm K (mol 2/kJ 2)-5.998391090.955q m (mg/g)70.52Eads (kJ/mol)9.13Fig.6Plots of q e versus C e for the adsorption of uranium(VI)on DETA-MCS (DETA-MCS =20mg,pH =3.5,and T =298K)Fig.7Dubinin–Radushkevich isotherm of sorption U(VI)onto DETA-MCS adsorbentpseudo-second-order models were used to test the experi-ment data [26,27].The pseudo-ﬁrst-order kinetic model is given as:ln ðq e Àq t Þ¼ln q e Àk 1tð9Þwhere q e is the amounts of adsorbed metal per unit mass (mg/g)at equilibrium and k 1is the rate constant of pseudo ﬁrst-order sorption (min -1).The value of the rate constant k 1and q e for the pseudo-ﬁrst-order sorption reaction can be obtained by plotting ln (q e -q t )versus t as well as further linear regression analysis (Fig.8).A series of parameters,including kinetic constants,correlation coefﬁcients and q e values,were obtained via linear regression analysis and shown in Table 2.The calculated q e value of ﬁrst-order kinetic model (54.35mg/L)cannot give reasonable values,which was lower greater than experimental value Q exp (62.75mg/L).Hence,this equation cannot provide an accurate ﬁt of the experimental data.The pseudo-second order kinetic model deﬁnes that the rate controlling mechanism formed by chemical reaction for the sorption of metal ions on adsorbents [28–31].In order to describe U(VI)sorption on the DETA-MCS resin for the initial U(VI)concentrations at constant temperature (298K),the kinetic data obtained from batch adsorption experiments have been analyzed using the pseudo-second order kinetic equation given below.t q t ¼t q e þ1k 2q eð10Þwhere k 2is the rate constant of pseudo-second-order adsorption (g/mg min).The pseudo-second-order plot (Fig.9)is also linear with correlation coefﬁcient of 0.994(Table 2),however the calculated value of adsorption capacity,q e ,cal (61.33mg/g)is close to the value of experimental adsorption capacity,q e,exp (62.75mg/g).Therefore,the pseudo-second-order rate kinetic model best described the experimental data.In other words,U(VI)sorption by DETA-MCS followed the pseudo-second-order kinetic reaction.The goodness of the ﬁt to the pseudo-second-order kinetic model indicates that U(VI)adsorption on the DETA-MCS resin occurred by chemical adsorption [32].Comparison of U(VI)sorption capacity with other adsorbentsTo evaluate the potential application prospect of DETA-MCS,the prepared magnetic adsorbent can be well dis-persed in the water and can be easily separated magneti-cally from the medium after adsorption.These unique features present this adsorbent as a novel,promising and feasible alternative for uranium removal as compared with other adsorbents [33–45](Table 3).This paper emphasizes the material of DETA-MCS as a novel adsorbent in envi-ronmental remediation.Although a direct comparison of DETA-MCS with other adsorbents is very difﬁcultowingFig.8Pseudo-ﬁrst-order kinetics of uranium(VI)adsorption on DETA-MCS ([UO 22?]=50mg/L,DETA-MCS =20mg,pH =3.5,and T =298K)Table 2Kinetic parameters of uranium(VI)adsorbed onto DETA-MCSPseudo-ﬁrst-order Pseudo-second-orderExperimental Q e valuek 1(min -1)Q e 1(mg/g)R k 2(g/(mg min))Q e 2(mg/g)R Q exp (mg/g)0.0326554.350.9490.004461.330.99462.75Fig.9Pseudo-second-order kinetics of uranium(VI)adsorption on DETA-MCS ([UO 22?]=50mg/L,DETA-MCS =20mg,pH =3.5,and T =298K)to different experimental conditions adopted,it is con-cluded that U(VI)sorption capacity of DETA-MCS is higher than that of chitosan grafted MWCNTs,magnetic Fe3O4@SiO2and Fe3O4/GO,but lower than that of ion-imprinted MCR.The lower uptake values observed may be attributed to the increased extent of protonation of amino groups in the acidic solution.All in all,it is noteworthy that DETA-MCS-bound U(VI)can be favorably and quickly separated from a solution by using the external magnetic ﬁeld and it is a prospective adsorbent for application in the ﬁeld of U(VI)removal.Resins regenerationTo make the process more effective and economically feasible,sorbent regeneration and U(VI)recovery must be evaluated.Regeneration of the DETA-MCS resins was carried out using0.5M HNO3.Adsorption/desorption cycles were carried out repeatedly.The regeneration efﬁ-ciency was calculated using equation[46,47].Regeneration efficiencyð%)¼Uptake of UðVIÞin the second cycleUptake of UðVIÞin the first cycleÂ100ð11ÞRegeneration efﬁciency of94.4%was achieved for DETA-MCS over three cycles with a standard deviation of ±1.1%.On the other hand,the DETA-MCS after being used for three cycles could still be aggregated very fast from the solution by a3,000G magnet.This higher regeneration efﬁciency along with easy separation from adsorption medium using external magneticﬁeld indicate promising application in theﬁeld of U(VI)removal.ConclusionsThe U(VI)adsorption capacity by DETA-MCS was strongly dependent on contact time,pH,and initial ura-nium(VI)concentration.The adsorption capacity of U(VI) onto DETA-MCS increases with an increase of contact time and reaches adsorption equilibrium within120min, the adsorption removal efﬁciency 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砷和氟污染地下水

ReviewArsenic andﬂuoride contaminated groundwaters:A review of current technologies for contaminants removalSachin V.Jadhav a,Eugenio Bringas b,Ganapati D.Yadav a,*,Virendra K.Rathod a, Inmaculada Ortiz b,Kumudini V.Marathe aa Department of Chemical Engineering,Institute of Chemical Technology,Nathalal Parekh Marg,Matunga,Mumbai,400019,Indiab Department of Chemical and Biomolecular Engineering,Universidad de Cantabria,Avda,Los Castros s/n.39005,Santander,Spaina r t i c l e i n f oArticle history:Received23November2014 Received in revised form26June2015Accepted7July2015Available online8August2015Keywords:FluorideArsenicMembraneAdsorptionElectrocoagulationIon exchange,Contaminant removal a b s t r a c tChronic contamination of groundwaters by both arsenic(As)andﬂuoride(F)is frequently observed around the world,which has severely affected millions of people.Fluoride and As are introduced into groundwaters by several sources such as water e rock interactions,anthropogenic activities,and groundwater recharge.Coexistence of these pollutants can have adverse effects due to synergistic and/or antagonistic mechanisms leading to uncertain and complicated health effects,including cancer.Many developing countries are beset with the problem of F and As laden waters,with no affordable tech-nologies to provide clean water supply.The technologies available for the simultaneous removal are akin to chemical treatment,adsorption and membrane processes.However,the presence of competing ions such as phosphate,silicate,nitrate,chloride,carbonate,and sulfate affect the removal efﬁciency.Highly efﬁcient,low-cost and sustainable technology which could be used by rural populations is of utmost importance for simultaneous removal of both pollutants.This can be realized by using readily available low cost materials coupled with proper disposal units.Synthesis of inexpensive and highly selective nanoadsorbents or nanofunctionalized membranes is required along with encapsulation units to isolate the toxicant loaded materials to avoid their re-entry in aquifers.A vast number of reviews have been published periodically on removal of As or F alone.However,there is a dearth of literature on the simultaneous removal of both.This review critically analyzes this important issue and considers stra-tegies for their removal and safe disposal.©2015Elsevier Ltd.All rights reserved.1.IntroductionWater is not only an essential component for life but also a basic building block to maintain quality of life.Water scarcity has already revealed adverse effects on all populations in every continent.More recently,UNICEF and WHO reports have conﬁrmed that748million people do not have adequate and safe water resource and over2.5 billion people have access to meagre water supply.The WHO also estimates that1.8billion people use faecally contaminated source of drinking water(UNICEF/WHO,2014).Groundwater is used for potable purposes by over50%of the global population.Thus, groundwater is sometimes described as the‘hidden sea’.This is indeed true to a greater extent in countries like India where local supply to~80%rural and~50%urban dwellings is provided by groundwater sources alone(Ayoob et al.,2008).Presence of several naturally occurring,anthropogenic and in-dustry generated ions such asﬂuoride,arsenic,nitrate,sulfate,iron, manganese,chloride,selenium,heavy metals,and radioactive materials may greatly compromise water quality,leading to health problems.The most signiﬁcant inorganic pollutants in groundwater affecting human health at the global scale,according to the WHO, are arsenic andﬂuoride(Thompson et al.,2007).In this context,ﬂuoride pollution of drinking water receives much less consider-ation than arsenic.1.1.Health effects of single and combined As and FFluoride is the only chemical in potable water that can causeAbbreviations:WHO,World Health Organization;USEPA,United States Envi-ronmental Protection Act;CPC,chemical precipitation/coagulation;EC,electro-coagulation;EF,electrocoagulation/ﬂotation;AD,adsorption;IE,ion exchange;MT, membrane technology;RO,reverse osmosis;FO,forward osmosis;NF,nano-ﬁltration;ED,electrodialysis.*Corresponding author.E-mail addresses:gdyadav@,gd.yadav@.in (G.D.Yadav).Contents lists available at ScienceDirect Journal of Environmental Management journal h omepage:w /locate/jenvman/10.1016/j.jenvman.2015.07.0200301-4797/©2015Elsevier Ltd.All rights reserved.Journal of Environmental Management162(2015)306e325different health effects depending upon its concentration in dis-solved form.A very small amount ofﬂuoride is beneﬁcial for bone and teeth development and dental health.However,concentrations higher than1.5mg/L are damaging to human health,causing dental or skeletalﬂuorosis(Miretzky and Cirelli,2011).Children below12 years are likely to be most exposed toﬂuorosis as their body tissues continue to grow during the formative age.Moreover,ﬂuorosis is non-reversible and the disorder has no medical treatment.The WHO permits aﬂuoride concentration of0.5e1mg/L in drinking water(WHO,2011).Efﬂuent limit of4mg/L for F from the waste-water treatment facilities has been set by USEPA(Shen et al.,2003). Fig.1depicts the statistics on population exposed to F contamina-tion.Clearly,China and India are the most affected countries where nearly35and26million,respectively,people are atﬂuoride risk.As for arsenic,As(V)(arsenate)and As(III)(arsenite)are the most predominant valence states,which are found in aerobic sur-face waters and anaerobic groundwaters,respectively.Between pH 4and10,major As(III)compound is charge neutral,whereas As(V) species exists as charge negative.The occurrence of As in ground-water poses even a greater danger than F hazards due to its extreme toxicity at low concentration which goes undetected especially in As(III)form(Camacho et al.,2011).Arsenic is well known for its carcinogenicity in kidney,lung,liver,skin,and bladder.At high concentrations,As causes gastrointestinal problems and arsen-icosis,which arise mainly via consumption of water containing As and its subsequent accumulation in the body(Sharma and Sohn, 2009;Villaescusa and Bollinger,2008).Therefore,the WHO rec-ommends As concentration of10m g/L as the upper permissible limit in water(WHO,2011).This limit is applicable in India,Japan Taiwan,USA and Vietnam(Hug et al.,2008;Reddy and Roth,2013) while other countries like China,Bangladesh,and most of South American nations have permitted a higher concentration of50m g/L (Camacho et al.,2011;Chakraborti et al.,2010).A report prepared by UNICEF consultant Ravenscroft(2007)estimates that natural As pollution of groundwater and surface water affects more than140 million people in at least70countries worldwide.A large popula-tion of South Asia is exposed to As toxicity(Fig.2).When two different types of harmful contaminants are ingested,they may function independently or synergistically or antagonistically to one another(Chouhan and Flora,2010).While the harmful effects of As and F individually have been widely studied,the exposure to both together has received little atten-tion.Rao and Tiwari(2006)reported that As and F in combina-tion affect integrity of cells genetic material more than the individual exposure.In animal studies for rats,co-exposure of As and F even at low concentrations resulted in decreased comet tail and detrimental effect on liver and kidney(Flora et al.,2009; Mittal and Flora,2006).Wang et al.(2007)reported that chil-dren's growth and intelligence were severely inﬂuenced by high concentrations of As or F.Hence,it is important to remove these toxicants from potable water.Despite the extreme seriousness of the issue,very less data exist on the populations facing simul-taneous toxicity of As and F.This article analyzes the genesis of the combined presence of geogenicﬂuoride and arsenic in groundwater and drinking water as well as the treatment methods for their removal.Also,it reviews the effectiveness of several treatment methods when these two contaminants are present together.2.Occurrence of F and As in groundwatersThere is an evidence of the presence ofﬂuoride in different latitudes such as south-east of Africa,United States,the Middle East of Asia,South America,and Asian countries.However,China and India are the worst affected countries(Fig.S-1,supplementary information).For past few decades,the areas with arid and semi-arid climates are suffering from the scarcity of water due to the fact that the uptake of groundwater is in far excess than water recharge as well as excessive evaporation leading to decreased availability of water (Jakariya et al.,2003,2007).Sinceﬂuoride primarily originates from ﬂuoride rich rocks,concentrations ofﬂuoride are directly propor-tional to the extent of leaching/dissolution of crystalline minerals through water e rock interactions.The sources ofﬂuoride in groundwater throughﬂuoride-rich rocks are(i)ﬂurospar(CaF2) from lime stones,sand stone,sedimentary rocks;(ii)cryolite Fig.1.Estimated population exposed to F contamination in selected countries(Â103)(CDC,1993;Diaz-Barriga et al.,1997;Fewtrell et al.,2006).*no recent data available.S.V.Jadhav et al./Journal of Environmental Management162(2015)306e325307(Na 3AlF 6)from igneous rocks,granite;(iii)ﬂuorapatite (Ca 5(PO 4)3F)from igneous rocks,metamorphic rocks,and (iv)sellaite (MgF 2)from bituminous dolomite e anhydrite rock.Arsenic occurs in the environment in several oxidation states as stated earlier.Parsa and Shahidi (2010)stated that unusual large proportions of As are present in potentially soluble forms.Studies have shown that over exploitation of shallow (or main)aquifer has been the source of many arsenic problems (Jakariya and Bhattacharya,2007).Fig.S-2(supplementary information)shows the identi ﬁed regions of world having arsenic contamination of groundwaters.Very high concentrations of As can be seen pre-dominant in Mexico,USA,China,Bangladesh,Vietnam and Pakistan.Recently,groundwaters in Japan and Korea were also found to be contaminated with As (Ahn,2012;Yoshizuka et al.,2010).The major mechanism responsible for the As contamina-tion in groundwater is desorption from iron oxides or hydroxides from natural rocks and their reductive dissolution (Kim et al.,2012;Li et al.,2012).Fig.S-3(supplementary information)shows the co-occurrence of ﬂuoride and arsenic worldwide.Arsenic and F are found to co-exist in groundwaters in China,Argentina,Mexico,and Pakistan among other countries where As concentrations up to 5300m g/L and F up to 29mg/L are noticed in the same groundwaters (Jing et al.,2012).Recently,Australia,Japan,Korea,and Chile have also been shown with elevated levels of As and F co-occurrence (Fig.S-3,supplementary information )(Ahn,2012;Chakraborti et al.,2011;Fernandez-Turiel et al.,2005;Richards et al.,2009;Yoshizuka et al.,2010).Table S-1(supplementary information)presents detailed information on As and F co-contamination in groundwa-ters with respect to geographical location and lists key ﬁndings of various publications.Arsenic is found to be frequently associated with ﬂuoride in shallow aquifers around the world.A high concentration of As is also found in semi-arid regions that contain oxidized groundwater (Currell et al.,2011).The correlations between these two toxicants are dissimilar as per redox potential (Kim et al.,2012).The Fe-hydroxides adsorption capacity of F and As decreases with in-crease in pH,releasing both components into groundwater indi-cating that the Fe-(hydr)oxides play an important role for hosting the co-contamination (Streat et al.,2008).High concentrations ofAs (10e 5300m g/L)and F (51e 7340mg/L)are reported in shallow groundwaters,indicating potential risk of arsenicosis and ﬂuorosis in Chaco-Pampean plain of Argentina where As and F borne dis-eases affect ~2e 8million people (Nicolli et al.,2012).The concur-rent presence of As and F in groundwater is connected to volcanic eruptions,geothermal currents and mining activities (Alarcon-Herrera et al.,2013).Selected data on As and F levels in waters from around the world are presented in Table 1.Excessive rates of evaporation in arid and semi-arid regions lead to generation of saline groundwaters and alkaline pH,which are related to high concentrations of As and F in groundwater (Nicolli et al.,2008a,2008b ).Presence of Na þand HCO 3Àare also correlated with simultaneous presence of high concentrations of As and F.On the other hand,studies in Yuncheng Basin,northern China,have found a moderately positive correlation between pH and As and F concentrations indicating that high pH may favor desorption of F and As,while HCO 3Àmay act as a sorption competitor (Currell et al.,2011).These authors also have found a strong correlation between As and F,suggesting that the enrich-ment of As and F is governed by a common mechanism and/or due to a set of aquifer conditions.The sediment geochemical data of 2046samples from northern Mexico were generated by Alarcon-Herrera et al.(2013).The region is surrounded by large basins ﬁlled with alluvium,and dominated by a large number of wells.A strong correlation between As and F was observed in these localities where As has been reported in high amounts indicating co-occurrences of As and F.Nicolli et al.(2008b)have found a strong link between F and As concentration in deep wells in Cordoba province,Argentina.3.Methodologies for toxicant removal from groundwaters 3.1.Fluoride removalBasically,de ﬂuoridation of water can be introduced at two organizational levels;as household de ﬂuoridation,carried out by individual households for their own water consumption,and as community de ﬂuoridation,carried out at village,town,sub-urban area levels.As it has been previously reported in the reviews of Ayoob et al.(2008)and Mohapatra et al.(2009),manytechnologiesFig.2.Estimated population exposed to As contamination in selected countries (Â103)(Ravenscroft,2007).S.V.Jadhav et al./Journal of Environmental Management 162(2015)306e 325308have been employed and are currently being used to carry out the F removal from potable water such as ion exchange(Chubar,2011), adsorption(Bhatnagar et al.,2011;Gong et al.,2012a),coagulation and electrocoagulation(Behbahani et al.,2011;Gong et al.,2012b), and membrane processes(Richards et al.,2010).Each of these technologies has its own merits and demerits.Fig.3presents a typical initial concentration range treatable by a particular tech-nology and its corresponding removal efﬁciency.It can be seen that chemical treatment is capable of treating very high concentrations of F;since it cannot bring F concentration within WHO permissible limit,it can be coupled with other technologies as a primary treatment method(Islam and Patel,2007).This approach offers advantage of better life expectancy of secondary or tertiary treat-ment due to reduced load.3.1.1.Chemical precipitation/coagulation(CPC)The roots of the deﬂuoridation processes can be traced to the early1930's since when researchers from all around the world have been trying to develop a sustainable cost effective technology to reduce F concentrations in water.CSIR-NEERI in Nagpur,India, developed the Nalgonda process for the deﬂuoridation of drinking water.The technique has been operated in a number of villages in India asﬁll and draw type and hand-pump attached plants (Meenakshi and Maheshwari,2006).Chemical precipitation technique involves addition of aluminum salts along with lime to the F rich water followed by ﬂocculation and sedimentation orﬁltration.In theﬁrst step,lime reacts with F impurities such as NaF,HF,etc.to form insoluble calciumﬂuoride.Ca(OH)2þ2FÀ/CaF2þ2OHÀ(1) Ca(OH)2(aq)þ2NaF(aq)/CaF2(s)þ2NaOH(aq)(2) Essentially,in the second step,aluminum sulfate or aluminum chloride or both together,is added.Aluminum salt acts as aTable1Data on groundwaters containing As and F. Parameter Geographical locationZhijiliang,Inner Mongolia,China Meoqui City,Chihuahua,Mexico2sitesKalalanwala andKot Asadullah,Pakistan2depthsUn-known site,IndiaKyushu,Japan16sitesChihuahua,MexicoGeumsanCounty,KoreaMin.andMax.Santiago del EsteroProvince,ArgentinapH7.577.27.487.557.2 2.8-8.57.17 5.788.68 6.4-9.3 Temperature, C13e e25±225±220.5e251223eArsenic,mg/L 1.100.1340.0750.2350.060.130.01e3.230.435e1130.01-15Fluoride,mg/L 1.59 5.9 4.811 1.4750.04e3.7611.8e7.540.7-22Sodium,mg/L119e e630273e 1.1e1501250 1.6466.8eCarbonate,mg/L10121126857410e e e8.64239eTurbidity,NTU311 1.4 1.1e e320e1e e eHardness,mg/L51.524.558.3e e e e e e e eReference Zhang et al.Pinon-Miramonteset al.Farooqui et al.Devi et al.Yoshizuka et al.Nevarez et al.Ahn Bhattacharya et al. Year20032003200720082010201120122006Fig.3.Fluoride removal performance of various technologies.Where,CPC-11¼Reardon and Wang(2000);CPCþAD-12¼Islam and Patel(2007);ECþEF-13¼Zuo et al.(2008); EC/EF-14¼Bennajah et al.(2009);EC-15¼Khatibikamal et al.(2010);AD-16¼Cengeloglu et al.(2002);AD-17¼Tripathy et al.(2006);AD-18¼Tripathy and Raichur(2008);AD-19¼Thakre et al.(2010);AD-110¼Ganvir and Das(2011);IE-111¼Solangi et al.(2009);IE-112¼Viswanathan and Meenakshi(2009);RO-113,NF-113¼Kettunen and Keskitalo (2000);RO-114,NF-114¼Dolar et al.(2011);ED-115¼Zeni et al.(2005);ED-116¼Ergun et al.(2008).S.V.Jadhav et al./Journal of Environmental Management162(2015)306e325309coagulant and is often being used for viable and effective F removal from water.The basicity present in water with alum yields an aluminum salt,[Al(OH)3],which is insoluble.It has been reported that the pH of the contaminated water increases up to12;however,the best F removal is achieved be-tween the pH range of6e7(Aoudj et al.,2012).The dose of alum is typically around20fold the lime required.Theﬂuoride from groundwater can be removed upto96%from the initial concen-tration of109mg/L using lime(Reardon and Wang,2000).Similar ﬁndings on F removal by calcium salts have been conﬁrmed by Jadhav et al.(2014).The Nalgonda technique has the advantages of low initial cost and effectiveness.However,Nalgonda technique is not recom-mended owing to high maintenance cost,unpleasant water taste, requirement of large area for sludge drying,and a very high residual aluminum(Liu et al.,2013).The residual aluminum in waters ranges from2.01to6.86mg/L under different operating conditions. Any amount over0.2mg/L of aluminum in drinking water can cause serious health problems,including dementia(Shrivastava and Vani, 2009).3.1.1.1.Electrocoagulation(EC)and electrocoagulation/ﬂotation(EF)Passage of an electric current into an aqueous medium helps to destabilize suspended,dissolved and emulsiﬁed impurities,and the process is known as electrocoagulation(EC).During the past decade,the use of electrocoagulation(EC)as well as electro-coagulation/ﬂotation(EF)process is on the rise.It can be effectively employed to treat oily wastewaters,dye and textile industry ef-ﬂuents and removal of organic matter,heavy metals andﬂuoride (Hu et al.,2003,2005).EC is advantageous since no impurities are introduced and useful contents existing in raw water can be retained during deﬂuoridation.Electrochemical cell(also known as“electrolytic cell”)is the basis of electrocoagulation and electroﬂotation techniques.An electrocoagulation reactor typically consists of an electrolytic cell with an anode and a cathode.Passage of the electric current leads to the deterioration of the anode and cathode which may be made of the same or different materials and which act as‘sacriﬁcial electrodes’(Mollah et al.,2001).It is reported that the following three routes are adapted for the electrocoagulation/ﬂotation tech-niques,namely,(i)electrode oxidation,(ii)generation of gas bub-ble,(iii)ﬂotation and sedimentation ofﬂocs(Emamjomeh and Sivakumar,2009).Electrocoagulation deals with electrochemical production of destabilization agents that lead to neutralization of electric charge to remove pollutants.Charged particles coagulate together to make a mass.Effective removal of pollutants byﬂotation and sedimen-tation is augmented by metal based coagulants having a similar effect as the metal cations produced by anode.The key reactions are as follows(Zuo et al.,2008):Al/Al3þþ3e at the anode(3) Al3þþ3H2O/Al(OH)3þ3Hþ(4) Al(OH)3þxFÀ/Al(OH)3Àx F xþxOHÀ(5) 2H2Oþ2e/H2þ2OHÀat the cathode(6) Evolution of hydrogen bubbles at the cathode improves the F ion mass transfer rates and also leads toﬂoating of the aluminum complex(Al(OH)3Àx F x)ﬂocs at the top of the electrocoagulation system.Effective F removal can be achieved by isolating the aluminum complex from the aqueous phase periodically.There-fore,formation of aluminum complex[Al(OH)3Àx F x]results in the deﬂuoridation of the source waters.Developments during the past decade have demonstrated that EC is an effective technique for F removal in drinking waters and industrial wastewaters (Abuzaid et al.,2002;Essadki et al.,2009;Han and Kwon,2002; Holt et al.,2002;Khatibikamal et al.,2010;Mameri et al.,2001). It has been reported that by using EC,ﬂuoride concentrations could be reduced to values lower than1.5mg/L from initial con-centrations ranging from10to20mg/L(Bennajah et al.,2009). Another advantage is that EC processes are characterized by lower amount of sludge and absence of chemical handling.However, electricity is required to operate the plant which increases its operating cost.3.1.2.Adsorption(AD)Amongst other methods,adsorption(AD)is a conventional technique which is widely used for deﬂuoridation of water because it is economical,robust,environmentally benign and efﬁcient.New advanced materials have been developed recently for effective and cheapﬂuoride removal from potable water.Many low cost adsor-bents have also been employed forﬂuoride removal like alumina, red mud,clays,soils,activated carbon,calcite,brick powder,acti-vated coconut-shell,activated kaolinites,oxides ores,modiﬁed chitosan,bone char,and some other low cost materials(Mohapatra et al.,2009).A common mechanism by which the adsorption of F ions occurs onto solid particles can be given by the following three steps:(i)mass transfer of F ions to the external surface of the adsorbent,(ii)F ion adsorption onto external particle surface, and(iii)intra-particle diffusion of F ions from the exterior surface and possible exchange with elements on the pore surface inside particles(Fan et al.,2003).Activated alumina happens to be one of the most popular and widely used solid adsorbents for the deﬂuoridation of potable water and many reports are available on large-scale installations(Chauhan et al.,2007). Alumina is popular amongst other adsorbents because it main-tains its structural stability without shrinkage,swelling or disintegration in water(Serbezov et al.,2011).Granules of acti-vated alumina having a very high surface area of~200e300m2/g possess a substantial number of active sites to facilitate adsorp-tion.Atﬁrst,the acidiﬁcation of alumina is carried out with HCl as follows:Alumina¤H2OþHCl/Alumina¤HClþH2O(7) (¤represents activated state)The acidic alumina on contact with F ions displaces the Cl ions and gets bonded to alumina.Alumina¤HClþNaF/Alumina¤HFþNaCl(8) Further,regeneration is carried out byﬁrst treating the adsor-bent with alkaline solution followed by acid wash.A few reports mention that F removal is due to ion-exchange as well as adsorp-tion following both Freundlich and Langmuir isotherms(Bansiwal et al.,2010;Ghorai and Pant,2005).Meenakshi and Maheshwari(2006)have reported the devel-opment of domestic and hand-pump units with activated alumina for deﬂuoridation.The unit consists of~20L capacity bucket,ﬁtted with a micro-ﬁlter at the bottom containing5kg of activated alumina.Activated alumina coated with manganese dioxide was found to reduce F concentration to0.2mg/L from10mg/L at pH5.5 (Tripathy and Raichur,2008).Magnesia-modiﬁed activated alumina granules were also used successfully to reduce F from drinking water with a maximum adsorption capacity of10.12mg/g at F concentration of about150mg/L(Maliyekkal et al.,2008).S.V.Jadhav et al./Journal of Environmental Management162(2015)306e325 310Biswas et al.(2009)characterized the synthetic hydrated iron(III)e tin(IV)mixed oxide(HITMO)for F removal.It was found that the F adsorption capacity decreased with increasing initial pH from3.0to 5.0,and it remained constant up to pH of7.5.Also,the high bicar-bonate content showed adverse effects on F removal by HITMO.It was the pseudo-second order kinetics with multiple-stage which deﬁned rate-limiting step.The equilibrium data was best portrayed by Langmuir isotherm model with adsorption capacity10.50mg/g and adsorption energy~9kJ/mol.Theﬂuoride-rich material was successfully regenerated up to75%in desorption studies.All in all,the choice of the adsorbent seems to be reliant on factors such as the ability to adsorb from dilute solutions,pH, removal duration,adsorbent stability,regeneration,and adsorption capacity in the presence of competing ions,and the economics (Mohapatra et al.,2009).The main disadvantage of the adsorption process is that the adsorbent gets exhausted soon and considerable time is required for regeneration.Moreover,regeneration steps leads to secondary pollution because F containing aqueous solution is discarded as a waste.3.1.3.Ion exchange(IE)Many reports have highlighted the efﬁcacy of ion exchange with other techniques(Onyango et al.,2005,2006).Usually,ion ex-change technique removes F by adsorption rather than exchanging ions.The fundamental reason is that theﬂuoride concentration is comparably lower than other ions present in water.Cation ex-change resins are more selective for F removal than anion exchange resins(Meenakshi and Viswanathan,2007).However,the deﬂuoridation capacity and selectivity for F is dependent on the type of resin.The loading of metal ions inﬂuences theﬂuoride removal drastically,owing to variations in their properties(Luo and Inoue,2004).Thus,it is difﬁcult to maximize the deﬂuoridation capacity(DC)of ion exchange resins while simultaneously enhancing the F selectivity.Viswanathan and Meenakshi(2008)used a widely available ion exchanger Indion FR10that has considerable F removal capabil-ities.It was chemically altered with Ce3þ,Fe3þ,La3þ,and Zr4þspecies to understand their selectivity for deﬂuoridation.The maximum deﬂuoridation capacity of all the modiﬁed resins was measured around0.5mg/g.The authors pointed out that the deﬂuoridation was due to electrostatic adsorption and complexa-tion.Recently,they have modiﬁed Indion FR10into Naþand Al3þtypes by loading the metal ions in Hþtype of resin(Viswanathan and Meenakshi,2009).Chubar et al.(2005)obtained a new ion exchanger from double hydrous oxide(Fe2O3$Al2O3$x H2O)by the sol e gel method from easily available raw materials,which was used for adsorption of FÀ, ClÀ,BrÀ,and BrO3Àsimultaneously.It was found that the pH effect of FÀand BrÀwas dependent on ion speciation.Sufﬁcient capacity for sorption of all these anions was in the pH range of3e10.The F adsorption of88mg/g was the highest among all species.Solangi et al.(2009)modiﬁed anionic resin Amberlite XAD-4™,in which the amino group is introduced to the aromatic ring of the resin and utilized effectively for F extraction.They found that the modiﬁed resin was efﬁcient particularly at pH9and was also effective in the presence of other anions such as BrÀ,NO2À,NO3À,HCO3Àand SO42À. Subsequently,the authors modiﬁed Amberlite XAD-4™resin by adding thio-urea binding sites into the aromatic rings(Solangi et al.,2010).The modiﬁed resin had high efﬁciency for F removal from aqueous solution at a wide range of pH from4to10.The resin could be regenerated several times and used as an ion exchange material inﬁlters for F removal from potable water.Ion exchange treatment has a great F removal potential(up to 95%).However,the resins are costly,thereby making the entire ion exchange method expensive.Regenerating the resins is simple but it generates a large volume of F loaded waste which again is a problem.3.1.4.Membrane technology(MT)In recent years,membrane-based techniques have got a lot of attention due to their performance and reliability in operation for the removal of F from groundwater.At present,nanoﬁltration(NF), reverse osmosis(RO),and electro-dialysis(ED)are the most pop-ular membrane processes for F removal.3.1.4.1.Reverse osmosis(RO)and nanoﬁltration(NF)In membraneﬁltration,the water containing high concentration of pollutants is passed through a semipermeable membrane.The membrane discards atoms on the criteria of size and electric charge. The pollutants are removed from the water and collected at retentate side;whereas,clean water is recovered through permeate.In RO,pressure greater than the natural osmotic pressure is applied to the concentrated side of the membrane.Kettunen and Keskitalo(2000)evaluated the performance of a low energy RO membrane with a NF membrane.A membraneﬁltration plant of 16e25m3/h capacity was constructed in Laitila,Finland,to control F and Al concentration in drinking water.The comparative removal was above95and76%for F and Al,respectively and RO needed ~1.6bar excess feed pressure than that used in NF.Sehn(2008) studied a large scale RO plant in southern Finland for three years. The system was operated at a pressure of6e11atm and tempera-ture range of5e11 C,which was until recently possible only with NF membrane.This resulted in low power requirements and the plant was operated at80%recovery without using a scaling inhib-itor.Richards et al.(2011)used four commercial NF/RO membranes in Australian groundwaters.Their investigations showed the in-ﬂuence of solar irradiance levels on retention of F,Mg,NO3,K and Na where convection/diffusion predominated the retention.About 85%of all solutes were retained during solar irradiance conditions.Nanoﬁltration(NF)is a relatively new process in contrast to RO, ultraﬁltration(UF)and microﬁltration(MF),and it is emerging as a practically functional technology in treating industrial wastewa-ters.Nanoﬁltration is characterized by attributes between reverse osmosis(RO)and ultraﬁltration(UF).Regardless of the similarity in operation with RO,NF is operated at a comparatively lower pres-sure yielding identical permeateﬂux even at lesser pressure.NF removes less than60%of the monovalent ions as opposed to90%by RO membranes.RO can completely demineralize water with very low or practically no selectivity for monovalent ions but it suffers from high operating pressure,low permeateﬂux and high energy requirements(Alarcon-Herrera et al.,2013).In particular,fractional deﬂuoridation can be attained by altering the operating variables of the NF process;while simultaneously keeping the required F con-tent in the water.Two of the popular NF membranes,NF90and NF400,were used to practically remove F from groundwaters(Tahaikt et al.,2007). The quality of water achieved by the NF400membrane was found to be satisfactory especially for lower F content.For higher F con-tent,a double pass was required to reduce it to an acceptable level. Further,these authors calculated the economics of a100m3/h NF plant corresponding to a recovery rate of84%,F removal of97.8% and pressure of10bar(Elazhar et al.,2009).The capital cost was estimated to be748,000V with operating cost of0.212V per m3. Hu and Dickson(2006)investigated performance of negatively-charged commercial thin-ﬁlm composite(TFC)nanoﬁltration membranes.They conﬁrmed that higher pressures exhibit sub-stantial positive effect on F removal along with increase in theﬂux.Malaisamy et al.(2011)modiﬁed a commercially available NF membrane by layer-by-layer assembly of alternating poly-electrolyte thinﬁlms in order to promote removal and selectivityS.V.Jadhav et al./Journal of Environmental Management162(2015)306e325311。

磁性壳聚糖凝胶微球对水中Pb(Ⅱ)的吸附性能

中国环境科学 2018,38(4)：1364~1370 China Environmental Science 磁性壳聚糖凝胶微球对水中Pb(Ⅱ)的吸附性能蒲生彦1,2,3*,王可心1,2,马慧1,2,杨曾1,2,候雅琪1,2,陈虹宇1,2(1.成都理工大学,地质灾害防治与地质环境保护国家重点实验室,四川成都 610059；2.成都理工大学,国家环境保护水土污染协同控制与联合修复重点实验室,四川成都 610059；3.香港理工大学,土木及环境工程学系,中国香港)摘要：以壳聚糖为原材料,通过原位共沉淀法和柠檬酸钠交联法制备了一种新型多孔磁性壳聚糖凝胶微球吸附剂CS-citrate/Fe3O4.利用扫描电镜(SEM)、透射电镜(TEM)、傅里叶红外光谱(FTIR)、热重分析(TG)对吸附剂进行了表征.结果表明,吸附剂内部具有发达的孔隙结构,并均匀分布有平均直径为(4.79±1.09) nm的Fe3O4纳米颗粒;吸附剂中引入Fe3O4后,仍存在羟基、氨基和羧基等功能基团,且吸附剂磁性良好可用于磁场分离;吸附剂对Pb(II)的吸附等温线和动力学研究表明,吸附过程以化学吸附为主,最大吸附容量可达178.25mg/g.关键词：多孔结构；生物质吸附剂；磁性壳聚糖；凝胶微球；重金属中图分类号：X703 文献标识码：A 文章编号：1000-6923(2018)04-1364-07Adsorption properties ofmagnetic chitosan hydrogelmicrospheres to Pb(II) from aqueous solutions. PU Sheng-yan1,2,3*, WANG Ke-xin1,2, MA Hui1,2, YANG Zeng1,2, HO U Ya-qi1,2, CHEN Hong-yu1,2 (1.State Key Laboratory of Geological Prevention and Geological Environment Protection, Chengdu University of Technology, Chengdu 610059, China；2.State Environmental Protection Key Laboratory of Synergetic Control and Joint Remediation for Soil & Water Pollution, Chengdu University of Technology, Chengdu 610059, China；3.Department of Civil and Environment Engineering, The Hong Kong Polytechnic University, Hong Kong, China). China Environmental Science, 2018,38(4)：1364~1370Abstract：In this study, the magnetic porous chitosan hydrogel microsphere was fabricated by a combination of in situ-coprecipitation and sodium citrate crosslinking technique using chitosan as raw material. The scanning electron microscopy (SEM), transmission electron microscopy (TEM), Fourier transform infrared spectroscopy (FTIR) and thermogravimetric analysis (TG) were conducted for the characterization of this novel adsorbent. The hydrogel microsphere present a well developed porous inner structure and the Fe3O4nanoparticles with an average diameter of (4.79±1.09) nm dispersed uniformly. The functional group of chitosan, the hydroxyl, amino and carboxyl groups, remained after the introduction of the Fe3O4, and the magnetic adsorbent could be separated by the addition of external magnetic field. The adsorption isotherm and kinetic study for the Pb () removal from the aquatic environment indicatingⅡthat the adsorption process was dominated by the chemical adsorption and the maximum adsorption capacity was calculated as 178.25mg/g.Key words：porous structure；biomass adsorbent；magnetic chitosan；hydrogel microsphere；heavy metal重金属在生物物质循环和能量交换中不能被分解破坏,只能改变其物理化学形态或转移其存在位置,加之重金属在环境中的迁移转化几乎涉及了所有可能的物理、化学和生物过程,因而治理难度很大[1].伴随现代工业的快速发展,重金属废水已成为对环境污染最严重的工业废水之一[2].现有重金属废水处理技术,如离子交换法、电解法、化学沉淀法等常规方法普遍存在处理工艺复杂,运行成本高,对低浓度重金属废水处理效果差的问题[3].相比之下,吸附法则具有适用范围广、反应速度快、可适应不同反应条件、环境友好等优点,受到了研究人员的高度关注[4].近年,研究较多的吸附材料有活性炭[5]、沸石[6]、膨润土[7]等.这些吸附剂对废水中重金属有一定的去除效果,但吸附完成后难以与水体分离,容易造成环境收稿日期：2017-09-18基金项目：国家自然科学基金资助项目(51408074,41772264)* 责任作者, 教授, pushengyan@4期蒲生彦等：磁性壳聚糖凝胶微球对水中Pb(Ⅱ)的吸附性能 1365二次污染.与非生物质吸附剂相比,生物质基吸附剂富含大量吸附功能基团,对重金属离子有很强的吸附能力和较高吸附容量,而且具有资源丰富,可再生易降解,环境友好成本低的优点,较为适合水中重金属离子的富集与分离[8-9].常见的天然高分子吸附剂,如:壳聚糖[10]、纤维素[11]、木质素[12]等,其中以壳聚糖及其衍生物研究最为活跃.壳聚糖是一种成本低廉,环境友好的天然生物高分子,其分子主链上大量氨基、羟基等官能团可络合金属离子,且这些官能团具有良好的反应性,可功能化改性[13].若将壳聚糖赋予磁性后,采用磁分离技术可使吸附剂回收和再生变得简易[14].目前,磁性壳聚糖吸附剂的制备方法有原位共沉淀法[15]、微乳液法[16]和水热法[17]等,其中原位共沉淀法通过溶液中的化学反应直接得到均一的材料,相比其他方法制备过程简单且环境友好,是应用最普遍的方法之一[18].目前已有的原位共沉淀法包括电喷雾技术[19],静电液滴(ESD)技术[20]和反向共沉淀法[21]等.本研究采用原位共沉淀法结合柠檬酸钠交联制备了一种新型多孔磁性壳聚糖凝胶微球,在对其微观结构、物化性能进行充分表征的基础上,选取Pb(II)作为目标污染物考察了该凝胶微球的吸附性能,以期能为水中重金属富集去除提供一种新的思路和方法.1 材料与方法1.1 材料试剂:壳聚糖(Chitosan,CS,脱乙酰度80%~ 95%)购于上海阿拉丁生化科技股份有限公司;冰醋酸、氢氧化钠、柠檬酸钠购于成都科龙化学试剂厂;六水合氯化铁、四水合氯化亚铁和硝酸铅购于志远化学试剂厂;实验用水均采用超纯水.仪器:KW-400恒温水浴振荡器,上虞佳星仪器厂;SCIENTZ-50F冷冻干燥机,宁波新芝生物科技股份有限公司;GGX-9火焰原子吸收分光光度计,北京海光仪器有限公司.1.2 多孔磁性壳聚糖凝胶微球的制备将0.8g壳聚糖溶于24mL 2%的乙酸溶液中,机械搅拌30min,使得壳聚糖充分溶解;之后向溶液中加入2mL摩尔比为2:1的Fe3+/Fe2+混合溶液,继续搅拌30min,溶液由亮黄色变为棕红色后,将混合溶液用蠕动泵滴入NaOH/柠檬酸钠混合浸泡液(NaOH 1.25mol/L, 柠檬酸钠0.1mol/L)中,静置陈化10h;磁分离后用超纯水多次洗涤,除去残余的NaOH和柠檬酸钠,冷冻干燥30h.无磁壳聚糖凝胶微球(CS)在不加Fe3+/Fe2+混合溶液的条件下以相同方法制得作为实验对照组.1.3 表征方法采用德国Sigma300型扫描电子显微镜(SEM)观察样品表面形貌,采用日本FEI Tecnai-G20型透射电子显微镜(TEM)观察样品内部形貌,采用美国Nicolet-1170SX型傅里叶红外光谱仪(FTIR)进行红外谱图分析,采用美国STA6000型热重分析仪(TGA)考察在壳聚糖凝胶微球中引入Fe3O4纳米颗粒的热力学效应.1.4 对Pb(II)的静态吸附实验将0.05g磁性壳聚糖凝胶微球投加到50mL200mg/L的Pb(II)溶液中,在25℃下恒温振荡(150r/min),测定吸附量q随时间t的变化情况.吸附量采用公式(1)进行计算.()/tq c c V M=− (1) 式中:c0和c t为在Pb(II)溶液的初始浓度和吸附t时间后的浓度,mg/L;V为Pb(II)溶液的体积,L;M为吸附剂的投加量,g.2 结果与讨论2.1 多孔磁性壳聚糖凝胶微球制备首先,壳聚糖溶液与Fe3+/Fe2+(摩尔比为2:1)经螯合作用形成Fe3+-CS-Fe2+混合溶胶,然后,将混合溶胶通过蠕动泵滴入NaOH/柠檬酸钠混合浸泡液形成凝胶微球.在此过程中,发生Fe3+/Fe2+原位共沉淀反应生成Fe3O4纳米颗粒,壳聚糖和柠檬酸钠发生交联反应.最后,将制得吸附剂用于水中重金属离子的静态吸附,并利用外加磁场将吸附剂分离回收,从而达到回收再利用,减少二次污染的目的.多孔磁性壳聚糖凝胶微球制备及重金属吸附实验流程如图1所示.本实验所制得的磁性壳聚糖凝胶微球平均1366 中国环境科学 38卷粒径约为(2.91±0.65)mm,将其冷冻干燥处理后平均粒径约为(2.42±0.51)mm,干燥后平均粒径约减小16.8%,冷冻干燥后凝胶微球较好地保留原有圆球状形态及内部多孔结构.图1 磁性壳聚糖凝胶微球制备与重金属吸附实验流程Fig.1 Schematic illustration of preparation of magnetic chitosan hydrogel microspheres and its adsorption process2.2 多孔磁性壳聚糖凝胶微球吸附剂表征图2(a)~(d)为不同放大倍数下多孔磁性壳聚糖凝胶微球吸附前后外观扫描电镜图,由图可知多孔磁性壳聚糖凝胶微球为形态良好、表面光滑的球型结构,在吸附后吸附剂表面粗糙,覆盖有Pb(II)的结晶产物,孔道堵塞严重;图2(e)和(f)为多孔磁性壳聚糖凝胶微球和壳聚糖凝胶微球的内部结构扫描电镜图.由图可知,与壳聚糖凝胶微球内部紧密的结构相比,磁性壳聚糖凝胶微球内部具有良好的多孔结构,增大了吸附剂的比表面积,有利于吸附作用发生.图2 壳聚糖凝胶微球和磁性壳聚糖凝胶微球SEMFig.2 SEM characterization results of chitosan hydrogel microsphere and magnetic chitosan hydrogel microspheres磁性壳聚糖凝胶微球(a)吸附前外观,×30;(b)吸附前外观,×300;(c)吸附后外观,×30;(d)吸附后外观,×300;(e)内部结构SEM 图,×250;(f) 壳聚糖凝胶微球内部结构SEM 图,×7004期蒲生彦等：磁性壳聚糖凝胶微球对水中Pb(Ⅱ)的吸附性能 1367为深入了解多孔磁性壳聚糖凝胶微球中Fe3O4纳米颗粒的形态,对样品进行了TEM分析.由图3可知,Fe3O4纳米颗粒在壳聚糖微球内部分布较均匀,未出现明显团聚现象,其平均粒径约为(4.79 ±1.09)nm(图3(b)).图3 磁性壳聚糖凝胶微球TEM图Fig.3 TEM characterization results of magnetic chitosanhydrogel microspheres(a)TEM图;(b)Fe3O4纳米颗粒粒径分布图图4为壳聚糖凝胶微球及吸附前后多孔磁性壳聚糖凝胶微球红外光谱图.壳聚糖、柠檬酸钠和Pb(II)之间的相互作用会影响特征峰的位置和强度,在壳聚糖凝胶微球的光谱中,1082cm-1, 1027cm-1两处为C-OH键的伸缩振动吸收峰, 1383cm-1处为伯醇组-C-O键的伸缩振动吸收峰.1425cm-1处为C-N键的伸缩振动峰.壳聚糖固有的O-H和N-H伸缩振动峰出现在3444cm-1附近,在多孔磁性壳聚糖凝胶微球的两个光谱中也可观察到这一较宽的吸收峰.在吸附前多孔磁性壳聚糖凝胶微球的光谱中,1648cm-1处的N-H 伸缩振动吸收峰移动到1640cm-1处.由于柠檬酸钠的交联和Fe3O4与壳聚糖之间的弱相互作用,导致酰胺峰强度降低.在吸附前、后磁性壳聚糖凝胶微球光谱中,586cm-1处出现了Fe3O4的特征吸收峰,对应的是Fe-O的伸缩振动峰,说明磁性纳米颗粒Fe3O4已成功嵌入吸附剂中.而吸附了Pb(II)的凝胶微球光谱图4(b)和(d)中,1383cm-1和1425cm-1处特征吸收峰形状发生变化表明Pb(II)离子和壳聚糖发生络合反应,同时说明多孔磁性壳聚糖凝胶微球的羟基、氨基和羧基可以高效吸附金属阳离子.图4 吸附前后壳聚糖凝胶微球和磁性壳聚糖凝胶微球的红外谱Fig.4 FTIR spectra of pure chitosan hydrogelmicrospheres, and magnetic hydrogel chitosanmicrospheres before and after adsorption(a)壳聚糖凝胶微球Cs;(b) 壳聚糖凝胶微球/吸附后Cs/Pb;(c)磁性壳聚糖凝胶微球Cs-citrate/Fe3O4;(d) 磁性壳聚糖凝胶微球/吸附后Cs-citrate/Fe3O4/Pb通过热重分析表征了壳聚糖凝胶微球中引入的Fe3O4纳米颗粒的热力学效应.图5为壳聚糖凝胶微球和多孔磁性壳聚糖凝胶微球的热重曲线图.图5 磁性壳聚糖凝胶微球的热重分析曲线Fig.5 Thermo gravimetric curves of magnetic chitosanhydrogel microspheres由热重分析曲线可知,壳聚糖凝胶微球重量损失发生在3个阶段.第一阶段,当温度升至90℃左右,吸附剂中的自由水及通过氢键形成的结合水减少;第二阶段,在90~320℃范围内,壳聚糖发生分解;第三阶段,壳聚糖发生碳化分解,1368 中国环境科学 38卷在800℃时所对应的重量为热解最终产物残余碳.多孔磁性壳聚糖凝胶微球在25~120℃的范围内脱去自由水和结合水;在320℃时壳聚糖完全分解;在600℃时,Fe3O4与碳反应生成单质铁;在800℃时残余重量为碳和单质铁.多孔磁性壳聚糖凝胶微球的分解起始温度比壳聚糖凝胶微球高,说明Fe3O4的存在有效地提高了吸附剂的热稳定性.2.3 吸附时间和初始浓度对吸附效果的影响图6(a)讨论了在0~840min内壳聚糖凝胶微球和多孔磁性壳聚糖凝胶微球对Pb(II)吸附量的变化.图6 时间和初始浓度对Pb(II)吸附的影响Fig.6 Influence of time and initial concentration on adsorption of Pb(II)可以看出,壳聚糖凝胶微球的吸附作用主要发生在0~120min内,在120min后达到吸附平衡,平衡吸附量为16.1mg/g.多孔磁性壳聚糖凝胶微球对Pb(II)的吸附与壳聚糖凝胶微球呈现相同的变化趋势,但平衡吸附量达到了45.3mg/g,为壳聚糖凝胶微球的2.8倍,这是磁性复合吸附剂的高度多孔结构提供了较大的比表面积,使更大数目的活性基团与Pb(II)接触产生的结果.吸附剂吸附量在0~120min内升高较快,说明Pb(II)与多孔磁性壳聚糖凝胶微球的基团发生螯合反应,被成功地吸附到样品表面上,使得溶液中Pb(II)浓度下降.随着吸附反应的进行,吸附到多孔磁性壳聚糖凝胶微球的Pb(II)逐渐占据了大部分活性基团,导致活性基团的数目下降,在120min时吸附量趋于平衡.本研究将铅离子初始浓度设置为100、200、300、400、500mg/L对多孔磁性壳聚糖凝胶微球的吸附性能进行了考察(图6(b)).多孔磁性壳聚糖凝胶微球对不同初始浓度Pb(II)的吸附呈现类似的变化趋势,随着初始浓度增大,多孔磁性壳聚糖凝胶微球对Pb(II)的吸附量逐渐增大.2.4 吸附动力学及吸附等温线表1 Pb(II)吸附动力学方程的拟合Table 1 Fitting results of lead ions adsorption kinetics equations准一级动力学准二级动力学初始浓度(mg/L) q exp(mg/g)k1(min-1)(×10-2) q cal(mg/g) R2k2(min-1)(×10-2)q cal(mg/g) R2 100 30.88 0.805 16.53 0.794 0.163 31.34 0.9983 200 45.73 1.976 23.52 0.845 0.263 46.23 0.9995 300 64.97 0.799 23.34 0.562 0.136 65.40 0.9990 400 89.16 2.003 51.94 0.799 0.123 90.01 0.9994 500 105.99 1.651 35.39 0.775 0.182 106.61 0.9999 2.4.1 吸附动力学采用准一级和准二级动力学模型对动力学数据进行拟合,计算出相应的速4期蒲生彦等：磁性壳聚糖凝胶微球对水中Pb(Ⅱ)的吸附性能 1369率常数,研究其吸附过程的动力学行为并探讨吸附机理.所用拟合方程的线性表达式如下: e e 1ln()ln t q q q k t −=− (2)22e e /1/()/t q k q t q =+ (3)式中:q t 为t 时刻吸附剂对Pb(II)的吸附量,mg/g;q e为平衡吸附量,mg/g;k 1为准一级速率常数,min -1; k 2为准二级速率常数,mg/(g ⋅min).q e 、k 2可分别由截距和直线斜率求得.分析结果见表1和图7.准一级动力学相关系数最高为0.845,而准二级动力学相关系数均高于0.99,因此准二级动力学方程能更好地描述整个吸附过程.这证实了多孔磁性壳聚糖凝胶微球吸附剂对Pb(II)的吸附为化学吸附,比表面积是吸附的重要影响因素.l n (q e-q t )图7 吸附动力学曲线 Fig.7 Absorption kinetic curve2.4.2 吸附等温线使用Langmuir 和Freundlich 吸附等温线模型来解释吸附机理. Langmuir 方程假设吸附过程为单层吸附,线性表达式如下:e e m e m /1/()/c q kq c q =+ (4) 式中:q e 表示吸附质的吸附量,mg/g;c e 表示其在溶液中的平衡浓度,mg/L;b 为Langmuir 吸附平衡常数,L/mg;q m 为在吸附剂上单层形成的最大吸附能力,mg/g.c e/q e (g /L )c e (mg/L)图8 等温吸附曲线 Fig.8 Sorption isothermFreundlich 等温线是用于描述非均相表面的经验方程,它的线性表达式如下:e F e ln ln (1/)ln q K n c =+ (5) 式中:K F 是Freundlich 常数;1/n 为吸附指数.将Pb(II)起始浓度范围为100~500mg/L 的5组吸附实验数据进行吸附等温线拟合,结果见图8和表2. Langmuir 模型和Freundlich 模型均具有良好的拟合度,且后者线性拟合相关系数更高,大于0.95,可能是由于吸附剂表面基团分布不均,导致吸附过程呈现非均质吸附特性.吸附指数1/n 的值为0.3749,有报道称1/n <1表明吸附容易进行[22].相较其他生物质基吸附材料吸附铅离子研究,如甘蔗渣对铅最大吸附量为41.32mg/g [23],改性木质素磺酸钠对铅最大吸附量为55.22mg/ g [24],N -(2-磺乙基)壳聚糖对铅最大吸附量为1370 中国环境科学 38卷99.79mg/g [25],本研究中磁性壳聚糖微球对铅的吸附容量178.25mg/g 均大于上述吸附材料.表明磁性壳聚糖微球对Pb(II)具有良好的吸附效果.表2 Pb(II)吸附等温线拟合参数Table 2 Fittingof lead ions adsorption isotherm equationsLangmuir 等温吸附 Freundlich 等温吸附q m (mg/g)178.25K F (mg/g) 3.2541 k 0.0093 1/n 0.3749R 20.9219 R 2 0.95363 结论3.1 采用原位共沉淀法和柠檬酸钠交联法制备得到的多孔磁性壳聚糖凝胶微球内部孔隙丰富,比表面积大,磁性良好,可外加磁场分离.Fe 3O 4纳米颗粒在壳聚糖基质中分布均匀,增加了吸附剂的热稳定性.3.2 多孔磁性壳聚糖凝胶微球对水中Pb(II)具有良好的吸附性能,约在2h 达吸附平衡,当Pb(II)初始浓度从100mg/L 增加到500mg/L 时,平衡吸附量从30.88mg/g 增加到105.99mg/g.3.3 吸附剂对水中Pb(II)的吸附过程满足准二级动力学方程,并较好的符合Freundlich 等温吸附方程,最大吸附容量可达178.25mg/g.参考文献：[1] 唐黎,李秋华,陈椽,等.贵州普定水库沉积物重金属分布及污染特征 [J]. 中国环境科学, 2017,37(12):4710-4721. [2] 范小杉,罗宏.工业废水重金属排放区域及行业分布格局 [J].中国环境科学, 2013,33(4):655-662.[3] Fu F, Wang Q. 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三氯化铁和硫酸亚铁除砷的比较研究

在 pH 值为 8、n（ Fe） / n（ As） = 2、反应时间为 60 min 的条件下，采用三氯化铁处理模拟含砷废水；在 pH 值为 7、n（ Fe） / n（ As） = 1. 5、反应时间为 30 min 的条件下，采用硫酸亚铁处理模拟含砷废水。并对各个试样中的滤液和滤渣分别进行测定。其结果见表 1。
分别取 100 mL 模拟废水置于若干个烧杯中，按照不同的 n （ Fe ） / n （ As ）比例，依次加入 w（ FeCl3 ）为 35% 的三氯化铁溶液，用固体氢氧化钠调节 pH 值为 8，搅拌 1 h，过滤；在其他对比试验的烧杯中，按照不同的 n（ Fe） / n（ As）比例加入硫酸亚铁，用氢氧化钠调节 pH 值为 7，搅拌 0. 5 h，过滤；测各个样品滤液中的砷含量，计算砷去除率，结果见图 2。
蒋明磊等 . 三氯化铁和硫酸亚铁除砷的比较研究
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为 2 时的砷去除率和硫酸亚铁在 n（ Fe） / n（ As）为 1. 5 时的砷去除率均能达到 99% 以上。 2. 3 反应时间对砷去除率的影响
分别取 100 mL 模拟水样置于若干个烧杯中，按照 n（ Fe） / n（ As）为 2 加入 w（ FeCl3 ）为 35% 的三氯化铁溶液，用固体氢氧化钠调节 pH 值到 8；在其他对比试验的烧杯中，按照 n（ Fe） / n（ As）为 1. 5 加入硫酸亚铁，用氢氧化钠调 pH 值为 7。各个试样分别反应不同的时间过滤，测各个试样滤液中的砷含量，计算砷去除率，结果见图 3。
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Adsorption of heavy metal ion from aqueous single metal solution by chemically

Adsorption of heavy metal ion from aqueous single metal solutionby chemically modiﬁed sugarcane bagasseOsvaldo Karnitz Jr.a ,Leandro Vinicius Alves Gurgel a ,Ju´lio Ce ´sar Perin de Melo a ,Vagner Roberto Botaro a ,Taˆnia Ma ´rcia Sacramento Melo a ,Rossimiriam Pereira de Freitas Gil b ,Laurent Fre´de ´ric Gil a,*aDepartamento de Quı´mica,Instituto de Cie ˆncias Exatas e Biolo ´gicas,Universidade Federal de Ouro Preto,35400-000Ouro Preto,Minas Gerais,BrazilbDepartamento de Quı´mica,Instituto de Cie ˆncias Exatas,Universidade Federal de Minas Gerais,31270-901Belo Horizonte,Minas Gerais,BrazilReceived 22November 2005;received in revised form 28April 2006;accepted 2May 2006Available online 14July 2006AbstractThis work describes the preparation of new chelating materials derived from sugarcane bagasse for adsorption of heavy metal ions in aqueous solution.The ﬁrst part of this report deals with the chemical modiﬁcation of sugarcane bagasse with succinic anhydride.The carboxylic acid functions introduced into the material were used to anchor polyamines,which resulted in two yet unpublished modiﬁed sugarcane bagasse materials.The obtained materials were characterized by elemental analysis and infrared spectroscopy (IR).The sec-ond part of this reports features the comparative evaluation of the adsorption capacity of the modiﬁed sugarcane bagasse materials for Cu 2+,Cd 2+,and Pb 2+ions in aqueous single metal solution by classical titration.Adsorption isotherms were studied by the Freundlich and Langmuir models.Ó2006Elsevier Ltd.All rights reserved.Keywords:Adsorption;Modiﬁed sugarcane bagasse;Polyamines;Isotherm;Heavy metals1.IntroductionWater pollution is a major environmental problem faced by modern society (Baird,1995)that leads to eco-logical disequilibrium and health hazards (Kelter et al.,1997).Heavy metal ions such as copper,cadmium,lead,nickel,and chromium,often found in industrial waste-water,present acute toxicity to aquatic and terrestrial life,including humans.Thus,the discharge of eﬄuents into the environment is a chief concern.The methods commonly used to remove toxic heavy metal from municipal and industrial wastewater are based on the adsorption of ions onto insoluble compounds and the separation of the sed-iments formed.Many eﬀorts have been made recently toﬁnd cheaper pollution control methods and materials(Panday et al.,1985;Ali and Bishtawi,1997;Acemiog˘lu and Alma,2001).The new material world trends point to the importance of using industrial and agricultural residues as production starting materials.Reusing and recycling these residues can minimize the environmental problems associated with their build-up and reduce the use of noble starting materi-als.This trend has contributed to the reconsideration of the use of traditional biomaterials such as natural lignocellu-losic ﬁbers to substitute synthetic polymers,for example,since in many cases they have a better performance.Brazil is the world leading producer of sugarcane for both the alcohol and the sugar industries.These industries produce a large amount of sugarcane bagasse and although it is burned to produce energy for sugar mills,leftovers are still signiﬁcant.Thus,on account of the importance of0960-8524/$-see front matter Ó2006Elsevier Ltd.All rights reserved.doi:10.1016/j.biortech.2006.05.013*Corresponding author.Tel.:+553135591717;fax:+55315511707.E-mail address:laurent@iceb.ufop.br (L.F.Gil).Bioresource Technology 98(2007)1291–1297bagasse sugar as an industrial waste,there is a great interest in developing chemical methods for recycling it.Sugarcane bagasse has around50%cellulose,27%polyoses,and23% lignin(Caraschi et al.,1996).These three biological poly-mers have many hydroxyl and/or phenolic functions that can be chemically reacted to produce materials with new properties(Xiao et al.,2001;Navarro et al.,1996).Despite the many studies of the chemical modiﬁcation of cellulose published around the world in this area(Gurnani et al.,2003;Gellerested and Gatenholm,1999),only a few have investigated the modiﬁcation of bagasse sugar(Krish-nan and Anirudhan,2002;Orlando et al.,2002).This work describes the preparation and the evaluation of new chelating materials from sugarcane bagasse to adsorb heavy metal ions in aqueous solution.In a prelimin-ary study,it has been chosen to study the adsorption of Cu2+,Cd2+,and Pb2+.Theﬁrst part of this work describes the modiﬁcation of sugarcane bagasse with succinic an-hydride to introduce carboxylic functions to sugarcane bagasse and the chemical introduction of commercial linear polyamine via the formation of amide functions.It is well known that polyamines have powerful chelating properties, mainly towards ions such as Cu2+,Zn2+,and Pb2+(Bian-chi et al.,1991;Martell and Hancock,1996).The second part of this work evaluates the adsorption of Cu2+,Cd2+,and Pb2+onto three modiﬁed sugarcane bag-asses(MSBs)from aqueous single metal ion solutions by classical titration.The results were analyzed by the Lang-muir and Freundlich models(Ho et al.,2005).2.Methods2.1.MaterialsPolyamines ethylenediamine3and triethylenetetramine 4were used in this work.Succinic anhydride,1,3-diiso-propylcarbodiimide(DIC),and triethylenetetramine,from Aldrich,were used without puriﬁcation.Ethylenediamine and dimethylformamide were distilled before use.Pyridine was reﬂuxed with NaOH and distilled.2.2.Sugarcane bagasse preparationSugarcane bagasse was dried at100°C in an oven for approximately24h and nextﬁber size was reduced to pow-der by milling with tungsten ring.The resulting material was sieved with a4-sieve system(10,30,45,and60mesh). Then,the material was washed with distilled water under stirring at65°C for1h and dried at100°C.Finally,it was washed anew in a sohxlet system with n-hexane/ ethanol(1:1)as solvent for4h.2.3.Synthesis of MSBs1and2Washed and dried sugarcane bagasse(5.02g)was trea-ted with succinic anhydride(12.56g)under pyridine reﬂux (120mL)for18h.The solid material wasﬁltered,washed in sequence with1M solution of acetic acid in CH2Cl2, 0.1M solution of HCl,ethanol95%,distilled water,and ﬁnally with ethanol95%.After drying at100°C in an oven for30min and in a desiccator overnight,MSB1(7.699g) was obtained with a mass gain of53.4%.MSB2was obtained by treatment of1with saturated NaHCO3solu-tion for30min and afterwards byﬁltering using sintered ﬁlter and washing with distilled water and ethanol.2.4.Synthesis of MSBs5and6The process used to introduce amine functions was the same as that used to prepare MSB5and6.MSB1was trea-ted with5equiv of1,3-diisopropylcarbodiimide(DIC)and 6equiv of polyamine in anhydrous DMF at room tempera-ture for22h under stirring.Afterﬁltration,the materials were washed with DMF,a saturated solution of NaHCO3, distilled water,andﬁnally with ethanol.Next,they were dried at80°C in an oven for30min and in a desiccator overnight.2.5.Kinetic study of metal ion adsorption of MSBs2,5,and6Experiments with each material and metal ion were per-formed to determine the adsorption equilibrium time from 10to90min in10min intervals.The amount of100mg MSB was placed in a250-mL Erlenmeyer with100.0mL metal ion solution with concentration of200mg/L under stirring.The experiments were done at pHs5.8for Cu2+, 7.0for Cd2+,and6.2for Pb2+,optimal values to obtain the best adsorption.To adjust pH values,was added NaOH solution(0.01mol/L)into metal solutions with MSB.Afterﬁltration,metal ion concentration was deter-mined by EDTA titration.2.6.pH study of metal ion adsorption of MSBs2,5,and6Experiments with each material and metal ion were per-formed to determine the eﬀect of pH on ion adsorption.An amount of100mg MSB was placed into a250-mL Erlen-meyer with100.0mL of metal ion solution200mg/L under stirring.pH was calibrated with HCl or NaOH solutions (0.1–1.0mol/L).The reaction times used were30min (MSB2)or40min(MSB5and6)for Cu2+and Cd2+, and40min(MSB2)or50min(MSB5and6)for Pb2+. Metal ion concentration was determined afterﬁltration by EDTA titration.No signiﬁcative variation of pH was observed at the end of each experiment.2.7.Adsorption isotherms of MSBs2,5,and6Experiments were performed for each material and metal ion to determine adsorption isotherms.In each experiment,100mg of MSB was placed into a250-mL Erlenmeyer with100.0mL of metal ion solution in speciﬁc concentrations(between200mg/L and400mg/L)under stirring.Each experiment was performed at the pH of1292O.Karnitz Jr.et al./Bioresource Technology98(2007)1291–1297larger ion adsorption during the time necessary for equilib-rium (Tables 3and 4).After ﬁltration,the metal ion con-centration was determined by EDTA titration.2.8.Characterization of the new obtained materials MSB 1,2,5,and 6were characterized by IR spectro-scopy in a Nicolet Impact 410equipment with KBr.Elemental analyses were accomplished in Analyzer 2400CHNS/O Perkin Elemer Series II.3.Results and discussion3.1.Synthesis of MSBs 1,2,5,and 6The synthesis route used to prepare MSBs 1,2,5,and 6are presented in Scheme 1.Prewashed sugarcane bagasse was succinylated for various periods of time.The degree of succinylation of the bagasse ﬁbers was determined by measuring the quantity of acid function.The results are shown in Fig.1.The concentration of carboxylic functions per mg of bagasse was determined by retro titration.For this,MSB 1was initially treated with an excess solution of NaOH (0.01mol/L)for 30min.Soon afterwards the material was ﬁltered and the obtained solution was titrated with an HCl solution (0.01mol/L).The highest degree of succinylation was reached after 18-h ing this reaction time,sugarcane bagasse was succinylated to pro-duce MSB 1,which presented a weight gain of 54%and a concentration of carboxylic acid function per mg of 3.83·10À6mol.Next,MSB 1was treated with a saturated NaHCO 3solution to produce MSB 2.Starting from MSB 1,two polyamines were introduced:ethylenediamine 3and triethylenetetramine 4.The method-ology used to introduce the polyamines was the same for the two MSBs 5and 6,as shown in Scheme 1.Concentra-tions of 2.4·10À6mol (5)and 2.6·10À6mol (6)of amine function per mg of material were determined by back titra-tion with excess HCl solution.The introduction of the amine functions was also veriﬁed by IR spectroscopy (Table 1)and elemental analysis (Table 2).3.2.Characterization of MSBs 1,5,and 6Characterization of carboxylated MSB 1was accom-plished by IR spectroscopy.The spectrum of unmodiﬁed sugarcane bagasse and MSB 1are presented in Fig.2.The spectrum of MSB 1displayed two strong bands at 1740and 1726cm À1in relation to that of unmodiﬁed sug-arcane bagasse.This demonstrated the presence of two types of carbonyl functions,one relative to carboxylic acid and another relative to the ester.The acid and ester IR bands indicate that succinic anhydride acylated theO.Karnitz Jr.et al./Bioresource Technology 98(2007)1291–12971293hydroxy group of bagasse to generate an ester bond with consequent release of a carboxylic acid functional group.The spectra of MSBs5and6(Figs.3and4,respectively) showed three new strong bands at1550–1650cmÀ1(see data in Table1)corresponding to the presence of amide and amine functions,and one band at1060cmÀ1 corresponding to C–N stretch.The bands at1635and 1650cmÀ1(Fig.3)correspond to the axial deformation of the carbonyl of the amide function and the angular deformation of the N–H bond of the amine function.The band at1575cmÀ1corresponds to the angular deformation of the N–H bond of the amide function.The band at 1159cmÀ1(Fig.4)corresponds to the asymmetric stretch of C–N–C bond.The main bands observed in all MSBs are presented in Table1.MSB elemental analysis data presented in Table2show a modiﬁcation in the carbon and hydrogen composition of MSB1and a larger proportion of nitrogen as the number of amine functions in the used polyamine increases.3.3.Study of adsorption of Cu2+,Cd2+and Pb2+on MSBs2,5,and6The study of the MSB adsorption properties was accom-plished for each material and metal ion.A kinetic study and an adsorption study as a function of pH wereﬁrst carried out.3.3.1.Eﬀect of contact timeThe kinetic study of MSB2with Cu2+,Cd2+,and Pb2+ ions in aqueous solution is presented in Fig.5.Adsorption equilibrium was reached after20min for Cd2+ions.A time of30min was chosen for all studies of MSB2with Cd2+. The adsorption equilibrium times chosen for pH and con-centration dependent experiments are presented in Table3.Similar studies were accomplished for MSBs5and6for Cu2+,Cd2+,and Pb2+.The results are presented in Table3.3.3.2.pH EﬀectThe removal of metal ions from aqueous solutions by adsorption is dependent on solution pH as it aﬀects adsor-Table1Main IR spectrum bands observed in MSBs1,5,and6MSB Main bands observed(cmÀ1)11740,172651745,1650,1635,1575,1423,1060 61738,1651,1635,1560,1400,1159,1060 Table2Elemental analysis of MSBs1,2,5,and6C(%)H(%)N(%) Sugarcane bagasse43.98 6.020.13MSB145.41 5.620.10MSB238.04 5.140.01MSB544.01 6.51 2.21MSB646.88 6.65 3.431294O.Karnitz Jr.et al./Bioresource Technology98(2007)1291–1297bent surface charge,the degree of ionization,and the species of adsorbates.The study of adsorption of Cd 2+,Cd 2+,and Pb 2+on MSB 2as a function of pH was accom-plished with the reaction times given in Table 3;the results are presented in Fig.6.The adsorption of the three metal ions increases with the increase in pH.Maximum removal of Cd 2+was observed above pH 6and in the case of Pb 2+and Cu 2,above pH 5and 5.5.Similar studies were accomplished for MSBs 5and 6and Cu 2+,Cd 2+and Pb 2+with similar results,as shown in Table 4.3.3.3.Adsorption isothermsThe Langmuir (Ho et al.,2005)(Eq.(1))and Freundlich (Eq.(2))isotherms were evaluated by adsorption experi-ments as a function of the initial metal ion concentrations in aqueous solution under equilibrium time and pH condi-tions given in Tables 3and 4.The results of each material and metal ion are presented in Fig.7(Langmuir)and Fig.8(Freundlich)and Table 5.c q ¼1Q max Âb þc Q maxð1Þln q ¼ln k þ1nln cð2ÞTable 3Adsorption equilibrium times of MSBs 2,5and 6MSB Equilibrium time (min)Cu 2+Cd 2+Pb 2+230304054040506404050Table 4pH of largest adsorption of MSBs 2,5and 6MSB pH of largest adsorption Cu 2+Cd 2+Pb 2+2 5.5–6.0 6.5–7.5 5.0–6.05 5.5–6.0 6.5–7.5 5.0–6.065.5–6.06.5–7.55.0–6.0O.Karnitz Jr.et al./Bioresource Technology 98(2007)1291–12971295where q(mg/g)is the concentration of adsorbed metal ions per gram of adsorbent,c(mg/L)is the concentration of metal ion in aqueous solution at equilibrium,Q max and b are the Langmuir equation parameters and k and n are the Freundlich equation parameters.High correlation coeﬃcients of linearized Langmuir and Freundlich equations indicate that these models can explain metal ion adsorption by the materials satisfactorily. Therefore,both models explained metal ion adsorption by MSBs2,5,and6as can be observed in Table5,with the exception of the Freundlich model for Pb2+adsorption by MSB2.The Langmuir isotherm parameter Q max indicates the maximum adsorption capacity of the material,in other words,the adsorption of metal ions at high concentrations. It can be observed in Table5that MSB5presents the larg-est Cu2+adsorption capacity while MSB6adsorbs Cd2+ and Pb2+the ngmuir parameter b indicates the bond energy of the complexation reaction of the material with the metal ion.It can be observed that MSB2presents the largest bond energy for Cu2+and Cd2+,while three materials do not diﬀer signiﬁcantly for Pb2.The Freundlich isotherm parameter k indicates the adsorption capacity when the concentration of the metal ion in equilibrium is unitary,in our case1mg/L.This parameter is useful in the evaluation of the adsorption capacity of metal ions in dilute solutions,a case closer to the characteristics of industrial eﬄuents.The values of k of MSB2and5are much similar for Cu2+and Cd2+ and much higher than that for MSB6.This shows the superiority of both materials in the adsorption of these metal ions in low concentrations.MSB5has a higher k value for Pb2+when compared to those of the other materials.These results were compared with those of Vaughan et al.(2001)for a commercial macroreticular chelating resin with thiol functional groups,Duolite GT-73.The Q max of Duolite GT-73for Cu2+,Cd2+,and Pb2+were 62mg/g,106mg/g,and122mg/g,respectively.Duolite GT-73exhibited Q max lower than those of MSBs(Table5).4.ConclusionsThrough a fast,eﬀective,and cheap methodology,it was possible to devise a strategy to introduce chelating func-tions(carboxylic acid and amine)to sugarcane bagasse. Modiﬁed sugarcane bagasses presented a good adsorption capacity for Cu2+,Cd2+,and Pb2+ions with maximum adsorption capacity observed for MSB6.It has been dem-onstrated that metal ion adsorption eﬃciency is propor-tional to the number of amine functions introduced into the material.MSB2,which contained only carboxylate functions,showed an eﬃciency similar to that of MSB5, a material of much more complex synthesis. AcknowledgementsWe thank FAPEMIG forﬁnancial support,CAPES and UFOP.Table5The Langmuir and Freundlich parameters for Cu2+,Cd2+and Pb2+ adsorptionMetalion MSB Langmuir FreundlichQ max (mg/g)b(L/mg)r2k(mg/g)n r2Cu2+21140.431191.623.90.919351390.1730.999898.315.80.906161330.0140.992722.8 3.640.9635Cd2+21960.1030.993459.4 4.160.977351640.0680.995762.8 5.490.983463130.0040.9528 5.15 1.630.9856Pb2+21890.1100.994566.0 4.660.757951890.1250.999914724.510.98163130.1210.9994121 5.210.8771296O.Karnitz Jr.et al./Bioresource Technology98(2007)1291–1297ReferencesAcemiog˘lu,B.,Alma,M.H.,2001.Equilibrium studies on adsorption of Cu(II)from aqueous solution onto cellulose.Journal of Colloid and Interface Science243,81–83.Ali,A.A.,Bishtawi,R.,1997.Removal of lead and nickel ions using zeolite tuﬀ.Journal of Chemical Technology and Biotechnology69, 27–34.Baird,C.,1995.Environmental Chemistry.W.H.Freeman and Company, New York.Bianchi,A.,Micheloni,M.,Paoletti,P.,1991.Thermodynamic aspects of the polyazacycloalkane complexes with cations and anions.Coordi-nation Chemistry Reviews110,17–113.Caraschi,J.C.,Campana,S.P.,Curvelo, A.A.S.,1996.Preparac¸a˜o e Caracterizac¸a˜o de Polpas Obtidas a Partir de Bagac¸o de Cana de Ac¸u´car.Polı´meros:Cieˆncia e Tecnologia3,24–29.Gellerested,F.,Gatenholm,P.,1999.Surface properties of lignocellulosic ﬁbers bearing carboxylic groups.Cellulose6,103–121.Gurnani,V.,Singh,A.K.,Venkataramani,B.,2003.2,3-Dihydroxypyri-dine-loaded cellulose:a new macromolecular chelator for metal enrichment prior to their determination by atomic absorption spectrometry.Analytical and Bioanalytical Chemistry377,1079–1086. Ho,Y.S.,Chiu,W.T.,Wang,C.C.,2005.Regression analysis for the sorption isotherms of basic dyes on sugarcane dust.Bioresource Technology96,1285–1291.Kelter,P.B.,Grundman,J.,Hage,D.S.,Carr,J.D.,Castro-Acun˜a,C.M., 1997.A discussion of water pollution in the United States and Mexico;with High School Laboratory Activities for the analysis of lead, atrazine,and nitrate.Journal of Chemical Education74,1413–1421. Krishnan,K.A.,Anirudhan,T.S.,2002.Removal of mercury(II)from aqueous solutions and chlor-alkali industry eﬄuent by steam activated and sulphurised activated carbons prepared from bagasse pith:kinetics and equilibrium studies.Journal of Hazardous Materials92,161–183. Martell, A.E.,Hancock,R.D.,1996.Metal complexes in aqueous solutions.Plenum,New York.Navarro,R.R.,Sumi,K.,Fujii,N.,Matsumura,M.,1996.Mercury removal from wastewater using porous cellulose carrier modiﬁed with polyethyleneimine.Water Research30,2488–2494.Orlando,U.S.,Baes,A.U.,Nishijima,W.,Okada,M.,2002.Preparation of chelating agents from sugarcane bagasse by microwave radiation as an alternative ecologically benign procedure.Green Chemistry4,555–557.Panday,K.K.,Gur,P.,Singh,V.N.,1985.Copper(II)removal from aqueous solutions byﬂy ash.Water Research19,869–873. Vaughan,T.,Seo,C.W.,Marshall,W.E.,2001.Removal of selected metal ions from aqueous solution using modiﬁed corncobs.Bioresource Technology78,133–139.Xiao,B.,Sun,X.F.,Sun,R.,2001.The chemical modiﬁcation of lignins with succinic anhydride in aqueous systems.Polymer Degradation and Stability71,223–231.O.Karnitz Jr.et al./Bioresource Technology98(2007)1291–12971297。

弱磁场强化零价铁去除水中砷的效果

５期
胡一帆等：弱磁场强化零价铁去除水中砷的效果
１０７５
４ｍｇ·ｋｇ－１便可致命［４］．从２００８年７月起，我国《生活饮用水卫生标准》中对砷执行了更严格的安全标准，将砷的限制浓度从５０ μｇ·Ｌ－１下调至１０ μｇ·Ｌ－１．人类饮用水源地的砷污染问题是当前人类面临的最严峻的环境污染问题，它严重威胁着人类饮用水的安全［５⁃６］，因此必须高度重视水体中砷污染的问题，以此预防地方性砷中毒现象．
近年来的研究表明，外加弱磁场可以明显改善零价铁的反应活性［１５⁃１７］，且该方法无需额外能量和化学药
品输入，操作简单，绿色环保，具有广阔的应用前景．为了探究磁场促进零价铁去除砷的效果和影响因素的作用，本文做了如下的研究：（１）探究磁场对
不同初始ｐＨ、不同粒径条件下，零价铁去除砷的反应动力学规律，分析磁场的作用．（２）通过对反应剩余固体的微观表征和反应过程中铁离子、ｐＨ的变化，探究零价铁的腐蚀程度的变化，分析磁场促进零价铁去除砷的机理．
于北京伊诺凯科技有限公司，用去离子水稀释至浓度１０００ μｇ·Ｌ－１，作为砷储备液．实验所用的ＨＣｌ和ＮａＯＨ均为分析纯，购买于国药集团化学试剂有限公司．磁场由一块环形永磁铁提供，磁铁外径６０ｍｍ，内径３２ｍｍ，厚度１０ｍｍ，最大磁场强度为２０ｍＴ．１．２实验方案１．２．１零价铁粉除砷批实验
∗国家自然科学基金（４１２７２２６１）资助．ＳｕｐｐｏｒｔｅｄｂｙｔｈｅＮａｔｉｏｎａｌＮａｔｕｒａｌＳｃｉｅｎｃｅＦｏｕｎｄａｔｉｏｎｏｆＣｈｉｎａ（４１２７２２６１）．
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Adsorption of arsenic (V) by iron (III)-modified natural zeolitic tuff

ORIGINAL PAPERAdsorption of arsenic (V)by iron (III)-modiﬁednatural zeolitic tuffTanja Stanic ´ÆAleksandra Dakovic ´ÆAleksandar Zˇivanovic ´ÆMagdalena Tomas ˇevic ´-C ˇanovic ´ÆVera Dondur ÆSonja Milic´evic ´Received:25October 2007/Accepted:7April 2008/Published online:14May 2008ÓSpringer-Verlag 2008Abstract Adsorption of arsenic (V)by natural zeolitictuff,modiﬁed with iron (III),was investigated.Also,theiron (III)adsorption characteristics by natural zeolitic tuffwas evaluated.It was determined that iron (III)adsorptionby starting zeolitic tuff was best represented by the Fre-undlich type of isotherm,having correlation coefﬁcient (r 2)of 0.990.Arsenic (V)adsorption by iron (III)-modiﬁedzeolitic tuff followed a nonlinear type of isotherm.The bestﬁt of the experimental data was obtained using the Lang-muir–Freundlich model (r 2=0.99),with the estimatedmaximum of arsenic (V)adsorption to iron (III)-modiﬁedzeolitic tuff of 1.55mg/g.Keywords Zeolite ÁClinoptilolite ÁIron (III)ÁArsenic (V)ÁAdsorptionIntroductionArsenic has been known as a toxic element for centuriesdue to the risk to plants,animals,and human health (Karim2000).In natural waters,arsenic exists predominantly ininorganic form,as trivalent arsenite,As (III),and penta-valent arsenate,As (V),while organic forms of arsenic are rarely quantitatively important (Thirunavukkarasu et al.2001;Dous ˇova ´et al.2006).Actual valence states of arsenic depend on the redox environment in water systems.Pentavalent arsenic is more prevalent in surface water while trivalent arsenic is more likely to occur in anaerobic ground waters (Xu et al.2002).The stability and domi-nance of arsenic compounds depend directly on the pH of the solution.Pentavalent arsenic is stable at pH 0–2as neutral H 3AsO 4,while H 2AsO 4-,HAsO 42-,AsO 43-exist as stable species in the pH intervals 3–6,7–11,and 12–14,respectively.In most countries,the arsenic level of water is limited to the value of 0.05mg/dm 3.Therefore,in order to reduce the arsenic concentration in water,water treatment for its removal from industrial wastes should be necces-sarily applied (Altundog ˘an et al.2002).Various treatment methods such as ion exchange,reverse osmosis,ultraﬁltration,adsorption,coagulation–precipitation and adsorption–coprecipitation by metals (predominately ferric chloride)followed by coagulation have been proposed for removal of arsenic from water (Harper and Kingham 1992;Viraraghavan et al.1992).Among others,one of the most effective method for removal of inorganic arsenic compounds from water is their adsorption by both natural and synthetic adsorbents (Altundog ˘an et al.2002;Dous ˇova ´et al.2003).The most commonly used adsorbents for this purpose are natural and modiﬁed alumosilicate minerals––zeolites,bentonite,kao-line,etc.The main advantage of these materials is their low cost,simple modiﬁcation,and availability in huge amounts in many parts of the world.Natural and synthetic alumosilicates modiﬁed with iron (II/III)are recently considered as the efﬁcient adsorbents for arsenic,based on the well-known strong adsorptionT.Stanic´(&)ÁA.Dakovic ´ÁM.Tomas ˇevic ´-C ˇanovic ´Áic´evic ´Institute for Technology of Nuclear and Other Mineral RawMaterials,P.O.Box 390,11000Belgrade,Serbiae-mail:t.stanic@itnms.ac.yuA.Zˇivanovic ´VMA Institute for Hygiene,Crnotravska 17,11000Belgrade,SerbiaV.DondurFaculty of Physical Chemistry,P.O.Box 137,11000Belgrade,Serbia 123Environ Chem Lett (2009)7:161–166DOI 10.1007/s10311-008-0152-3afﬁnity of As oxyanions to hydrated Fe(II–III)oxi-(hydroxides),through formation of very stable inner-sphere bidentate As–Fe complexes(Bruce et al.1998).Arsenic sorption to natural zeolites(unmodiﬁed and modiﬁed forms by acid washing or by addition of iron)was studied by Elizalde-Gonza´lez et al.(2001a,b).They found that both arsenite and arsenate were efﬁciently removed in the pH interval4–11after contact with zeolite ZH(clinoptilolite) and ZMA(clinoptilolite+erionite).More effective for arsenite retention among the unmodiﬁed zeolite samples was the natural clinoptilolite ZMS,while the mixed phase zeolite ZMA showed the highest efﬁciency in the iron-modiﬁed group.They also reported that removal of arsenite by some iron-modiﬁed zeolites was comparable with the amount removed by iron hydroxide.Elizalde-Gonza´lez et al.(2001c)also studied the effect of the content of iron added to clinoptilolites on the uptake of arsenic species, and found that zeolite samples with higher iron content did not exhibit greater uptake of the arsenate species.Dousˇova et al.(2006)investigated the sorption of arsenate on alu-mosilicates(natural metakaoline,natural clinoptilolite,and synthetic zeolite)in both untreated and Fe(II)-treated forms and found that sorption capacity of Fe(II)-treated sorbents increased signiﬁcantly in comparison to the untreated material.Batch removal of arsenate and arsenate from water by iron-treated activated carbon and natural zeolites(chabazite and clinoptilolite)was studied by Payne and Abdel-Fattah(2005).Activated carbon removed *60%while clinoptilolite and chabazite removed*20 and*50%of the arsenate,respectively.Arsenate removal by iron-treated activated carbon and clinoptilolite bestﬁt the Langmuir model while arsenate removal by iron-treated chabazite gave better Freundlich modelﬁts.Onyango et al. (2003)investigated sorption kinetics of arsenic(V)onto iron-conditioned synthetic zeolite and found that the adsorption reaction between arsenic and the binding sur-faces is the essential rate-controlling step.However,at the present time,very limited literature data are available on the iron(III)adsorption by natural zeolitic tuff and the arsenic(V)by iron(III)-modiﬁed clinoptilolita.In this paper,the adsorption of arsenic(V)by natural zeolitic tuff(clinoptilolite),modiﬁed with iron(III),was investigated.The aim of this research included:(1)the evaluation of iron(III)adsorption characteristics by natural zeolitic tuff and(2)determination of arsenic(V)adsorption mechanism by selected iron(III)-modiﬁed clinoptilolite.ExperimentalThe natural material used in these experiments was the raw zeolitic tuff from the Bala Mare deposit in Romania.After crushing and grinding,the sample was prepared below 0.2mm in size.The mineralogical composition of the starting zeolitic tuff was primarily clinoptilolite,with smaller amounts of quartz and plagioclase,as measured by X-ray powder diffraction analysis(XRPD).The total cation exchange capacity(CEC)of the zeolitic tuff was167meq/100g,measured with1M NH4Cl.Cal-cium(86meq/100g)and potassium(57meq/100g)were the dominant ions in exchangeable position,while mag-nesium(11meq/100g)and sodium(13meq/100g)were present at a lower amount in exchangeable position in starting zeolitic tuff.The chemical composition of starting zeolitic tuff was determined by atomic absorption spec-trophotometry(AAS)using Perkin–Elmer730instrument (Table1).Iron(III)chloride for adsorption and modiﬁcation experiments was supplied by Merck.For the evaluation of iron(III)adsorption by natural zeolitic tuff,the iron(III) adsorption experiment was carried out using the batch equilibrium by weighting20g of starting material in 50cm3laboratoryﬂasks.The initial iron(III)concentra-tions added to the samples were:0.00,2.22,3.34,5.34, 8.38,14.86,and18.52g/dm3.Twenty cm3of the initial concentrations were added to the zeolitic tuff;suspensions were shaken for15min,and then left in contact for24h, at constant temperature(20°C).After the reaction time, suspensions were centrifuged at10,000rpm for15min. The amount of iron(III)adsorbed by zeolitic tuff(mg/g) was calculated by the difference between iron(III)added and that remaining in equilibrium solutions(mg/dm3).Iron(III)-modiﬁed clinoptilolite was obtained by treat-ment of the starting material with iron(III)chloride(FeCl3) solution using the following procedure:an aqueous solu-tion(125cm3)of2.67g/dm3FeCl3was shaken with100g of zeolite at room temperature(20°C).The mixture was stirred for15min and then left for24h.Afterﬁltration,the iron(III)-modiﬁed zeolite was rinsed with distilled water, until Cl-ions were no longer detected,and dried at70°C. The iron(III)concentrations in solutions were measured by AAS.For adsorption of arsenic(V)by iron(III)-modiﬁed clinoptilolite,stock solutions of arsenate with differentTable1Chemical content of starting zeolitic tuffConstituents SiO2Al2O3Fe2O3CaO MgO TiO2Na2O K2O I.L. Content(%)66.7212.28 1.54 4.370.680.330.59 3.0710.38 I.L.ignition loss123concentrations were prepared from arsenate solution (Titrisol,Merck)using distilled water.The concentration of the arsenate is always given as the concentration of elemental arsenic.Adsorption isotherm for arsenic(V)on the iron(III)-modiﬁed clinoptilolite was obtained using the batch equilibration technique.In this experiment,1g of the iron(III)-modiﬁed clinoptilolite was shaken with100cm3 of arsenic(V)solutions with concentrations of:0.1,0.25, 0.5,1,2,3,5,and10mg/dm3.The suspensions were mixed for30min at room temperature(20°C)andﬁltered. The initial and the equilibrium arsenic(V)concentrations in the supernatant were determined by AAS.The amount of arsenic(V)adsorbed by iron(III)-modiﬁed clinoptilolite (mg/g)was calculated from the difference between the initial andﬁnal arsenic(V)concentration in aqueous solution after equilibrium.Results and discussionTo study the iron(III)adsorption by starting zeolitic tuff, the adsorption isotherm was ually,the equilibrium relationships between an adsorbent and an adsorbate are described by sorption isotherms,which rep-resent the ratio between the amount adsorbed and that remaining in the solution at aﬁxed temperature under equilibrium conditions(Lemic´et al.2006).The isotherm for iron(III)adsorption by natural zeolitic tuff,obtained by plotting the equilibrium concentration of iron(III)in solution,against the amount of iron(III)adsorbed per unit of weight of adsorbent,is presented in Fig.1a,while Fig.1b presents the percentage of adsorbed iron(III) concentrations by natural zeolitic tuff.It can be seen,from Fig.1a,that the amount of iron(III) adsorbed by the starting material increased with increasing the initial iron(III)concentration,and the iron(III) adsorption by starting zeolitic tuff followed a nonlinear type of isotherm.Within the investigated concentration range,the adsorption isotherm consists of two regions: in theﬁrst,the amount of adsorbed iron(III)increases gradually up to3.34mg/g,while in the second part,the adsorption does not reach the plateau,although the slope of the isotherm is gradually decreased.Furthermore,from Fig.1a,b,it can also be concluded that,when iron(III) concentration is lower that5.34g/dm3(5.34mg/g),the starting zeolitic tuff has the ability to adsorb almost the entire iron(III)solution concentration.Langmuir and Freundlich isotherms are the most com-monly used models,since they are applicable over a wide range of concentrations.Often,these isotherms are not able to explain equilibrium adsorption because of irregu-larities on the surface of the mineral(Jaroniec1983).The Langmuir isotherm describes adsorption on strongly homogeneous surfaces(which is not the case for natural zeolite);thus,for heterogeneous surfaces,the heterogeneity factor is introduced,resulting in the Freundlich or Lang-muir–Freundlich isotherm.Over the examined iron(III) concentration range,the experimental results wereﬁtted to the following isotherm equations:Langmuir:n¼KÂCÂM=ð1þKÂCÞð1ÞFreundlich:n¼ðKÂCÞbð2ÞLangmuirÀFreundlich:n¼ðKÂCÞbÂM=ð1þðKÂCÞbÞð3Þwhere n is the amount adsorbed per gram of adsorbent(mg/ g),K is the surface adsorption equilibrium constant(dm3/ mg),M is the maximum amount adsorbed per gram of adsorbent(mg/g),C is the equilibrium solution concen-tration(mg/dm3),and b is the heterogeneity factor which is related to the afﬁnity of the surface.Based on the correlation coefﬁcients,r2,the applica-bilities of different isotherm equations were compared.The bestﬁt of the experimental data was obtained using the Freundlich model,having a r2value of0.990(Table2).123Similar results were obtained(Sheta et al.2003)for iron (II)adsorption by natural zeolites(clinoptilolite,phillipsite, chabazite,and analcime)and bentonite.They found that sorption data followed the Freundlich adsorption equation and theﬁt was better for all the studied mineral species in comparison with theﬁt for the Langmuir equation.Addi-tionally,they also found that the percentage of Fe sorbed/ added increased with the increase in initial Fe added up to 250mg/dm3(5mg/g),and then decreased for all samples, and reported that most of the added Fe at initial concen-trations\250mg/dm3could be retained on the exchanged sites or precipitated as insoluble Fe compounds.Thus,to ensure quantitative adsorption of iron(III)by zeolitic tuff, in our study,the initial iron(III)concentration of3.34g/ dm3(3.34mg/g)was selected for production of iron(III)-modiﬁed clinoptilolite.The amount of Fe(III)in the obtained material,calculated from the difference between the amount of Fe(III)added and the amount of Fe(III) retained in the supernatant after adsorption,was3.34mg/g. The amount of Fe(III)in iron(III)-modiﬁed clinoptilolite, determined additionally by chemical analysis,was 3.32mg/g,conﬁrming that iron(III)was quantitatively adsorbed onto clinoptilolite.In order to preliminarily test the arsenic(V)adsorption by the starting zeolitic tuff and iron(III)-modiﬁed clinop-tilolite,the single arsenic(V)concentration was used.The pH of the arsenic(V)solutions during the experiments was found to be between6and7.It is well known that arsenic speciation is highly dependent on pH(Viraraghavan et al. 1992).Thus,in the investigated pH region,the anionic species of arsenic(V)(H2AsO4-and HAsO42-)are the dominant components sorbed by the obtained materials. The results of arsenic(V)adsorption(C0=0.1mg/dm3, C susp=10g/dm3)showed that the arsenic(V)adsorption index on the starting zeolitic tuff and iron(III)-modiﬁed clinoptilolite was50and[99%,respectively.Based on the fact that iron(III)-modiﬁed clinoptilolite can adsorb larger amounts of arsenic(V)than unmodiﬁed material,and that iron-modiﬁed minerals are recently used as efﬁcient adsorbents for arsenic anionic species,the iron(III)-mod-iﬁed clinoptilolite was selected for further investigations.For experiments of arsenic(V)adsorption by iron(III)-modiﬁed clinoptilolite,a wide range of arsenic(V)initial concentrations(0.1–10mg/dm3)was examined.These arsenic(V)concentrations were much higher than10l g/ dm3,which is the World Health Organization’s(WHO’s) recommended maximum level in drinking water(Thiru-navukkarasu et al.2001).In this study,a certain amount of iron(III)clinoptilolite was added to the solution containing each of eight different concentrations of arsenic(V).From kinetic experiments,it was found that adsorption of arsenic (V)was very fast and that most of the arsenic(V)was adsorbed in less than30min(Fig.2).Figure2shows that more than96.5%of initial arsenic(V)concentrations(C0) of2mg/dm3were adsorbed in theﬁrst15min,while after 30min,the residual arsenic(V)concentration remained unchanged.Thus,for isotherm experiments,the suspensions were mixed for30min at room temperature andﬁltered.The isotherms were obtained by plotting the equilibrium con-centration of arsenic(V)in solution,against the amount of arsenic(V)adsorbed per unit of weight of adsorbent.The arsenic(V)adsorption isotherm and the percentage of adsorbed arsenic(V)concentrations for iron(III)-modiﬁed clinoptilolite are presented in Fig.3a,b.As can be seen from Fig.3a,the increase of arsenic(V) adsorption by iron(III)-modiﬁed clinoptilolite,by increasing the initial arsenic(V)concentration,was observed.For low arsenic(V)initial concentration(\3mg/ dm3),iron(III)-modiﬁed clinoptilolite removes almost the entire concentrations of arsenic(V),whereas the percent-age of adsorption decreases with an increase in arsenic(V) initial concentration.The maximum adsorbed arsenic(V) by iron(III)-modiﬁed clinoptilolite,under our experimental conditions,was0.45mg/g.From the shape of the isotherm presented at Fig.3a,it can be seen that the arsenic(V) adsorption by iron(III)-modiﬁed clinoptilolite followed a nonlinear type of isotherm.Table2Parameters and correlation coefﬁcients of various isotherms for adsorption of iron(III)by natural zeolitic tuffParameters r2K(10-3dm3/mg),b=1M(mg/g)bLangmuir0.764 2.0228.53 Freundlich0.990 5.5790.17 Langmuir–Freundlich0.988 1.772910-549.310.19123Nonlinear isotherms were reported by Elizalde-Gonza-les(2001c)for adsorption of both arsenite and arsenate on several zeolites with different contents of iron.They found that some curves showed the saturation trend,while other curves exhibited a similar form involving an initial step slope followed by gentle upward growth.The adsorption of arsenite varied among the zeolitic materials,and was the highest for zeolite ZMS which contained the highest Fe2O3 amount(1.88%).Zeolites ZMS,ZMA(1.41%Fe2O3),and ZMT(1.38%Fe2O3)adsorbed similar amounts of arsenate, while ZME(0.74%Fe2O3)and ZH(0.60%Fe2O3)dis-played the lowest uptake capacity.The theoretically calculated maximum adsorption capacity(Langmuirﬁt to the data)were16.7and100l g/g on the zeolitic tuff ZMS for arsenite and arsenate,respectively.In our study,over the examined arsenic(V)concentra-tion range,the experimental results presented at Fig.3a wereﬁtted to the Langmuir(Eq.1),Freundlich(Eq.2), and Langmuir–Freundlich(Eq.3)isotherm equations. The bestﬁt of the experimental data was obtained using the Langmuir–Freundlich model,having an r2value of 0.99(Table3).The estimated maximum of arsenic(V) adsorption to iron(III)-modiﬁed clinoptilolite,based on Langmuir–Freundlichﬁt to the data,was1.55mg/g. ConclusionThe results presented in this paper showed that efﬁcient adsorbent for arsenic(V)was prepared by appropriate modiﬁcation of natural zeolitic tuff with iron(III)chloride. The evaluation of iron(III)adsorption characteristics by natural zeolitic tuff showed that the amount of iron(III) adsorbed by the starting material increased with increasing the initial iron(III)concentration,and iron(III)adsorption by starting zeolitic tuff followed a nonlinear type of iso-therm.The bestﬁt of the experimental data was obtained using the Freundlich model,having an r2value of0.990. The initial iron(III)concentration of3.34g/dm3(3.34mg/ g)was used for the production of iron(III)-modiﬁed clin-optilolite.From kinetic experiments,it was found that adsorption of arsenic(V)on iron(III)-modiﬁed clinoptilo-lite was very fast and that most of the arsenic(V)was adsorbed in less than30min.Adsorption of arsenic(V) by iron(III)-modiﬁed clinoptilolite followed a nonlinear adsorption isotherm.The bestﬁt for the experimental adsorption data were given by the Langmuir–Freundlich model,having an r2value of0.99.The estimated maximum of arsenic(V)adsorption,based on Langmuir–Freundlichﬁt to the data,was1.55mg/g.The availability of this mineral, high adsorption capacity for iron(III),simple modiﬁcation, and excellent arsenic(V)adsorption capacity makes this material suitable for potential practical applications.Acknowledgments Funding for this research was provided by the Ministry of Science of the Republic of Serbia under the project6702. ReferencesAltundog˘an HS,Altundog˘an S,Tu¨men F,Bildik M(2002)Arsenic adsorption from aqueous solutions by activated red mud.Waste Manag22:357–363Bruce AM,Scott EF,Sabine G(1998)Surface structures and stability of arsenic(III)on goethite:spectroscopic evidence for inner-sphere complexes.Environ Sci Technol32:2383–2388 Dousˇova´B,MachovicˇV,Kolousˇek D,Kovanda F,Dornicˇa´k V (2003)Sorption of As(V)species from aqueous systems.Water Air Soil Pollut149:251–267Dousˇova´B,Grygar T,Martaus A,Fuitova´L,Kolousˇek D,MachovicˇV(2006)Sorption of As V on aluminosilicates treated with Fe II nanoparticles.J Colloid Interface Sci302:424–431Table3Parameters and correlation coefﬁcients of various isothermsfor adsorption of As(V)by iron(III)-modiﬁed clinoptiloliteParameters r2K(dm3/mg),b=1M(mg/g)bLangmuir0.9212.0290.41Freundlich0.980.3260.24Langmuir–Freundlich0.990.012 1.550.29123Elizalde-Gonza´lez MP,Mattusch J,Einicke WD,Wennrich R(2001a) Sorption on natural solids for arsenic removal.Chem Eng J 81:187–195Elizalde-Gonza´lez MP,Mattusch J,Wennrich R(2001b)Application of natural zeolites for preconcentration of arsenic species in water samples.J Environ Monit3:22–26Elizalde-Gonza´lez MP,Mattusch J,Wennrich R,Morgenstern P (2001c)Uptake of arsenite and arsenate by clinoptilolite-rich tuffs.Microporous Mesoporous Mater46:277–286Harper TR,Kingham NW(1992)Removal of arsenic from waste-water using chemical precipitation methods.Water Environ Res 64(3):200–203Jaroniec M(1983)Current state in adsorption from multicomponent solutions of nonelectrolytes on solids.Adv Colloid Interface Sci 18:149–225Karim MM(2000)Arsenic in groundwater and health problems in Bangladesh.Water Res34:304–310Lemic´J,Kovacˇevic´D,Tomasˇevic´-Cˇanovic´M,Kovacˇevic´D,Stanic´T,Pfend R(2006)Removal of atrazine,lindane and diazinone from water by organo-zeolites.Water Res40:1079–1085Onyango MS,Matsuda H,Ogada T(2003)Sorption kinetics of arsenic onto iron-conditioned zeolite.J Chem Eng Jpn 36(4):477–485Payne KB,Abdel-Fattah TM(2005)Adsorption of arsenate and arsenite by iron-treated activated carbon and zeolites:effects of pH,temperature,and ionic strength.J Environ Sci Health Part A 40(4):423–749Sheta AS,Falatah AM,Al-Sewailem MS,Khaled EM,Sallam ASH (2003)Sorption characteristics of zinc and iron by natural zeolite and bentonite.Microporous Mesoporous Mater61:127–136 Thirunavukkarasu OS,Viraraghavan T,Subramanian KS(2001) Removal of arsenic in drinking water by iron oxide-coated sand and ferrihydrite-batch studies.Water Qual Res J Can36:55–70 Viraraghavan T,Jin YC,Tonita PM(1992)Arsenic in water supplies.Int J Environ Stud41:159–167Xu YH,Nakajima T,Ohki A(2002)Adsorption and removal of arsenic(V)from drinking water by aluminium-loaded Shirasu-zeolite.J Hazard Mater B92:275–287123。

Adsorption of heavy metal ions from aqueous solutions by activated carbon prepared

Short CommunicationAdsorption of heavy metal ions from aqueous solutions byactivated carbon prepared from apricot stoneM.Kobya a ,E.Demirbasb,*,E.Senturk a ,M.InceaaDepartment of Environmental Engineering,Gebze Institute of Technology,41400Gebze,TurkeybDepartment of Chemistry,Gebze Institute of Technology,41400Gebze,TurkeyReceived 17July 2004;received in revised form 24November 2004;accepted 10December 2004Available online 25February 2005AbstractApricot stones were carbonised and activated after treatment with sulphuric acid (1:1)at 200°C for 24h.The ability of the acti-vated carbon to remove Ni(II),Co(II),Cd(II),Cu(II),Pb(II),Cr(III)and Cr(VI)ions from aqueous solutions by adsorption was investigated.Batch adsorption experiments were conducted to observe the eﬀect of pH (1–6)on the activated carbon.The adsorp-tions of these metals were found to be dependent on solution pH.Highest adsorption occurred at 1–2for Cr(VI)and 3–6for the rest of the metal ions,respectively.Adsorption capacities for the metal ions were obtained in the descending order of Cr(VI)>Cd(II)>Co(II)>Cr(III)>Ni(II)>Cu(II)>Pb(II)for the activated carbon prepared from apricot stone (ASAC).Ó2005Elsevier Ltd.All rights reserved.Keywords:Apricot stone;Adsorption;Heavy metals;Aqueous solution;pH1.IntroductionHeavy metal ions such as cobalt,copper,nickel,chro-mium and zinc are detected in the waste streams from mining operations,tanneries,electronics,electroplating and petrochemical industries,as well as in textile mill products (Patterson and Passino,1987).Heavy metals have a harmful eﬀect on human physiology and other biological systems when they exceed the tolerance levels.Heavy metals are not biodegradable and tend to accu-mulate in living organisms,causing various diseases and disorders.The most widely used methods for removing heavy metals from wastewaters include ion exchange,chemical precipitation,reverse osmosis,evaporation,membrane ﬁltration and adsorption.Most of these methods suﬀer from some drawbacks,such as high capital and opera-tional cost or the disposal of the residual metal sludge,and are not suitable for small-scale industries.Many reports have appeared on the development of low cost activated carbon from cheaper and readily available materials (Bailey et al.,1999).Activated carbons,with their high surface area,micro porous character and chemical nature of their surface,have made them poten-tial adsorbents for the removal of heavy metals from industrial wastewater.Studies on the adsorption of heavy metals by acti-vated carbon and various low-cost materials have been reported in the literature.These include activated car-bon prepared from peat,coconut shells,coal (Paajanen et al.,1997),anthracite (Petrov et al.,1992),hazelnut shell activated carbon (Cimino et al.,2000;Demirbas,2003;Kobya,2004),kaolinite (Yavuz et al.,2003),coir-pith (Kadirvelu and Namasivayam,2003),bambara nut and rice husks (Ajmal et al.,2003),almond shells,olive and peach stones (Ferro-Garcia et al.,1988),and vari-ous commercially activated carbons (Netzer and0960-8524/$-see front matter Ó2005Elsevier Ltd.All rights reserved.doi:10.1016/j.biortech.2004.12.005*Corresponding author.Fax:+902627542385.E-mail address:erhan@.tr (E.Demirbas).Bioresource Technology 96(2005)1518–1521Hughes,1984)which have been used for removal of heavy metals from aqueous solutions.The purpose of this study was to remove selected heavy metals,namely Ni(II),Co(II),Cd(II),Cu(II), Pb(II),Cr(III)and Cr(VI)from aqueous solutions and to evaluate the inﬂuence of pH on activated carbon pre-pared from apricot stone(ASAC).2.MethodsChemical activation utilizes chemicals,such as H2SO4,H3PO4,ZnCl2,KOH and CaCl2,that have dehydration and oxidation characteristics(Kim et al., 2001).Carbonisation and activation are usually carried out simultaneously in the chemical activation process.Apricot stones were obtained from Malatya in Turkey.The material was ground in a micro hammercutter mill(Glen Mills)and sieved to a size range of 2.0mm·0.5mm particle size prior to activation.Chem-ical activation using H2SO4at moderate temperatures produces a high surface area and high degree of micro-porosity(Demirbas,2003).The materials were mixed in a1:1wt ratio with concentrated H2SO4,placed in an oven and heated to200°C for24h.After this,the sam-ples were allowed to cool to room temperature,washed with distilled water and soaked in1%NaHCO3solution to remove any remaining acid.The samples were then washed with distilled water until pH of the activated carbon reached6,dried at105°C for5h and sieved to obtain the desired particle size(1.00–1.25mm).The sur-face area of the activated carbons was measured by BET (Brunauer-Emmett-Teller nitrogen adsorption tech-nique).Characteristics of the carbon are presented in Table1.Higher surface area(>642m2/g)and lower mes-oporosity are obtained with H3PO4for carbonisation of apricot stone at higher temperature(Philip and Girgis, 1996)when it is compared with H2SO4,but H2SO4is al-most as eﬃcient as H3PO4in the carbonisation process.The size distribution of micropores of activated carbon is understood to be one of the critical factors determining its applicability(Blacher et al.,2000).The microstructures of the carbon were observed by SEM (Philips XL30S-FEG)and are shown in Fig.1.The sample was gold coated prior to SEM observation.This ﬁgure shows that the adsorbent had an irregular and porous surface,indicating relatively high surface areas. This observation is supported by the BET surface area of the activated carbon.The pore size distributions were calculated using t-plots(Gregg and Sing,1982).The porosity of the carbon was around74%.Batch experiments were conducted to investigate the eﬀect of pH on adsorption of metals on ASAC.All reagents used were of AR grade(Sigma-Aldrich,Ger-many).Salts used were cadmium sulphate for Cd(II), nickel sulphate for Ni(II),copper sulphate for Cu(II), cobalt nitrate for Co(II),lead nitrate for Pb(II),chro-mium nitrate for Cr(III)and potassium dichromate for Cr(VI).Samples were prepared by dissolving each metal salt at a known concentration in deionised water to obtain a stock solution.The initial pH of the solution was adjusted by using either0.1M NaOH or0.05M H2SO4.50ml solution of known concentration(C0) and initial pH was placed in a100ml screw-cap conical ﬂask with0.1g of adsorbent and was agitated at a speed of200rpm in a thermostatic shaker bath at25°C for 48h.The initial concentration of metal ions and corre-sponding concentrations afterﬁxed time periods were measured by atomic absorption spectrophotometry (Perkin Elmer SIMAA6000).Chromium was deter-mined spectrophotometrically using diphenylcarbazide. The solution pH was measured using a Mettler Toledo 340pH meter.The metal concentration retained in the adsorbent phase(q e,mg/g)was calculated by using the following equationqe¼ðC0ÀC eÞVW sð1Þwhere C0and C e are the initial andﬁnal concentrations of metal ion in solution(mg/l),V is the volume ofTable1The characteristics of the activated carbonParameters Value Bulk density(g/ml)0.43 Ash content(%) 2.21 pH 6.00 Moisture content(%)7.18 Surface area(m2/g)566 Solubility in water(%)0.85 Solubility in0.25M HCl(%) 1.22 Decolorising power(mg/g)22.8 Iodine number(mg/g)548 Particle size(mm)1.00–1.25Fig.1.SEM image of ASAC.M.Kobya et al./Bioresource Technology96(2005)1518–15211519solution(l)and W s is the mass of the adsorbent(g). Each experiment was carried out in duplicate and the average of two values was used in the calculations. The maximum diﬀerence between the two values was less than3%of the mean.3.Results and discussionThe pH of the solution is an important factor in determining the rate of surface reactions.The variation in adsorption capacity in this pH range is largely due to the inﬂuence of pH on the surface adsorption character-istics of ASAC.The eﬀect of pH on the adsorption of metal ions on the adsorbent is presented in Table2.For Cr(VI),the amount adsorbed decreased from 34.70to7.86mg/g as the pH increased from1to6. For the rest of the metal ions,the amount adsorbed increased from3.08to33.57mg/g for Cd(II),7.74to 30.07mg/g for Co(II),2.83to29.47mg/g for Cr(III), 2.50to27.21mg/g for Ni(II),4.86to24.21mg/g for Cu(II)and6.69to22.85mg/g for Pb(II),respectively, as the pH increased from1to6.The pH experiments also showed maximum removalof99.99%for Cr(VI)at pH1,99.86%for Pb(II)at pH3, 99.67%for Cd(II)at pH5,99.11%for Co(II)at pH6, 98.56%for Cr(III)at pH4,97.59%for Ni(II)at pH4 and96.24%for Cu(II)at pH4,respectively.Chromium exists mostly in two oxidation states which are Cr(VI)and Cr(III)and the stability of these forms is dependent on the pH of the system(Cimino et al.,2000;Selomulya et al.,1999).It is well known that the dominant form of Cr(VI)at pH2is HCrOÀ4.Increas-ing the pH will shift the concentration of HCrOÀ4toother forms,CrO2À4and Cr2O2À7.Maximum adsorptionat pH1.0indicates that it is the HCrOÀ4form of Cr(VI) which is the predominant species between pHs1and2.At pH values lower than3,there is excessive proton-ation of the carbon surface resulting in a decrease in the adsorption of Ni2+,Co2+,Cd2+,Pb2+and Cu2+(=M) ions.This is consistent with the results obtained by Petrov et al.(1992).On increasing the pH of M(II) solutions from3,the percentage removal increased and become quantitative over the pH range3–6.The increase in metal removal as pH increased can be explained on the basis of a decrease in competition between proton(H+)and positively charged metal ion at the surface sites,and by decrease in positive charge which results in a lower repulsion of the adsorbing metal ion.The removal eﬃciencies of metal ions are aﬀected by the initial metal ion concentration with the removal decreasing as the concentration increases at constant pH.In addition to that,at solution pH above3the pre-ponderance of OHÀgenerates a competition between the carbon surface and the solution OHÀions for M(II)ions,which causes a decrease in the adsorption of M(II)ions on the carbon surface.In other words, the increase in M(II)removal above pH3for the carbon may be due to the retention of M(OH)2species into pores of the carbon particles.A change in pH at the end of the adsorption experi-ments indicates that a lowerﬁnal pH was reached for higher adsorbent concentration(Table2).The increase in adsorbent concentration results in greater removal of metal ions from the solution,leading to a higher H+concentration and accounting for the decrease in theﬁnal pH value.As solution pH was increased,the onset of metal hydrolysis and precipitation began at a pH of3for all metal ions and precipitation was dominant at higher pHs(P5).This indicated that the adsorption capacity of the adsorbent is clearly pH dependent.4.ConclusionActivated carbon prepared from apricot stone,an agricultural waste,could be used as potential adsorbentTable2Adsorption of the metal ions onto ASAC at various pH valuesMetal Adsorptionparameters pH123456Nickel C e(mg/l)50.2533.65 6.66 1.33 1.100.20q e(mg/g) 2.5010.8024.3026.9626.9727.21Removal(%)9.0539.1087.9597.5997.6398.51pHﬁnal 1.09 1.91 3.46 4.73 6.00 6.15Cobalt C e(mg/l)45.2033.2711.40 2.09 1.240.54q e(mg/g)7.7413.7024.6429.3029.7230.07Removal(%)25.5145.1781.2196.5697.9699.11pHﬁnal 1.18 2.13 3.60 4.91 5.02 5.24Cadmium C e(mg/l)61.2031.627.457.580.220.22q e(mg/g) 3.0817.8729.9629.8933.5733.57Removal(%)9.1553.0688.9488.7599.6799.68pHﬁnal 1.03 2.08 3.58 4.33 5.65 5.47Lead C e(mg/l)32.3516.680.0620.0280.0400.034q e(mg/g) 6.6914.5222.8322.8522.8422.85Removal(%)29.2563.5299.8699.9499.9199.93pHﬁnal 1.02 2.12 4.36 5.73 6.53 6.50Copper C e(mg/l)39.9722.40 3.17 1.87 1.51 1.25q e(mg/g) 4.8613.6423.2623.9124.0824.21Removal(%)19.5554.9193.6196.2496.9697.48pHﬁnal 1.07 1.98 3.34 4.59 4.62 4.87Cr(III)C e(mg/l)53.8838.03 4.740.860.940.60q e(mg/g) 2.8310.7527.4029.3429.3029.47Removal(%)9.4936.1292.0498.5698.4298.99pHﬁnal 1.06 2.04 3.12 4.08 5.02 5.98Cr(VI)C e(mg/l)0.007 2.1028.2031.8757.1962.81q e(mg/g)34.7033.3020.5018.7011.357.86Removal(%)99.9996.9959.7154.4718.3010.27PHﬁnal 1.02 2.01 3.22 4.42 5.01 6.021520M.Kobya et al./Bioresource Technology96(2005)1518–1521for the removal of heavy metal ions from aqueous solu-tions.Batch experiments were conducted to assess the eﬀect of pH on ASAC.Adsorptions of the metal ions were found to be highly pH dependent and the results indicated that the optimum pH for removal was1for Cr(VI)while that for the rest of the metal ions varied from3to6.Cost analysis for the preparation of acti-vated carbon was not carried out,but as apricot stone is found in abundance in Turkey,carbon cost is expected to be economical.ReferencesAjmal,M.,Rao,R.A.K.,Anwar,S.,Ahmad,J.,Ahmad,R.,2003.Adsorption studies on rice husk:removal and recovery of Cd(II) from wastewater.Bioresour.Technol.86,147–149.Bailey,S.E.,Olin,T.J.,Bricka,R.M.,Adrian,D.D.,1999.A review of potentially low-cost sorbents for heavy metals.Water Res.33, 2469–2479.Blacher,S.,Sahouli,B.,Heinrichs,B.,Lodewyckx,P.,Pirard,R., Pirard,J.P.,2000.Micropore size distributions of activated ngmuir16,6754–6756.Cimino,G.,Passerini,A.,Toscano,G.,2000.Removal of toxic cations and Cr(VI)from aqueous solution by hazelnut shell.Water Res.34, 2955–2962.Demirbas,E.,2003.Adsorption of cobalt(II)from aqueous solution onto activated carbon prepared from hazelnut shells.Adsorpt.Sci.Technol.21,951–963.Ferro-Garcia,M.A.,Rivera-Utrilla,J.,Rodriguez-Gordillo,J., Bautista-Toledo,I.,1988.Adsorption of zinc,cadmium and copperon activated carbons obtained from agricultural by-products.Carbon26,363–373.Gregg,S.J.,Sing,K.S.W.,1982.Adsorption Area and Porosity,second ed.Academic Press Inc.,New York,USA,p.303.Kadirvelu,K.,Namasivayam, C.,2003.Activated carbon from coconut coirpith as metal adsorbent:adsorption of Cd(II)from aqueous solution.Adv.Environ.Res.7,471–478.Kim,J.W.,Myoung,H.S.,Dong,S.K.,Seung,M.S.,Young,S.K., 2001.Production of granular activated carbon from waste walnut shell and its adsorption characteristics for Cu2+ion.J.Hazard.Mater.B85,301–315.Kobya,M.,2004.Adsorption,kinetic and equilibrium studies of Cr(VI)by hazelnut shell activated carbon.Adsorpt.Sci.Technol.22,51–64.Netzer,A.,Hughes,D.E.,1984.Adsorption of Cr,Pb and Co by activated carbon.Water Res.18,927–933.Paajanen,A.,Lehto,J.,Santapakka,T.,Morneau,J.P.,1997.Sorption of cobalt on activated carbons from aqueous solution.Sep.Sci.Technol.32,813–826.Patterson,J.,Passino,R.,1987.Metal Speciation,Separation,and Recovery.Lewis Publishers,Inc.,Chelsea,MI,USA.Petrov,H.,Budinova,T.,Khovesov,I.,1992.Adsorption of zinc, cadmium,copper and lead ions on oxidised anthracite.Carbon30, 135.Philip,C.A.,Girgis,B.S.,1996.Adsorption characteristics of micro-porous carbons from apricot stones activated by phosphoric acid.J.Chem.Technol.Biotechnol.67,248–254.Selomulya,C.,Meeyoo,V.,Amal,R.,1999.Mechanisms of Cr(VI) removal from water by various types of activated carbons.J.Chemical Technol.Biotechnol.74,111–122.Yavuz,O.,Altunkaynak,Y.,Guzel,F.,2003.Removal of copper, nickel,cobalt and manganese from aqueous solution by kaolinite.Water Res.37,948–952.M.Kobya et al./Bioresource Technology96(2005)1518–15211521。

科学家成功解析硫氧镁水泥新物相结构

第06期
膨胀珍珠岩吸附Cu2+和Cr3+的研究
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Adsorption of Cu2+ and Cr3+ using expanded perlite

除去水中的砷

Journal of Hazardous Materials 182 (2010) 156–161Contents lists available at ScienceDirectJournal of HazardousMaterialsj o u r n a l h o m e p a g e :w w w.e l s e v i e r.c o m /l o c a t e /j h a z m atAs(III)removal using an iron-impregnated chitosan sorbentDaniel Dianchen Gang a ,∗,Baolin Deng b ,LianShin Lin caDepartment of Civil Engineering,University of Louisiana at Lafayette,Lafayette,LA 70504,USAbDepartment of Civil and Environmental Engineering,University of Missouri,Columbia,MO 65211,USA cDepartment of Civil and Environmental Engineering,West Virginia University,Morgantown,WV 26506,USAa r t i c l e i n f o Article history:Received 18December 2009Received in revised form 28May 2010Accepted 1June 2010Available online 9 June 2010Keywords:Trivalent arsenic Iron-chitosan AdsorptionAs(III)adsorption kinetics Adsorption isotherma b s t r a c tAn iron-impregnated chitosan granular adsorbent was newly developed to evaluate its ability to remove arsenic from water.Since most existing arsenic removal technologies are effective in removing As(V)(arsenate),this study focused on As(III).The adsorption behavior of As(III)onto the iron-impregnated chi-tosan absorbent was examined by conducting batch and column studies.Maximum adsorption capacity reached 6.48mg g −1at pH =8with initial As(III)concentration of 1007␮g L −1.The adsorption isotherm data ﬁt well with the Freundlich model.Seven hundred and sixty eight (768)empty bed volumes (EBV)of 308␮g L −1of As(III)solution were treated in column experiments.These are higher than the empty bed volumes (EBV)treated using iron-chitosan composites as reported by previous researchers.The investi-gation has indicated that the iron-impregnated chitosan is a very promising material for As(III)removal from water.© 2010 Elsevier B.V. All rights reserved.1.IntroductionArsenic,resulting from industrial and mine waste discharges or from natural erosion of arsenic containing rocks,is found in many surface and ground waters [1].Common chemical forms of arsenic in the environment include arsenate (As(V)),arsenite (As(III)),dimethylarsinic acid (DMA),and monomethylarsenic acid (MMA).Inorganic forms of arsenic (As(V)and As(III))are more toxic than the organic forms [2].Arsenite can be predominant in ground-water with low oxygen levels and is generally more difﬁcult to be removed than arsenate [3].Due to the negative impacts of arsenic on human health that range from acute lethality to chronic and car-cinogenic effects,the U.S.Environmental Protection Agency revised the maximum contaminant level (MCL)of arsenic in drinking water from 50to 10␮g L −1[4].This new regulation has posed a chal-lenge for the research of new technologies capable of selectively removing low levels of arsenic.Existing technologies that are being used for arsenic removal include precipitation [5],membrane separation,ion exchange,and adsorption [6–9].While these approaches can remove arsenic to below 10␮g L −1under optimal conditions,most of the systems are expensive,not suitable for small communities with limited resources.Of these methods,much work has been done on arsenic removal through adsorption because it is one of the most effec-∗Corresponding author.Tel.:+13374825184;fax:+13374826688.E-mail addresses:ddgang@ ,digang@ ,Gang@ (D.D.Gang).tive and inexpensive methods for arsenic treatment [7].Therefore,development of highly effective adsorbents is a key for adsorption-based technologies.Several iron(III)oxides,such as amorphous hydrous ferric oxide [5]and crystalline hydrous ferric oxide [10]are well known for their ability to remove both As(V)and As(III)from aqueous solutions.In general,arsenate is more readily removed by ferric (hydr)oxides than arsenite [11].Reported mechanisms for arsenic removal include adsorption onto the hydroxide surfaces,entrapment of adsorbed arsenic in the ﬂocculants,and formation of complexes and ferric arsenate (FeAsO 4)[12].The presence of other anions such as sulfate,chloride,and in particular,silicates,phosphate,and natural organic matters,can signiﬁcantly affect arsenic adsorption [13–15].The use of iron (hydr)oxides in ﬁne powdered or amor-phous forms was found to be effective for arsenic removal,but the process requires follow-up solid/water separation.For packed-bed adsorption systems,high-efﬁcient granular forms of adsorbent are essential.Recently,several iron based granular materials and processes have been developed for arsenic removal.Dong et al.[16]devel-oped iron coated pottery granules (ICPG)for both As(III)and As(V)removal from drinking water.The column tests showed that ICPG consistently removed total arsenic from test water to below 5␮g L −1level.In another study,Gu et al.[17]used iron-containing granular activated carbon for arsenic adsorption.This iron-containing granular activated carbon was shown to remove arsenic most efﬁciently when the iron content was approximately 6%.Viraraghavan et al.[18]reported a green sand ﬁltration process and found a strong correlation between inﬂuent Fe(II)concen-0304-3894/$–see front matter © 2010 Elsevier B.V. All rights reserved.doi:10.1016/j.jhazmat.2010.06.008D.D.Gang et al./Journal of Hazardous Materials182 (2010) 156–161157tration and arsenic removal percentage.The removal percentage increased from41%to above80%as the ratio of Fe/As was increased from0to20.Granular ferric hydroxide(GFH),another iron based granular material,showed a high treatment capacity for arsenic removal in a column setting before the breakthrough concentration reached10␮g L−1[19].It was found that complexes were formed upon the adsorption of arsenate on GFH[20].Selvin et al.[21]con-ducted laboratory-scale tests over50different media for arsenic removal and found GFH with a particle size of0.8–2.0mm was the most effective one among the tested media.However,some disad-vantages with GFH exist,including quick head loss buildup within 2days because of theﬁne particle size,and signiﬁcant reduction (50%)in adsorption capacity with larger sized media(1.0–2.0mm).Chitin and its deacetylated product,chitosan,are the world’s second most abundant natural polymers after cellulose.These polymers contain primary amino groups,which are useful for chemical modiﬁcations and can be used as potential separa-tors in water treatment and other industrial applications.Many researchers focused on chitosan as an adsorbent because of its non-toxicity,chelating ability with metals,and biodegradability[22]. Several studies have demonstrated that chitosan and its deriva-tives could be used to remove arsenic from aqueous solutions [23,24].Based on the fact that both iron(III)oxides and chitosan exhib-ited high afﬁnity for arsenic,this study focused on examining the effectiveness of an iron-impregnated chitosan granular adsorbent for arsenic removal.Most arsenic removal technologies are more effective for removing arsenate than for arsenite[12].We found in this study that the iron-impregnated chitosan was effective for arsenite removal from experiments in both batch and column set-tings.2.Experimental2.1.Preparation of iron-chitosan beadsThe experimental procedure for the preparation of iron-chitosan beads was described in detail by Vasireddy[25].To summarize, approximately10g of medium molecular weight chitosan(Aldrich Chemical Corporation,Wisconsin,USA)was added to0.5L of0.01N Fe(NO3)3·9H2O solution under continuous stirring at60◦C for2h to form a viscous gel.The beads were formed by drop-wise addition of chitosan gel into a0.5M NaOH precipitation bath under room temperature.Maintaining this concentration of NaOH was critical for forming spherically shaped beads[25].The beads were then separated from the0.5M NaOH solution and washed several times with deionized water to a neutral pH.The wet beads were then dried in an oven under vacuum and in air.Theﬁnal iron content of the chitosan bead was about8.4%.2.2.Arsenic measurementAn atomic absorption spectrometer(AAS)(Thermo Electron Corporation)equipped with an arsenic hollow cathode lamp was employed to measure arsenic concentration.An automatic inter-mittent hydride generation device was used to convert arsenic in water samples to arsenic hydride.The hydrides were then purged continuously by argon gas into the atomizer of an atomic absorption spectrometer for concentration measurements.As(III)stock solution(1000mg L−1)was prepared by dissolving 1.32g of As2O3(obtained from J.T.Baker)in distilled water con-taining4g NaOH,which was then neutralized to pH about7with 1%HCl and diluted to1L with distilled water.All the working solu-tions were prepared with standard stock solution.To50mL of each sample solution(i.e.,reagent blank,standard solutions,and water samples),5mL1%HCl and5mL of100g L−1NaI solution were used to convert arsenic in water samples to arsenic hydride.2.3.Arsenic adsorption experimentsEach arsenic solution(100mL)of desired concentration was mixed with the iron-chitosan beads in a250mL conicalﬂask.The solution pH was adjusted with0.1M HCl or0.1M NaOH to obtain the desired pHs.A pH buffer was not used to avoid potential com-petition of buffer with As(III)sorption.One sample of the same concentration solution without adsorbent(blank),used to estab-lish the initial concentration of the samples,was also treated under same conditions as the samples containing the adsorbent.The solu-tions were placed in a shaker for aﬁxed amount time,followed by ﬁltration to remove the adsorbent.Theﬁltrate was then analyzed for theﬁnal concentration of arsenic using the atomic absorption spectrometer.The solid phase concentration was calculated using the following formula:q=(C i−C f)VM(1) where,q(␮g g−1)is the solid phase concentration,C i(␮g L−1)is the initial concentration of arsenic in solution,C f(␮g L−1)is theﬁnal concentration of arsenic in treated solution;V(L)is the volume of the solution,and M(g)is the weight of the iron-chitosan adsorbent.2.4.Kinetic experimentsAdsorption kinetics was examined with various initial concen-trations at25◦C.The pH of the solutions was chosen at8.0for optimal adsorption.The adsorbent loading for three different ini-tial concentrations of306,584,and994␮g L−1was all0.2g L−1.A predetermined quantity of iron-chitosan adsorbent(20mg)was placed in separate conicalﬂasks with pH-adjusted As(III)solution. The conicalﬂasks were covered with paraﬁlm and placed in a shaker (150rpm),and sub-samples of the solutions were then removed periodically andﬁltered prior to arsenic analysis.To determine the reaction rate constants of arsenic adsorption onto iron-chitosan,both the pseudo-ﬁrst-order and pseudo-second-order models were used.Kinetics of the pseudo-ﬁrst-order model can be expressed as[26]:ln(q e−q t)=ln q e−k1t(2) where,k1(min−1)is the rate constant of pseudo-ﬁrst-order adsorp-tion,q t(mg g−1)is the amount of As(III)adsorbed at time t(min), and q e(mg g−1)is the amount of adsorption at equilibrium.The model parameters k1and q e can be estimated from the slope and intercept of the plot of ln(q e−q t)vs t.The pseudo-second-order model can be expressed as follow[27]:tq t=tq e+1k2q2e(3)where,k2(g mg−1min−1)is the pseudo-second-order reaction rate. Parameters k2and q e can be estimated from the intercept and slope of the plot of(t/q t)vs t.2.5.Isotherm modelsAdsorption isotherms such as the Freundlich or Langmuir mod-els are commonly utilized to describe adsorption equilibrium.The Freundlich isotherm model is represented mathematically as:q e=k f C1/ne(4) where,q e(mg g−1)is the amount of As(III)adsorbed,C e(␮g L−1) is the concentration of arsenite in solution(␮g L−1),k f and1/n158 D.D.Gang et al./Journal of Hazardous Materials182 (2010) 156–161Fig.1.Scanning electron micrograph(SEM)of iron-chitosan bead.are parameters of the Freundlich isotherm,denoting a distribu-tion coefﬁcient(L g−1)and intensity of adsorption,respectively.The Langmuir equation is another widely used equilibrium adsorption model.It has the advantage of providing a maximum adsorption capacity q max(mg g−1)that can be correlated to adsorption proper-ties.The Langmuir model can be represented as:q e=q maxK L C e1+K L C e(5)where,q max(mg g−1)and K L(L mg−1)are Langmuir constants representing maximum adsorption capacity and binding energy, respectively.2.6.Column studyColumn study was conducted to investigate the use of iron-chitosan as a low-cost treatment technology for arsenite removal. Experiments were conducted with a12-mm-ID glass column packed with1.5g iron-chitosan as aﬁxed bed.The inﬂuent solu-tion had an inlet As(III)concentration of308␮g L−1at pH8,and was passed the column at aﬂow rate of25mL h−1.Efﬂuent solu-tion samples were collected and analyzed for arsenic concentration during the column test.3.Results and discussion3.1.Structure characterization of iron-chitosan beadsThe prepared iron-chitosan beads were examined by scanning electron microscope(SEM)(AMRAY1600)for the surface morphol-ogy.A working distance of5–10mm,spot size of2–3,secondary electron(SE)mode,and accelerating voltage of20keV were used to view the samples.It can be seen from Fig.1that the beads are porous in structure.X-ray Photoelectron Spectroscopy(XPS),a sur-face sensitive analytic tool to determine the surface composition and electronic state of a sample,was used in this study.In XPS analysis,a survey scan was used to determine the elements exist-ing on the surface.The high resolution utility scans were then used to measure the atomic concentrations of Fe,C,N and O in the sam-ple.Fig.2shows the peak positions of carbon,nitrogen,oxygen,and iron obtained by the XPS for iron-chitosan beads.In Fig.2,the car-bon1s peak was observed at283.0eV with a FWHM(full width at maximum height)of2.015.The Fe peak was observed at730.0eV. The N-1s peak for iron-chitosan bead was found at398.0eV(FWHM 2.00eV),which can be attributed to the amino groups inchitosan.Fig.2.XPS spectrum of iron-chitosan bead.3.2.Effect of pHThe effect of pH on arsenite removal with the iron-chitosan adsorbent was examined using100mL As(III)solution with an initial concentration of314␮g L−1and a solid loading rate of 0.15g L−1.The solution pH was adjusted with0.1M HCl or0.1M NaOH to obtain pHs ranging from4to12.Lower pHs were avoided because the acid environments could lead to partial dissolution of the chitosan polymer and make the beads unstable[25,28]. The solutions were placed in a shaker(150rpm)for20h at room temperature(25◦C),followed byﬁltration to remove the adsor-bent.The amounts of As(III)adsorbed,calculated using Eq.(1),are present in Fig.3.Under the experimental conditions,approximately 2.0mg g−1of As(III)was adsorbed and that amount did not change signiﬁcantly in the pH range4–9.However,when pH was higher than9.2,arsenite removal decreased dramatically with increasing pH.The results can be explained using arsenic chemical speciation in different pH ranges[29].Arsenite remains mostly as a neutral molecule for pH<9.2,and negatively charged at pH>9.2.So at pH>9.2,arsenite sorption is less because of the unfavorable electro-static interaction with negatively charged surfaces.This adsorptive behavior is common for arsenite with other adsorbents[17,30].Gu et al.[17]reported that pH had no obvious effect on As(III)removal in the range of4.4–9.0,with removal efﬁciency above95%.Another study indicated that the uptake of As(III)by fresh andimmobi-Fig.3.Arsenite removal of the iron-chitosan adsorbent(0.15g L−1)as a function of pH for initial arsenite concentration of314␮g L−1at T=25◦C.D.D.Gang et al./Journal of Hazardous Materials182 (2010) 156–161159Fig.4.Adsorption kinetics for different initial arsenite concentrations with iron-chitosan adsorbent loading of0.2g L−1at pH=8and T=25◦C.lized biomass was not greatly affected by solution pH with optimal biosorption occurring at around pH6–8[30].Raven et al.[11] reported that a maximum adsorption of arsenite on ferrihydrite was observed at approximately pH9.3.3.Kinetics of adsorptionFig.4illustrates the adsorption kinetics for three different ini-tial arsenite concentrations.More than60%of the arsenite was adsorbed by iron-chitosan within theﬁrst30min,then adsorption leveled off after2h.Given the initial concentrations and adsorbent loading,equilibrium was reached after about2h.The adsorption capacity increased from1.51to4.60mg g−1as the initial arsen-ite concentration was increased from306to994␮g L−1.The rapid adsorption in the beginning can be attributed to the greater con-centration gradient and more available sites for adsorption.This is a common behavior with adsorption processes and has been reported in other studies[31].The sorption rate of As(III)on nat-urally available red soil was initially rapid in theﬁrst2h and slowed down thereafter[32].Elkhatib et al.[33]reported that the initial adsorption was rapid,with more than50%of As(III) adsorbed during theﬁrst0.5h in an arsenite adsorption study. Fuller et al.[34]reported that As(V)adsorption onto synthesized ferrihydrite had a rapid initial phase(<5min)and adsorption con-tinued for182h.Raven et al.[11]studied the kinetics of As(V) and As(III)adsorption on ferrihydrite and found that most of the adsorption occurred within theﬁrst2h.It has been reported that arsenite forms both inner-and outer-sphere surface complexes on amorphous Fe oxide[35].Another possible adsorption mech-anism is hydrogen bond formation between As(III)and chitosan bead[24].Figs.5and6illustrate modelﬁts of the kinetic data for the pseudo-ﬁrst-order and pseudo-second-order kinetic models. In general,the pseudo-second-order characterized the kinetic data better than the pseudo-ﬁrst-order model.Table1summa-Fig.5.Adsorption kinetics of the iron-chitosan adsorbent(0.2g L−1)for three initial arsenite concentrations at pH=8and T=25◦C,and corresponding pseudo-ﬁrst-ordermodels.Fig.6.Adsorption kinetics of the iron-chitosan adsorbent(0.2g L−1)for three initial arsenite concentrations at pH=8and T=25◦C,and corresponding pseudo-second-order models.rizes adsorption capacities determined from the modelﬁts.It is noted that the second order rate constant(k2)decreased from 3.19×10−2to 1.15×10−2g mg−1min−1as the initial concen-tration increased from306to994␮g L−1.The initial rate(k2q2e) increased from8.48×10−2to27.97×10−2with increasing initial As(III)concentration.Because as initial concentration increased,the concentration difference between the adsorbent surface and bulk solution increased.Jimenez-Cedillo et al.[36]investigated arsenic adsorp-tion kinetics on iron,manganese and iron-manganese-modiﬁed clinoptilolite-rich tuffs and concluded that the adsorption pro-cesses could be described by the pseudo-second-order model.Table1Adsorption capacities and parameter values of kinetic models for three initial arsenite concentrations and iron-chitosan loading of0.2g L−1at pH=8.Initial conc.(␮g L−1)Pseudo-ﬁrst order Pseudo-second orderk1×102(min−1)R2q e,exp(mg g−1)q e,col(mg g−1)k1×102(g mg−1min−1)R2q e,exp(mg g−1)q e,col(mg g−1)k2q2e×102306 2.630.98 1.51 1.24 3.190.99 1.51 1.638.48584 2.380.96 2.90 2.30 1.310.99 2.90 3.1913.28994 2.370.93 4.60 3.26 1.150.99 4.60 4.9327.97160 D.D.Gang et al./Journal of Hazardous Materials182 (2010) 156–161Fig.7.Adsorption isotherms of the iron-chitosan adsorbent (0.2g L −1)for three initial arsenite concentrations at pH =8,and corresponding isotherm models.Thirunavukkarasu et al.[37]examined As(III)adsorption kinet-ics with granular ferric hydroxide (GFH)and found that most of As(III)adsorption onto GFH occurred at pH 7.6,with 68%of As(III)removed within 1h and 97%removed at the equilibrium time of 6h.Kinetic data ﬁtted the pseudo-second-order kinetic model well with a kinetic rate constant of 0.003g GFH h −1␮g −1As,which is equivalent to 5.0×10−2g mg −1min −1[37].In our study,the kinetic rate constants were from 3.19×10−2to 1.15×10−2g mg −1min −1,which were smaller than using GFH.This could be attributed to the differences in adsorbent parti-cle size and initial arsenic concentrations between these two studies.3.4.Adsorption isothermsFig.7presents the adsorption isotherm data and two isotherm models at pH 8.The maximum adsorption capacity was found to increase from 1.97to 6.48mg g −1as the initial concentration of As(III)increased from 295to 1007␮g L −1.Maximum adsorp-tion capacity reached 6.48mg g −1with initial As(III)concentration of 1007␮g L −1.Chen and Chung [24]reported that the adsorp-tion capacity of As(III)was 1.83mg As g −1for pure chitosan bead.This study conﬁrmed that impregnating iron into chitosan could signiﬁcantly increase the As(III)adsorption capacity of the chi-tosan bead.In another study,Driehaus et al.[19]reported that the adsorption capacity could reach 8.5mg As g −1of granular fer-ric hydroxide (GFH).Model parameters and regression coefﬁcients are listed in Table 2.The Freundlich model agreed better with the experimental data compared to the Langmuir model.The adsorp-tion intensity (1/n )and the distribution coefﬁcient (k f )increased as the initial arsenite concentration increased.This indicated the dependence of adsorption on initial concentration.Low 1/n values (<1)of the Freundlich isotherm suggested that any large change in the equilibrium concentration of arsenic would not result in a signiﬁcant change in the amount of arsenic adsorbed.Selim and Zhang [38]reported that adsorption isotherms of three differ-ent soils for As(V)were better ﬁt to the Freundlich modelandFig.8.Breakthrough curve for an inlet arsenite concentration of 308␮g L −1at pH =8for a column reactor packed with the iron-chitosan adsorbent.adsorption intensity values ranged from 0.270to 0.340.Salim and Munekage [39]found that adsorptions of As(III)onto silica ceramic were well ﬁt by the Freundlich isotherm.Similarly low 1/n values for As(V)adsorption have been reported by others [40].3.5.Column studyFig.8shows a breakthrough curve for an inlet arsenite con-centration of 308␮g L −1at pH 8.The break point was observed after 768empty bed volumes (EBV)and adsorbent was exhausted at 1400bed volumes.In comparison,Boddu et al.[23]reported that the break through point was about 40and 120EBV for As(III)and As(V),respectively using chitosan-coated biosorbent.Gupta et al.[41]conducted column tests using iron-chitosan compos-ites for removal of As(III)and As(V)from arsenic contaminated real life groundwater.Their result showed that the iron-chitosan ﬂakes (ICF)could treat 147EBV of As(III)and 112EBV of As(V)spiked groundwater with an As(III)or As(V)concentration of 0.5mg L −1.Given the difference of the initial concentrations between the two studies,the numbers of EBV were lower than what we found in this study.This can be partially attributed to the difference of the water constituents in the real grounder water used in the previous study [41].Gu et al.[17]examined the arsenic breakthrough behaviors for an As-GAC sample prepared from Dacro 20×40LI with an inlet concentration of 56.1␮g L −1As(III).Their results demonstrated that the adsorbent could effectively remove arsenic from ground-water in a column setting.Dong et al.[16]also reported that average removal efﬁciencies for total arsenic,As(III),and As(V)for a 2-week test period were 98%,97%,and 99%,respectively,at an average ﬂow rate of 4.1L h −1and Empty Bed Contact Time (EBCT)>3min.Table 2Values of the Freundlich and Langmuir isotherm model parameters for three arsenite concentrations with iron-chitosan loading of 0.2g L −1at pH 8.Initial concentration(␮g L −1)Freundlich parameters Longmuir constants k f (L g −1)1/n R 2q max (mg g −1)K L (L mg −1)R 22950.590.240.98 2.000.120.985960.640.260.95 2.820.070.9410070.740.330.996.820.010.95D.D.Gang et al./Journal of Hazardous Materials182 (2010) 156–1611614.ConclusionsOverall,the study has demonstrated that iron-impregnated chi-tosan can effectively remove As(III)from aqueous solutions under a wide range of experimental conditions and removal efﬁciency depends on various factors including pH,adsorption time,adsor-bent loading,and initial concentration of As(III)in the solution. Results from the kinetic batch experiments indicated that more than60%of the arsenic was adsorbed by the iron-chitosan within 30min of adsorption.Kinetic resultsﬁt the pseudo-second-order model well.The second order reaction rate constants were found to decrease from3.19×10−2to1.15×10−2g mg−1min−1as the initial As(III)concentration increased from306to994␮g L−1.Adsorp-tion isotherm results indicated that maximum adsorption capacity increased from1.97to6.48mg g−1at pH=8as the initial concen-tration of As(III)increased from0.3to1mg L−1.The adsorption isotherm dataﬁt well to the Freundlich model.Column experi-ments of As(III)removal were conducted using12-mm-ID column at aﬂow rate of25mL h−1with an initial As(III)concentration of 308␮g L−1.This study corroborates that impregnating iron into chitosan can signiﬁcantly increase As(III)adsorption capacity of the chitosan bead.Advantages of using the iron-impregnated chitosan include its high efﬁciency for As(III)treatment and low cost compared with the pure chitosan bead.We expect that the iron-impregnated chi-tosan is a useful adsorbent for As(III)and could be used both in conventional packed-bedﬁltration tower and Point of Use(POU) systems.The possible concerns include the physicochemical sta-bility of the adsorbent because of the biodegradable nature of the chitosan material.Further research is underway to examine the adsorbent stability and whether the iron-impregnated chitosan can maintain its capability after several regeneration andCompeting adsorption of other ions will also be AcknowledgmentsThe authors would like to thank Mr.Ravi K.Kadari and Ms. Dhanarekha Vasireddy for conducting the laboratory experiments. 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天然磁黄铁矿机械活化强化除砷技术

第50卷第11期2019年11月中南大学学报(自然科学版)Journal of Central South University(Science and Technology)V ol.50No.11Nov.2019天然磁黄铁矿机械活化强化除砷技术张荥斐1,2，刘屾淼1,2，曹学锋1,2，韩海生1,2，孙伟1,2，成朋飞1,2，王丽1,2(1.中南大学资源加工与生物工程学院，湖南长沙，410083；2.中南大学战略含钙矿物资源清洁高效利用湖南省重点实验室，湖南长沙，410083)摘要：以天然磁黄铁矿为净化剂，研究As(III)在磁黄铁矿表面的吸附、转化行为；基于机械活化原理，开发高质量浓度含砷废水净化除砷技术。

研究结果表明：As(III)吸附在磁黄铁矿表面，并与表面S2−发生作用生成As2S3，矿物表面同时发生氧化反应生成少量As2S5，As2O3和As2O5，表面含砷组分及单质硫在一定程度上阻碍磁黄铁矿与溶液中砷的进一步反应。

通过对磁黄铁矿进行机械活化处理，一方面，可减小磁黄铁矿粒度、增大比表面积，另一方面，暴露具有较高活性的新鲜表面，极大地提高磁黄铁矿的除砷效率；当pH为3.0~3.5、温度为80℃、反应时间为24h、磁黄铁矿初始质量浓度为12g/L时，砷的去除率大于97%。

关键词：磁黄铁矿；As(III)；水处理；机械活化中图分类号：TD923文献标志码：A文章编号：1672-7207（2019）11-2623-10Arsenic removal technology based on mechanical activation ofnatural pyrrhotiteZHANG Xingfei1,2,LIU Shenmiao1,2,CAO Xuefeng1,2,HAN Haisheng1,2,SUN Wei1,2,CHENG Pengfei1,2,WANG Li1,2(1.School of Minerals Processing and Bioengineering,Central South University,Changsha410083,China;2.Key Laboratory of Hunan Province for Clean and Efficient Utilization of Strategic Calcium-containingMineral Resources,Central South University,Changsha410083,China)Abstract:Natural pyrrhotite was used as a purifying agent to study the adsorption and transformation behavior of As(III)on the surface of pyrrhotite.Based on the mechanical activation,the arsenic removal technology of acid wastewater with high mass concentration of arsenic was developed.The results show that As(III)adsorbs on thesurface of pyrrhotite and reacts with surface S2−to form As2S3.The surface of the mineral is simultaneouslyoxidized to form a small amount of As2S5,As2O3and As2O5.The presence of arsenic-containing components andelemental sulfur on the surface hinder the further reaction of pyrrhotite with arsenic in the solution to some extent.On the one hand,activated treatment of pyrrhotite by mechanical activation reduces the size of pyrrhotite and increases the specific surface area.On the other hand,it exposes the fresh surface with higher activity,and greatly improves the arsenic removal efficiency of pyrrhotite.The removal rate of arsenic is over98%when pH is3.0−3.5收稿日期：2019−01−26；修回日期：2019−03−26基金项目(Foundation item)：中南大学中央高校基本科研业务费专项资金资助项目(2019zzts702)；国家自然科学基金资助项目(51634009,51804340)(Project(2019zzts702)supported by the Fundamental Research Funds for the Central Universities of Central South University;Projects(51634009,51804340)supported by the National Natural Science Foundation of China)通信作者：韩海生，博士，副教授，从事湿法冶金及二次资源回收利用研究；E-mail:*************************DOI:10.11817/j.issn.1672-7207.2019.11.001第50卷中南大学学报(自然科学版)and temperature is80℃,with reaction time being24h,and an initial dosage of pyrrhotite being12g/L.Key words:pyrrhotite;As(III);water treatment;mechanically activated砷(As)是一种有毒非金属元素，在自然界中主要以硫化物的形式存在，在还原环境下，微量的砷可转入地下水中，经过迁移和富集可使地下水中的砷含量增加，长期饮用将会严重影响人的身心健康[1]。

Adsorption_of_nitrobenzene_from_aqueous_solution_by_MCM-41

Journal of Colloid and Interface Science315(2007)80–86/locate/jcisAdsorption of nitrobenzene from aqueous solution by MCM-41Qingdong Qin,Jun Ma∗,Ke LiuSchool of Municipal and Environmental Engineering,Harbin Institute of Technology,P.O.Box2627,202Haihe Road,Harbin150090,People’s Republic of ChinaReceived24April2007;accepted24June2007Available online30June2007AbstractAdsorption of nitrobenzene onto mesoporous molecular sieves(MCM-41)from aqueous solution has been investigated systematically using batch experiments in this study.Results indicate that nitrobenzene adsorption is initially rapid and the adsorption process reaches a steady state after1min.The adsorption isotherms are well described by the Langmuir and the Freundlich models,the former being found to provide the better ﬁt with the experimental data.The effects of temperature,pH,ionic strength,humic acid,and the presence of solvent on adsorption processes are also examined.According to the experimental results,the amount of nitrobenzene adsorbed decreases with an increase of temperature from278 to308K,pH from1.0to11.0,and ionic strength from0.001to0.1mol/L.However,the amount of nitrobenzene adsorbed onto MCM-41does not show notable difference in the presence of humic acid.The presence of organic solvent results in a decrease in nitrobenzene adsorption.The desorption process shows a reversibility of nitrobenzene adsorption onto MCM-41.Thermodynamic parameters such as Gibbs free energy are calculated from the experimental data at different temperatures.Based on the results,it suggests that the adsorption is primarily brought about by hydrophobic interaction between nitrobenzene and MCM-41surface.©2007Elsevier Inc.All rights reserved.Keywords:Nitrobenzene;Adsorption;MCM-41;Isotherm;Desorption1.IntroductionNitrobenzene has been widely adopted in the manufacture of dyes,plastic,pesticides,explosives,pharmaceuticals,and inter-mediates in chemical synthesis industries for years.After use, nitrobenzene in solution is generally discharged to wastewater treatment plants where a larger proportion of it cannot be re-moved andﬁnally is discharged into the surrounding aquatic en-vironment,which tends to persist in the environment and poses a potential toxic threat to ecological and human health[1,2]. Therefore,a variety of possible treatment technologies such as adsorption,ozonation,and advanced oxidation processes have been taken into account for the puriﬁcation of nitrobenzene-contaminated water[3–6].Particularly,adsorption with granu-lar activated carbon(GAC)materials is considered to be one of the most economical and efﬁcient methods for controlling the nitrobenzene in water[7].However,GAC is relativelyﬂam-*Corresponding author.Fax:+8645182368074.E-mail address:majun@(J.Ma).mable and difﬁcult to regenerate.Thus,an alternative adsorbent for nitrobenzene removal is required.Adsorption of organic compounds onto siliceous materials has been investigated intensively by many researchers[8–11]. Siliceous materials such as zeolite,silica,and clay are found widespread in the environment and they are environment-friendly materials for environmental contaminant remediation. In order to enhance the adsorption capacity,modiﬁcations of siliceous materials by cationic surfactants have been made in previous studies[12,13].The modiﬁed materials exhibit ex-cellent adsorption properties for organic pollutants like per-chloroethylene,atrazine,lindane,and diazinone.On the other hand,the synthesis of ordered mesoporous silica with high sur-face area and high hydrophobicity has attracted much attention in several areas because of its possible industrial application as reaction catalyst,catalyst support,chemical sensor,adsorbent for environmentally hazardous chemicals,and electrical and op-tical devices[14,15].MCM-41,one member of the mesoporous molecular sieve M41S family,possesses a high speciﬁc sur-face and a regular hexagonal array of cylindrical pores,which is largely used in shape-selective catalysis,selective adsorp-0021-9797/$–see front matter©2007Elsevier Inc.All rights reserved. doi:10.1016/j.jcis.2007.06.060Q.Qin et al./Journal of Colloid and Interface Science315(2007)80–8681tion and separation processes,and chemical sensors,as well as nanotechnology[15].Such material is characterized by large surface areas,narrow pore size distribution,and moderate hy-drophobic character.Recently,Cooper and Burch reported that M41S possessed large adsorption capacities for efﬁcient elim-ination of both cyanuric acid and p-chlorophenol from aque-ous solution and the adsorption capacity can be regenerated by ozonation[16].Additionally,Wang et al.pointed out that MCM-22could be an effective adsorbent for the removal of methylene blue,crystal violet,and rhodamine B from aqueous solution[17].Nevertheless,there are very scarce reports on the use of M41S for nitrobenzene removal from water and more ef-fort is required in this area in order for such adsorbent to be practically applied.Therefore,the purpose of the present research effort is to investigate the adsorption characteristics of nitrobenzene from aqueous solutions onto MCM-41material.As aﬁrst step in the adsorption of nitrobenzene,it is necessary to quantify the equi-librium time during adsorption.In turn,adsorption isotherms are conducted at batch experiments under different tempera-tures and then thermodynamic parameters are calculated.In-ﬂuencing parameters such as pH,ionic strength,humic acid, and the presence of solvent are evaluated to characterize the extent of nitrobenzene adsorption.Finally,the desorption of nitrobenzene by MCM-41has been studied to determine the reversibility of adsorption.2.Materials and methods2.1.MaterialsNitrobenzene(98%;without further puriﬁcation)was pur-chased from Tianjin Third Chemical Factory.Cetyltrimethy-lammonium bromide(CTMAB)was obtained from Xilong Chemical.Tetramethyl orthosilicate(TMOS)was purchased from Tianjin Kermel Chemical.All solvents used here were pure analytical HPLC grade solvents.Humic acid was pur-chased from Shanghai Chemical.The initial pH of background solution was adjusted by introducing appropriate amounts of acid(HCl)or base(NaOH)solutions.Deionized distilled water was puriﬁed by Millipore Milli-Q system and the ionic strength (I)was adjusted by NaCl solution.Unless otherwise stated,all reagents used in this study were analytical grade.2.2.Preparation of MCM-41The MCM-41was synthesized in an acidic medium using conventional literature recipes[14].Typically,CTMAB was dissolved in HCl solution,into which a TMOS solution was then added dropwise.The molar gel composition was as the following:3.5CTMAB:30SiO2:26.9H+:3800H2O.The white precipitate was homogenized by vigorous stirring for2h at am-bient temperature and then heated at373K in a Teﬂon-coated stainless-steel autoclave for48h under autogenous pressure. Subsequently,the white precipitate wasﬁltered,washed with distilled water,and dried at353K overnight.Finally,the as-synthesized sample was calcined in air in order to remove the surfactant molecules occluded in the pore system of the meso-porous silica.It was heated to813K at a rate of1K/min and kept for5h.The BET surface area of calcined MCM-41was 712m2/g and the pH zpc(pH of zero point of charge)was2.5.2.3.Adsorption experimentsThe effect of contact time on nitrobenzene adsorption onto MCM-41was studied based on the nitrobenzene concentration in the range from2to16µmol/L.The adsorption experiments were carried out by mixing1.0g MCM-41samples with a200-ml aqueous solution in a500-ml stirredﬂask at temperature of 298K.Samples were taken out andﬁltrated by a glassﬁber with0.7-µm pore size at different times.Then the residual con-centration of nitrobenzene was determined.The adsorption isotherm experiments of nitrobenzene onto MCM-41were performed on the basis of a batch experiment.A given amount of adsorbent(0.05g)was placed in a50-ml ﬂask,into which10ml of a nitrobenzene solution with varying concentrations was added.The experiments were performed in a temperature-controlled water bath shaker for4h at a mixing speed of180rpm.After the adsorption reached equilibrium,the solutions wereﬁltered and analyzed for the remaining concen-tration of nitrobenzene.Prior to all experiments,nitrogen was purged into solution to determine the effect of dissolved oxy-gen(DO)on nitrobenzene adsorption,showing that DO had no inﬂuence on the adsorption.Thenﬂowing experiments were undertaken under common conditions.The amount of nitroben-zene adsorbed onto MCM-41was calculated from the mass balance equation as(1) q e=(C0−C e)VM,where q e(µmol/g)is the amount of nitrobenzene adsorbed per gram of MCM-41at equilibrium;C0(µmol/L)and C e (µmol/L)are the initial and equilibrium liquid-phase concen-tration of nitrobenzene,respectively;V(L)is the volume of nitrobenzene solution,and M(g)is the mass of MCM-41used. To check reproducibility,nitrobenzene adsorption was carried out in duplicate.The relative deviations met with the require-ment of less than5%.2.4.Desorption experimentsDesorption isotherms were obtained from the adsorption samples in equilibrium.Solution containing MCM-41was transferred into centrifuge tubes that were centrifuged at4200g for20min with the centrifuge temperature being set at the incu-bation temperature.Then,5ml of the supernatant was removed and compensated for sampling by adding5ml of pure water or 0.1mol/L solution of NaCl.The tubes were placed on a recip-rocating shaker for1h at298K.Preliminary kinetic studies showed that desorption could reach equilibrium state within a time of1min.After shaking,the suspensions were centrifuged and1ml of the supernatant was collected for analysis.82Q.Qin et al./Journal of Colloid and Interface Science315(2007)80–86Fig.1.Effect of contact time on nitrobenzene adsorption by MCM-41(adsorp-tion conditions:adsorbent dosage=5g/L,pH5.8,and temperature=298K).2.5.Analytical methodThe concentration of nitrobenzene was determined by high performance liquid chromatography(Waters)equipped with UV-visible detection at a wavelength of263nm,using a sym-metry C-18column(4.6×150mm,5-µm spheres,Waters).The injection volume was20µl and the mobile phase was a mixture of methanol–water(60:40v/v)with a rate of1ml/min.Un-der these conditions,the retention time for nitrobenzene in the column was5.2min.3.Results and discussion3.1.Effect of contact time on nitrobenzene adsorptionAs is illustrated by Fig.1,an apparent adsorption equilib-rium is generally obtained within1min for the initial nitroben-zene concentration of2,4,and16µmol/L,respectively.There-after,no detectable concentration changes occur(<30min) after adsorption equilibrium(1min)and the average removal efﬁciency of nitrobenzene reaches about72.6,64.5,and56.1% when the initial concentrations of nitrobenzene are2,4,and 16µmol/L,respectively,during this period.It is assumed here that the adsorption takes place primarily at easily accessible surface sites,requiring no diffusion into micropores.And a hydrophobic interaction between adsorbent and organic com-pounds may be attributed to the rapid adsorption rate[18]. Additionally,Fig.1also illustrates that the removal of nitroben-zene at adsorption equilibrium decreases with increasing initial nitrobenzene concentration.The reason is the limited number of adsorption sites available for the uptake of nitrobenzene at a ﬁxed adsorbent dosage.3.2.Adsorption isothermsAdsorption equilibrium data,expressed by the mass of ad-sorbate adsorbed per unit weight of adsorbent and liquidphase Fig.2.Adsorption isotherm of nitrobenzene onto MCM-41at different temper-atures(adsorption conditions:adsorbent dosage=5g/L and pH5.8). equilibrium concentration of adsorbate,are usually represented by adsorption isotherms,which is of importance in the design of adsorption systems.The adsorption isotherms of nitrobenzene onto MCM-41at temperatures of278,288,298,and308K are shown in Fig.2.All adsorption isotherms are nonlinear with curvatures concave to the abscissa.The adsorption of nitroben-zene onto MCM-41decreases from2.018µmol/g(57.7%re-moval)to1.092µmol/g(31.2%removal)when temperature is increased from278to308K at an initial concentration of 17.5µmol/L.The decrease in the equilibrium adsorption of nitrobenzene with temperature demonstrates that nitrobenzene removal by adsorption onto MCM-41favors a low tempera-ture.In order to describe the adsorption isotherm,two important isotherms are selected in this study,the Langmuir and Fre-undlich isotherms,(2) q e=Q0K L C e1+K L C e,(3) q e=K F C1/n e,where q e(µmol/g)is the amount of nitrobenzene adsorbed per gram of MCM-41at equilibrium;C e(µmol/L)the equilib-rium concentration of nitrobenzene in solution;Q0(µmol/g) the maximum monolayer adsorption capacity;K L(L/µmol) the constant related to the free energy of adsorption;K F [µmol/g(L/µmol)1/n],a Freundlich isotherm constant for the system and the slope1/n,ranging between0and1,indicative of the degree of nonlinearity between solution concentration and adsorption.The isotherm parameters and linear regression obtained from theﬁtting curves by Langmuir and Freundlich models are given in Table1.It is clear that adsorption isotherms at different temperatures can beﬁtted well using two isotherm models(evidenced from the correlation coefﬁcients,>0.990). However,the Langmuir model is more suitable than the Fre-undlich model to describe the adsorption isotherm,as reﬂected with correlation coefﬁcients.It suggests that the adsorbent is homogeneous,and the adsorptionﬁlm is monomolecular.Then the following equations use Langmuir constants to calculate theQ.Qin et al./Journal of Colloid and Interface Science315(2007)80–8683 Table1Parameters of adsorption isotherms of nitrobenzene onto MCM-41at differenttemperaturesTempe-rature (K)Langmuir model Freundlich modelQ0(µmol/g)K L(L/µmol)R2K F[µmol/g(L/µmol)1/n]1/n R2278 3.7050.1430.99880.4630.750.9980 288 2.6820.1260.99950.3170.710.9958 298 2.1420.1240.99930.2550.690.9950 308 1.8410.1170.99980.2170.670.9914 thermodynamic parameters.On the other hand,the maximum monolayer adsorption capacity,Q0,deﬁned the total capacity of MCM-41for nitrobenzene adsorption and decreases with increasing temperature.Its maximum value is determined as 3.705µmol/g at278K.A change in the maximum monolayer adsorption capacity with temperature is possibly due to the low hydrothermal stability of MCM-41in water.At high tempera-tures,the structure of MCM-41could collapse by mechanical compression through the hydrolysis of siloxane bonds in the presence of adsorbed water,which may decrease the amount of active sites[15].Moreover,when the interaction is exothermic, the adsorbate has a tendency to escape from the solid phase to the solution.Thus,rise in temperature and excess energy supply will promote desorption.The free energy( G0)of nitrobenzene adsorption onto MCM-41at different temperatures is calculated based on the adsorption isotherm by[19](4) G0=−RT ln K L,where R(J/mol K)is the universal gas law constant;T(K)the absolute temperature of solution,and K L(L/mol)the Langmuir constant.The values of G0are calculated as from−27.4to −29.9kJ/mol.It is noted that the G0values are all nega-tive,which indicate the feasibility and spontaneous adsorption of nitrobenzene onto MCM-41.The enthalpy change( H0) and entropy change( S0)could not be established because a linear relationship of ln K L and1/T in the van’t Hoff equation could not be established.3.3.Effect of pH on nitrobenzene adsorptionThe pH of solution is one of the most important parameters affecting the adsorption process.In order to determine the ef-fect of pH on adsorption capacity of MCM-41,solutions were prepared at different pH values ranging from1.0to11.0.The dependence of pH on the adsorption of nitrobenzene at an ini-tial concentration of17.5µmol/L onto MCM-41is illustrated in Fig.3.Obviously,the amount of adsorbed nitrobenzene(q e) is decreased by increasing the pH value.At pH1.0,the uptake of nitrobenzene is54.3%(1.90µmol/g),while at pH11.0,the uptake of nitrobenzene is only18.1%(0.63µmol/g).It is ap-parent that using solutions at pH value lower than1.0gives the highest q e value.For adsorption onto a solid surface,six adsorption mecha-nisms are believed to exist(i.e.,electrostatic interaction,ionex-Fig.3.Effect of solution pH on the adsorption of nitrobenzene onto MCM-41 (adsorption conditions:adsorbent dosage=5g/L,temperature=298K,initial nitrobenzene concentration=17.5µmol/L,and ionic strength=0.1mol/L). change,ion–dipole interactions,coordination by surface metal cations,hydrogen bonding,and hydrophobic interaction[20]). Since nitrobenzene is an unionizable compound,electrostatic interaction and ion-exchange mechanisms are negligible.The ion–dipole interactions between the charged surface and the nonionic nitrobenzene are also expected to be negligible in this experiment.Coordination by surface metal cations is im-portant only when the organic ligand is a good electron donor relative to water.Therefore,the adsorption mechanisms can be contributed by hydrophobic interaction and hydrogen bonding. Nitrobenzene is an unionizable compound.Thus,the effect of pH on the adsorption of nitrobenzene onto MCM-41may result from a changed MCM-41surface.Based on the conclusions of Lu et al.[20],the maximum extent of nitrobenzene adsorp-tion is at a pH value close to its pH zpc(2.5)if hydrogen bond interaction is the dominant mechanism for nitrobenzene ad-sorption.However,it is noteworthy that the maximum extent of adsorption(Fig.3)is not at a pH value around its pH zpc in this study.Hence,it can be inferred that a speciﬁc interaction like hydrophobic interaction at the interface occurs between nitrobenzene and MCM-41.The possible explanation for the sharp decrease in nitrobenzene adsorption with the increase of pH could be that MCM-41material is relatively stable with high acid resistance,whereas it degrades readily in basic so-lution that will decrease the hydrophobic character and destroy the structure of MCM-41[15].3.4.Effect of ionic strength and humic acid on nitrobenzene adsorptionSince NaCl is often used as a stimulator in adsorption processes,the effect of ionic strength(adjusted by NaCl)on nitrobenzene adsorption has been determined and the results are presented in Fig.4.It indicates that the amount of ad-sorbed nitrobenzene onto MCM-41decreases with the increase of NaCl concentration of the solution.The Langmuir constants84Q.Qin et al./Journal of Colloid and Interface Science315(2007)80–86Fig.4.Effect of ionic strength on the adsorption of nitrobenzene onto MCM-41 (adsorption conditions:adsorbent dosage=5g/L,pH5.8,and temperature= 298K).(K L)at NaCl concentrations of0.001,0.01,and0.1mol/L are 0.128,0.127,and0.116L/µmol,respectively.As nitrobenzene is a nonpolar hydrophobic chemical,increasing ionic strength will lead to a decrease of the solubility of nitrobenzene in water.Therefore,for nitrobenzene,increasing ionic strength should usually be accompanied by an increase in partition into MCM-41.However,the amount of adsorbed nitrobenzene onto MCM-41decreases with the increase of ionic strength.Given the pH zpc value,the surface of MCM-41is expected to be dom-inantly negatively charged at the experimental pH of5.8.It is beneﬁcial for Na+to be adsorbed onto MCM-41by electro-static interaction.It infers that the additive(NaCl)prefers to occupy the strongest sites on MCM-41,and competes with ni-trobenzene for space in MCM-41due to the surface heterogene-ity,which leads to the decrease in the amount of nitrobenzene adsorption.Humic acid exists extensively in surface water and it inter-feres with the adsorption of trace organic compounds on porous adsorbents such as powdered activated carbon(PAC)by pore blockage and direct competition for adsorption sites[21].The effect of humic acid on nitrobenzene adsorption onto MCM-41is shown in Fig.5.The results suggest that the existence of humic acid at a concentration of10and50mg/L does not signiﬁcantly affect the extent of nitrobenzene adsorption.The amount of humic acid adsorbed at concentrations of10and 50mg/L is about0.15and0.26mg/g,respectively.This ob-servation can be explained by the large humic acid molecule. As the molecule of humic acid is larger than the pore diame-ter(4.3nm)of MCM-41,it would not be able to penetrate the pore spaces of MCM-41to form pore blockage and compete for adsorption sites.Hence,it does not reduce the adsorption efﬁciency for nitrobenzene.Similar results were observed by other researchers who examined the adsorption of Geosmin and 2-Methylisoborneol by ultrastable zeolite-Y in the presence of humic acid[22].Fig.5.Effect of humic acid on the adsorption of nitrobenzene onto MCM-41 (adsorption conditions:adsorbent dosage=5g/L,pH3.0,ionic strength= 0.1mol/L,and temperature=298K).Fig.6.Effect of solvent on the adsorption of nitrobenzene onto MCM-41(ad-sorption conditions:adsorbent dosage=5g/L,pH5.8,and temperature= 298K).3.5.Adsorption of nitrobenzene from solvent–water solutionsAdsorption of nitrobenzene from aqueous solutions contain-ing10%methanol or10%acetone was studied.Methanol is a protic solvent with hydrogen bonding character.Acetone,less polar than methanol,is a dipolar aprotic solvent that cannot act as a hydrogen bond donor.The adsorption isotherm from solvent–water solutions is illustrated in Fig.6.A decrease in the adsorption is evident when organic solvent is added in wa-ter solution.The maximum monolayer adsorption capacities (Q0)in water solution,methanol–water solution,and acetone–water solution are2.14,1.88,and0.63µmol/g,respectively.It also indicates that the presence of acetone appears to inhibit ni-trobenzene adsorption to an extent greater than the presence of methanol.Q.Qin et al./Journal of Colloid and Interface Science315(2007)80–8685Fig.7.Adsorption–desorption isotherms of nitrobenzene(adsorption condi-tions:adsorbent dosage=5g/L,pH5.8,and temperature=298K).The addition of organic solvent to water enhanced the inter-action between the nitrobenzene and the solvent,resulting in a decrease in adsorption capacity.This could be attributed to the difference in proton afﬁnity and solvation of these solvents. Thus,the inhibitory effect in less polar solution for nitroben-zene adsorption is greater than that in polar solution due to hydrophobic driving force being the dominant mechanism for nitrobenzene adsorption,where less polar solvent interacts with nitrobenzene more strongly than polar solvent.In other words, increasing nitrobenzene solubility in the mixed solvent results in a reduced driving force for nitrobenzene adsorption.Acetone interacts with nitrobenzene more strongly than methanol.Con-sequently,the hydrophobic driving force could be responsible for the adsorption of nitrobenzene onto MCM-41.3.6.Desorption of nitrobenzene by MCM-41Desorption is one of the key processes affecting the ulti-mate fate of contaminants in solids.In order to evaluate the reversibility of nitrobenzene adsorption onto MCM-41,desorp-tion characteristics were also determined.Desorption isotherms of nitrobenzene adsorbed onto MCM-41are shown in Fig.7. There is no hysteresis in Milli-Q water between adsorption and desorption.It suggests that the adsorption and desorption are reversible and the exchange of nitrobenzene between bulk wa-ter and MCM-41can be realized with a similar mechanism. Thus,desorption could be predicted on the basis of adsorp-tion isotherms.However,since adsorption of nitrobenzene onto MCM-41is ionic strength dependent,a change of suspension ionic strength would,of course,inﬂuence the desorption of adsorbed nitrobenzene.The observed desorption isotherm in 0.1mol/L NaCl is similar to the adsorption isotherm obtained in0.1mol/L NaCl.It seems that after adsorption,there are no strong bonding forces at the interface between nitrobenzene and MCM-41.Hydrophobic interaction may be the signiﬁcant bonding force in adsorption of nitrobenzene onto MCM-41.4.ConclusionsSynthetic MCM-41by a hydrothermal method was demon-strated to adsorb nitrobenzene from aqueous solution.The results showed that nitrobenzene could be adsorbed onto MCM-41quickly.The data obtained from adsorption isotherms at different temperatures wereﬁt to the Langmuir model.The maximum monolayer adsorption capacity for nitrobenzene de-creased from3.705to1.841µmol/g with an increase of tem-perature from278to308K at pH5.8.It was found that increas-ing pH from1.0to11.0decreased the adsorption capacities from54.3to18.1%.The result was attributed to destroying the hydrophobic characteristic of MCM-41in basic solutions.In-creasing ionic strength from0.001to0.1mol/L decreased the extent of nitrobenzene adsorption from2.12to1.81µmol/g for ions mainly occupied the strongest sites on MCM-41and led to a steric hindrance.The presence of humic acid had no inﬂu-ence on nitrobenzene adsorption and the presence of cosolvent (methanol and acetone)affected the adsorption process by a cosolvent effect.Desorption isotherms showed that adsorption and desorption of nitrobenzene were bined with pH effects,it was concluded that hydrophobic interaction was the main force to adsorb nitrobenzene from aqueous solution. Further research is needed to improve the stability of MCM-41 in aqueous solution.AcknowledgmentThe authors are grateful for theﬁnancial support provided by the National Natural Science Foundation of China(NSFC, Project50578051).References[1]S.B.Haderlein,R.P.Schwarzenbach,Environ.Sci.Technol.27(1993)316.[2]X.K.Zhao,G.P.Yang,X.C.Gao,Chemosphere52(2003)917.[3]S.A.Boyd,G.Y.Sheng,B.J.Teppen,C.T.Johnston,Environ.Sci.Tech-nol.35(2001)4227.[4]O.V.Makarova,T.Rajh,M.C.Thurnauer,A.Martin,P.A.Kemme,D.Cropek,Environ.Sci.Technol.34(2000)4797.[5]tifoglu,M.D.Gurol,Water Res.37(2003)1879.[6]F.J.Beltrán,J.M.Encinar,M.A.Alonso,Ind.Eng.Chem.Res.37(1998)25.[7]M.Franz,H.A.Arafat,N.G.Pinto,Carbon38(2000)1807.[8]M.A.Anderson,Environ.Sci.Technol.34(2000)725.[9]H.T.Shu,D.Y.Li,A.A.Scala,Y.H.Ma,Sep.Purif.Technol.11(1997)27.[10]C.Y.Chang,W.T.Tsai,C.H.Ing,C.F.Chang,J.Colloid Interface Sci.260(2003)273.[11]C.F.Chang,C.Y.Chang,K.H.Chen,W.T.Tsai,J.L.Shie,Y.H.Chen,J.Colloid Interface Sci.277(2004)29.[12]Z.H.Li,R.S.Bowman,Environ.Sci.Technol.32(1998)2278.[13]J.Lemi´c,D.Kovaˇc evi´c,M.Tomaševi´c-ˇCanovi´c,D.Kovaˇc evi´c,T.Stani´c,R.Pfend,Water Res.40(2006)1079.[14]X.S.Zhao,G.Q.Lu,lar,Ind.Eng.Chem.Res.35(1996)2075.[15]P.Selvam,S.K.Bhatia,C.G.Sonwane,Ind.Eng.Chem.Res.40(2001)3237.[16]C.Cooper,R.Burch,Water Res.33(1999)3689.[17]S.B.Wang,H.T.Li,L.Y.Xu,J.Colloid Interface Sci.295(2006)71.86Q.Qin et al./Journal of Colloid and Interface Science315(2007)80–86[18]M.Khalid,G.Joly,A.Renaud,P.Magnoux,Ind.Eng.Chem.Res.43(2004)5275.[19]K.P.Singh,D.Mohan,S.Sinha,G.S.Tondon,D.Gosh,Ind.Eng.Chem.Res.42(2003)1965.[20]M.C.Lu,G.D.Roam,J.N.Chen,C.P.Huang,Water Res.30(1996)1670.[21]G.Newcombe,J.Morrison,C.Hepplewhite,D.R.U.Knappe,Carbon40(2002)2147.[22]J.Ellis,W.Korth,Water Res.27(1993)535.。

红外光谱和X-射线衍射表征水铁矿吸附砷酸的性质

Infrared spectroscopic and X-ray diﬀraction characterizationof the nature of adsorbed arsenate on ferrihydriteYongfeng Jiaa,*,Liying Xu a ,Xin Wang a ,George P.Demopoulosb,*aInstitute of Applied Ecology,Key Laboratory of Terrestrial Ecological Process,Chinese Academy of Sciences,Shenyang 110016,ChinabDepartment of Mining,Metals and Materials Engineering,McGill University,Montreal,QC,Canada H3A 2B2Received 30March 2006;accepted in revised form 14December 2006;available online 30January 2007AbstractFourier transformed infrared (FTIR)spectroscopy was used to characterize arsenate–ferrihydrite sorption solids synthe-sized at pH 3–8.The speciation of sorbed arsenate was determined based on the As–O stretching vibration bands located at 650–950cm À1and O–H stretching vibration bands at 3000–3500cm À1.The positions of the As–O and O–H stretching vibra-tion bands changed with pH indicating that the nature of surface arsenate species on ferrihydrite was strongly pH dependent.Sorption density and synthesis media (sulfate vs.nitrate)had no appreciable eﬀect.At acidic pH (3,4),ferric arsenate surface precipitate formed on ferrihydrite and constituted the predominant surface arsenate species.X-ray diﬀraction (XRD)analyses of he sorption solids synthesized at elevated temperature (75°C),pH 3clearly showed the development of crystalline ferric arsenate (i.e.scorodite).In neutral and alkaline media (pH 7,8),arsenate sorbed as a bidentate surface complex (in both pro-tonated B FeO 2As ðO ÞðOH ÞÀand unprotonated B FeO 2As ðO Þ22Àforms).For the sorption systems in slightly acidic media (pH 5,6),both ferric arsenate and surface complex were probably present on ferrihydrite.It was further determined that the incor-porated sulfate in ferrihydrite during synthesis was substituted by arsenate and was more easily exchangeable with increasing pH.Ó2007Elsevier Ltd.All rights reserved.1.INTRODUCTIONFerrihydrite is a poorly ordered hydrous iron oxide commonly present in low-temperature geochemical pro-cesses.It is widely occurring in surface environments,e.g.in soils,lake and river sediments and water columns (Way-chunas et al.,1993;Cornell and Schwertmann,1996;Jam-bor and Dutrizac,1998).Arsenate is an important form of arsenic in both natural water systems and mineral pro-cessing tailings.Ferrihydrite shows strong aﬃnity for arse-nate at mineral–water interfaces.Adsorption on ferrihydrite is an important factor controlling transport,fate and bioavailability of arsenic in soils,groundwaterand surface water systems (Smedley and Kinniburgh,2002).Ferrihydrite is also a common product in many hydrometallurgical operations such as the coprecipitation process of iron with arsenate as well as in acid mine drain-age (Jambor and Dutrizac,1998;Carlson et al.,2002).Hence the adsorption of arsenate on ferrihydrite plays an important role in the removal and immobilization of ar-senic from industrial eﬄuents as well as the fate of arsenate in tailings impoundment.Adsorption of arsenate on ferrihydrite involves ligand exchange with H 2O and/or OH Àon the substrate surface (Jain et al.,1999).Factors inﬂuencing adsorption of arse-nate on ferrihydrite include medium pH,type and con-centration of co-ions,initial Fe/As molar ratios etc.pH is one of the most important factors that control aqueous arsenate speciation and surface functional groups of hy-drous oxides,consequently inﬂuencing the macroscopic characteristics of adsorption process and the microscopic characteristics such as bonding modes and the nature of0016-7037/$-see front matter Ó2007Elsevier Ltd.All rights reserved.doi:10.1016/j.gca.2006.12.021*Corresponding authors.Fax:+862483970436(Y.Jia),+15143984492(G.P.Demopoulos).E-mail addresses:yongfeng.jia@ (Y.Jia),george.demopoulos@mcgill.ca (G.P.Demopoulos)./locate/gcaGeochimica et Cosmochimica Acta 71(2007)1643–1654adsorbed arsenate on the surface of ferrihydrite(Mas-scheleyn et al.,1991;Hsia et al.,1992;Fuller et al., 1993;Bowell,1994;Wilkie and Hering,1996;Raven et al.,1998;Jain et al.,1999;Meng et al.,2000;Grafe et al.,2002;Dixit and Hering,2003).The degree of pro-tonation of arsenate anion in aqueous solution is a func-tion of pH with p K a1=2.3,p K a2=6.8and p K a3=11.6 (Goldberg and Johnston,2001),resulting in arsenate spe-cies varying from H3AsO4,H2AsO4À,HAsO42À,toAsO43Àwhen pH increases from acidic region to alkalineregion(Myneni et al.,1998;Raven et al.,1998;Goldberg and Johnston,2001).On the other hand,the presence and the density of surface groups of ferrihydrite,i.e. H2O,OHÀ,are also strongly pH dependent.The point of zero charge(PZC)is approximately8.5(Jain et al., 1999;Goldberg and Johnston,2001).Hence,the modes of complexation of arsenate anions with ferrihydrite by replacing surface hydroxyl groups and/or waters are lar-gely controlled by the pH of reaction medium.Ligand exchange in the mode of bidentate binuclear in-ner-sphere complexation is the widely accepted mechanism of the adsorption of arsenate on iron oxides.It has been proposed based on infrared(Harrison and Berkheiser, 1982)and extended X-ray absorptionﬁne structure(EX-AFS)(Waychunas et al.,1993)analyses and conﬁrmed to be the dominant interaction mode of arsenate–ferrihydrite and arsenate–goethite systems(Lumsdon et al.,1984;Man-ceau,1995;Sun and Doner,1996;Waychunas et al.,1996; Fendorf et al.,1997;Foster et al.,1998;Myneni et al.,1998; O’Reilly et al.,2001;Roddick-Lanzilotta et al.,2002;Sher-man and Randall,2003;Arai et al.,2004;Cance`s et al., 2005;Waychunas et al.,2005).However,most of the stud-ies were conducted at neutral to alkaline pH.The applica-bility of the conclusions to the acidic arsenate–ferrihydrite adsorption system may be questionable.Moreover,most of the studies have dealt mainly with characterizing the bonding mechanism between arsenate anions and surface iron polyhedra without identifying arsenate species on the surface of iron oxide.Based on macroscopic measurements of the adsorption process,arsenate was adsorbedas B FeO2AsðOHÞ2;B FeO2AsðOÞðOHÞÀ;B FeO2AsðOÞ22Àon the surface of ferrihydrite at mildly acidic and alkaline pH(Jain et al.,1999).In a recent work on direct character-ization of arsenate coordination on mineral(portlandite, gibbsite,ettringite,Fe-oxyhydroxides)surfaces using FTIR, both protonated and unprotonated arsenate species(i.e.HAsO42Àand AsO43Àwere present on the goethite surfaceat alkaline pH(Myneni et al.,1998).When arsenate was ad-sorbed on schwertmannite and ferrihydrite at acidic pH(i.e. pH3),surface precipitates were proposed to form and were termed as ferric hydroxyarsenate(FeOHAs)(Carlson et al., 2002).Evidence for surface precipitation of phosphate on goethite has been observed(Ler and Stanforth,2003).Sim-ilarly,surface precipitation of ferric arsenate on ferrihydrite is likely to occur in addition to bidentate binuclear com-plexation according to the XRD and Raman spectroscopic evidence we reported recently(Jia et al.,2006).The objective of this paper was to provide further evi-dence of surface precipitation of ferric arsenate on ferrihy-drite.This was done via characterization of the interactions between arsenate and ferrihydrite in terms of bonding modes and surface arsenate species as a function of pH and coverage density by Fourier transformed infrared spec-troscopy(FTIR)and evolution of crystallinity at elevated temperature(75°C)by X-ray diﬀraction(XRD)analysis. The eﬀect of ferrihydrite synthesis media(NO3Àvs.SO42À)was also evaluated since it signiﬁcantly inﬂuenced arsenate adsorption capacity(Jia and Demopoulos,2005). Poorly crystalline ferric arsenate was used as reference material in the study.Moreover,relatively high arsenic con-centration solutions were used in this study since this is the case in important hydrometallurgical operations where arsenic removal is practiced.2.MATERIALS AND METHODS2.1.Synthesis of poorly crystalline ferric arsenatePoorly crystalline ferric arsenate was synthesized at 21°C by adjusting a0.02M As(V)/0.02M Fe(III)solution (as sodium arsenate and ferric sulfate)from initial pH1.3to 1.8with NaOH solution and maintained at that pH for1h (Jia et al.,2006).The resultant solid was separated byﬁltra-tion,washed with de-ionized water(pH2)and vacuum-dried at60°C.The chemical formula(Fe1.02AsO4Æ2.4H2O) was determined by digestion with hydrochloric acid fol-lowed by ICP-AES analysis.2.2.Synthesis of arsenate–ferrihydrite sorption samplesTwo-line ferrihydrite samples were synthesized at21°C using a slightly modiﬁed procedure from that reported in the literature(Schwertmann and Cornell,1991).Both ni-trate(Fe(NO3)3Æ9H2O)and sulfate(Fe2(SO4)3Æ5H2O)salts were used as the sources of ferric iron[Fe(III)].Brieﬂy, the Fe(III)solution was prepared by dissolving ferric ni-trate or ferric sulfate in de-ionized water.The pH of the solution was raised to$7.5in about5min using1M NaOH solution and maintained at that pH for1h with the slurry mechanically agitated vigorously.The ferrihy-drite samples synthesized from sulfate and nitrate media were termed as‘‘sulfate–ferrihydrite’’and‘‘nitrate–ferrihy-drite’’,respectively,for simplicity.The prepared ferrihydrite slurry was adjusted to diﬀer-ent pH between3and8with NaOH and HNO3and al-lowed to equilibrate for1h.Arsenate solution was introduced into the ferrihydrite slurry from a burette over a10-min period with the slurry mechanically stirred moder-ately.The pH was controlled by addition of NaOH and/or HNO3solution and allowed to equilibrate at21°C for 2weeks.The volume of adsorption slurry was500mL for all experiments and the concentration of Fe(III)in the slur-ry system was4g/L.At each pH,three initial Fe/As molar ratios(i.e.Fe/As=2,4and8)were applied.Arsenate–fer-rihydrite sorption samples were also synthesized at75°C, Fe/As molar ratio of2and4.The pH of the slurry was con-trolled constant at pH3throughout the sorption reaction. Samples were taken at1day,3day,1week,2weeks and 2months of reaction time.The synthesized arsenate–fer-rihydrite sorption products wereﬁltered,DI water-rinsed1644Y.Jia et al./Geochimica et Cosmochimica Acta71(2007)1643–1654and vacuum-dried at60°C.The equilibrium concentration of arsenic was determined by ICP-AES analysis.2.3.FTIR analysisThe infrared spectra of the samples were obtained on a Bio-Rad FTS60Fourier Transformed Infrared Spectrome-ter with a MCT liquid nitrogen cooled detector.The KBr/ sample discs were prepared by mixing0.5%ofﬁnely ground samples in KBr.The sample chamber was purged by N2gas for10min before scans were started.The measurement res-olution was set at4cmÀ1and the spectra were collected in the range of400–4000cmÀ1with200co-added scans.2.4.X-ray diﬀraction(XRD)analysisThe powder XRD patterns were collected on a Rigaku D/Max2500PC X-ray diﬀractometer with graphite mono-chromated CuK a1radiation.The powder samples were scanned from10to90°2h with increments of0.02°2h.3.RESULTS AND DISCUSSION3.1.Eﬀect of pH on the nature of adsorbed arsenateAqueous arsenate species have no direct bearing on the surface arsenate species adsorbed on mineral surfaces (Myneni et al.,1998).However,the eﬀect of complexation of arsenate ions on oxide surfaces is similar to that of pro-tonation of aqueous arsenate species.A brief discussion on infrared absorption of aqueous arsenate can assist with understanding the infrared characteristics of surface species (Myneni et al.,1998;Goldberg and Johnston,2001;Rod-dick-Lanzilotta et al.,2002).The free arsenate anion,AsO43À,is present in highly alkaline(p K a3=11.6)aqueoussolution and belongs to T d symmetry.Only m3and m4funda-mental bands are infrared active in this form.The infrared spectrum of an AsO43Àdominated solution exhibits a major band at792cmÀ1(Roddick-Lanzilotta et al.,2002).Uponprotonation or complexation with metal cations,the sym-metry decreases and splitting of the m3band occurs(Harri-son and Berkheiser,1982;Myneni et al.,1998).AqueousHAsO42Àspecies belong to the C3v symmetry.Its infrared spectrum shows two broad bands at859and689cmÀ1, the latter was assigned to stretching vibration of As–OH (Myneni et al.,1998).Curveﬁtting of the former band gave two bands at865and846cmÀ1,which were assigned to asymmetric and symmetric stretching vibration of uncom-plexed As–O,respectively(Myneni et al.,1998).Poorly crystalline ferric arsenate was used as reference material in this study to identify the possible occurrence of a surface precipitate of arsenate on ferrihydrite.This compound is not well deﬁned and often termed loosely as amorphous ferric arsenate(Krause and Ettel,1989)or amorphous scorodite(Langmuir et al.,1999),because it possesses similar bonding structures to crystalline ferric arsenate,i.e.scorodite(FeAsO4Æ2H2O).It is an unstable arsenate phase with increasing pH and tends to convert to ferrihydrite(Krause and Ettel,1989).The poorly crystalline ferric arsenate synthesized in this work was determined to have the formula Fe1.02AsO4Æ2.4H2O.Fig.1shows the eﬀect of pH on the infrared spectra of the sorption samples of arsenate on sulfate–ferrihydrite(initial molar ratio of Fe/As=2).Both detailed display of the As–O stretching vibration region(500–1000cmÀ1)and the whole range of the scanning(400–4000cmÀ1)are shown in theﬁgure.Poorly crystalline ferric arsenate shows a strong well-resolved band at838cmÀ1.Within the crystalline ferric arsenate(i.e.scorodite)structure,AsO4tetrahedra and FeO4(OH2)2octahedra connect alternately at vertices(Kita-hama et al.,1975).The arsenate is coordinated with four iron octahedra with an average As–O bond length of1.68A˚.The band at838cmÀ1was attributed to the stretching vibration of As–O coordinating to iron atom,i.e.As–O–Fe.The weakThe nature of adsorbed arsenate on ferrihydrite1645shoulder at$750cmÀ1was probably caused by the hydrogen bonding between H2O and AsO4since the H-bonding re-sulted in increased bond length and a red shift of the wave number(Myneni et al.,1998).The1625cmÀ1band was due to water O–H bending mode whereas the stretching vibration bands of O–H were located at$3194and$3373cmÀ1.The bands between950and1250cmÀ1were assigned to struc-tural SO42Àions,which were incorporated into the poorlycrystalline ferric arsenate by substitution of AsO43Àions dur-ing synthesis from sulfate medium.Sulfate ions were incorpo-rated into crystalline scorodite synthesized from sulfate solution(Singhania et al.,2005).The infrared spectra of pH3and4sorption samples also exhibited a strong,well-resolved band in the As–O stretching vibration region at similar position($833cmÀ1)to that of poorly crystalline ferric arsenate,indicating similarities of the arsenate bonding structures between sorption samples and poorly crystalline ferric arsenate.This suggested the for-mation of a ferric arsenate surface precipitate in the arsenate–ferrihydrite sorption samples synthesized in acidic media. However,the weak shoulder at$750cmÀ1on the FTIR spec-trum of the poorly crystalline ferric arsenate was missing for the sorption solids for some unknown reasons.The band at$833cmÀ1for the pH3and4sorption samples was assigned to As–O stretching vibration of the As–O–Fe coordination of ferric arsenate precipitate on fer-rihydrite.The formation of ferric arsenate phase in pH3 and4sorption samples was also supported by the O–H stretching vibration band at$3190cmÀ1.All samples showed a strong broad O–H band at$3370,but only theacidic sorption samples displayed the$3190cmÀ1O–H stretching vibration band like the case of poorly crystalline ferric arsenate.This characteristic O–H band of poorly crystalline ferric arsenate at$3190cmÀ1was fading out with increasing pH,indicating the disappearance of ferric arsenate surface precipitate at neutral and alkaline pH.As pH increased from3to8,the As–O stretching vibra-tion band shifted gradually from$833cmÀ1down to $806cmÀ1.At the same time,a new band emerged at high-er frequency(870–880cmÀ1)and its intensity was more pronounced with increasing pH.At pH8,we could clearly see the splitting of the single band into two bands.Peak deconvolution and curveﬁtting of the band produced two peaks at$806–810and$878cmÀ1(Fig.2).It is well estab-lished that at mildly alkaline pH,arsenate is adsorbed on ferrihydrite via bidentate binuclear complexation with surface iron polyhedra(Harrison and Berkheiser,1982; Waychunas et al.,1993).The band at$878cmÀ1of the pH6–8sorption samples was assigned to uncomplexed/ unprotonated As–O,whereas the$806–808cmÀ1arose from the two As–O–Fe complexed to ferrihydrite surface. Two infrared bands at824/861and817/854cmÀ1were ob-served for the arsenate adsorbed on amorphous iron oxide at pH5and9,respectively(Goldberg and Johnston,2001). The lower frequency bands at817and824cmÀ1were as-signed to the stretching vibration of As–O–Fe and the high-er frequency bands at854and861cmÀ1were attributed to ‘‘non-surface-complexed’’As–O bonds of the adsorbed arsenate species(Goldberg and Johnston,2001).It is inter-esting to note that the lower frequency band increased from 817and824cmÀ1as pH increased from5to9.This obser-vation and the band assignments are in good agreement with present work.Roddick-Lanzilotta et al.(2002)also reported that the As–O stretching vibration band of the ad-sorbed arsenate on ferrihydrite shifted from$825to $800cmÀ1as pH increased from2.6to8.In the case of bidentate binuclear complexation,two of the four As–O bonding structures are complexed to iron atoms(i.e.As–O–Fe)and the remaining two are present either both as unprotonated As–O or one as unprotonated As–O and the other one as protonated As–O–H.In com-parison,arsenate ions are coordinated to four iron atoms in ferric arsenate.According to Myneni et al.(1998),the force constant of the As–OM bond increases with coordina-tion number and decreases compared to uncomplexed As–O.Hence,for the bidentate adsorbed arsenate ion,the force constant of the two coordinated As–O–Fe is lower than that of the As–O–Fe in ferric arsenate,whereas the uncomplexed/unprotonated As–O bond has larger force constant compared to ferric arsenate.Consequently,the stretching vibration frequency of the uncomplexed/unprot-onated As–O is located at higher position while the fre-quency of the complexed As–O–Fe band is located at lower position.The increase and decrease of the As–O force constant for the bidentate binuclear complexed arsenate compared to ferric arsenate is supported by the As–O bond length(two at1.62,1.67A˚and the other two at1.71A˚, compared to1.68A˚of ferric arsenate)(Sherman and Ran-dall,2003).The shorter bond distance results in a stronger force constant and consequently higher infrared frequency.1646Y.Jia et al./Geochimica et Cosmochimica Acta71(2007)1643–1654A very weak band was observed at$700cmÀ1on the infrared spectrum of pH8arsenate–ferrihydrite sorption sample(see Fig.1).This band was reasonably assigned to protonated As–O–H bond of the adsorbed arsenate species, which was located at similar position to that of aqueous protonated arsenate species(Myneni et al.,1998).Complex-ation with metals cannot give such a low As–O stretching vibration frequency.The presence of protonated arsenate species was also proposed for the adsorption of arsenate on freshly prepared hydrous iron oxide and goethite(Myn-eni et al.,1998).It was noted that the weak As–O–H band at$700cmÀ1was absent at acidic pH indicating the ab-sence of protonated adsorbed arsenate species on ferrihy-drite at acidic pH.In a previous study using a dispersion infrared instrument(Harrison and Berkheiser,1982),three bands at875,805and700cmÀ1were observed for the ad-sorbed arsenate on hydrous ferric oxide(HFO)at pH6.5. They are very similar to the infrared bands of the pH8 sorption samples of this work(878,806and700cmÀ1).pH3and8are the extreme cases for the adsorption of arsenate on ferrihydrite in this work.At pH3,a sur-face precipitate developed and the adsorbed arsenate spe-cies were present mainly as poorly crystalline ferric arsenate.The possibility of surface precipitation of arse-nate on ferrihydrite was also suggested previously by Stanforth(1999)and Carlson et al.(2002).According to the latter research,a poorly crystalline ferric hydroxy-arsenate(FeOHAs)surface precipitate was found to form during adsorption of arsenate on schwertmannite and fer-rihydrite at pH3(Carlson et al.,2002).Similarly,surface precipitation of phosphate on goethite has been proposed in recent studies(Zhao and Stanforth,2001;Ler and Stanforth,2003).It was suggested that the adsorption reaction consisted of two phases:theﬁrst phase of rapid surface complexation followed by the second phase of slow buildup of a surface precipitate(Zhao and Stan-forth,2001).At the other extreme(i.e.pH8)arsenate was adsorbed via bidentate binuclear complexation with surface iron atoms in the form of unprotonated and probably protonated arsenate species as well(i.e.B FeO2AsðOÞ22Àand B FeO2AsðOÞðOHÞÀ, where B Fe represents the surface of ferrihydrite).This is in good agreement with the model proposed previously by Myneni et al.(1998),Jain et al.(1999)and Goldberg and Johnston(2001).Myneni et al.(1998)suggested that both protonated and unprotonated arsenate species were present as surface arsenate species adsorbed on ferrihydrite at alkaline pH.The surface arsenate specieswere proposed to be XH2AsO4,XHAsO4À,XAsO42À(X was Al or Fe)(Jain et al.,1999;Goldberg and Johnston, 2001).As pH increased from3to8,surface arsenate species shifted from poorly crystalline ferric arsenate precipitates to bidentate surface complexes.This is reasonable since poorly crystalline ferric arsenate is stable only at acidic pH and decomposes with increasing pH(Krause and Ettel, 1989).For the sorption systems whose media pHs lay be-tween3and8,both types of arsenate species were probably present on the surface of ferrihydrite,with the feature of poorly crystalline ferric arsenate being more pronounced at lower pH and the feature of bidentate complexes being more pronounced at higher pH.Similar to the Fe/As=2systems,the single As–O stretching vibration band shifted down gradually and split into two bands as pH increased from3to8for the Fe/As= 4sorption samples(Fig.3).The single band at acidic pH was assigned to As–O–Fe of ferric arsenate surface pre-cipitate.The presence of$3190cmÀ1O–H stretching vibra-tion band also indicated the development of ferric arsenate in the sorption solids at acidic pH.This band was absent onThe nature of adsorbed arsenate on ferrihydrite1647the infrared spectra of sorption samples at alkaline pH.The two bands(808and878cmÀ1)at alkaline pH are attributed to As–O–Fe bidentate–binuclear coordinating to ferrihy-drite and the uncomplexed/unprotonated As–O bond, respectively.The presence of protonated As–O–H bond could not be ruled out,since there appeared to be a weak feature at$700cmÀ1.It was interesting to note that all sorption samples from pH3to8had very similar sorption density(i.e.As/Fe$0.25,see Table1),but the surface arse-nate species varied from poorly crystalline ferric arsenate to bidentate–binuclear surface complexes as pH increased. This was indicative that the nature of surface arsenate spe-cies strongly depended on the pH of the sorption media.The spectra of pH3and4sorption samples showed strong sulfate bands between950and1250cmÀ1.Sulfate ions were apparently adsorbed onto the ferrihydrite during synthesis from sulfate medium and displaced by arsenate during adsorption as discussed elsewhere(Jia and Demopo-ulos,2005).At higher pH(i.e.5–8),the sulfate ions were substituted by arsenate ions as indicated by the absence of sulfate bands on the infrared spectra.For the adsorption systems of Fe/As=2,no sulfate band was observed for all pH samples(see Fig.1).Ferrihydrite synthesized from sulfate media was found to adsorb signiﬁcantly more arsenate than that from nitrate media(Jia and Demopoulos,2005),which could not be ex-plained by the diﬀerence in BET surface areas.Therefore,it was of interest to compare the surface arsenate species of the two types of arsenate–ferrihydrite sorption samples. Fig.4shows the eﬀect of pH on the infrared spectra of arsenate adsorbed on nitrate–ferrihydrite.Apparently the spectra are similar to those of arsenate adsorbed on sul-fate–ferrihydrite(Fig.1).Both are dominated by the strong As–O stretching vibration bands at700–950cmÀ1and strong O–H stretching vibration bands at3000–3500cmÀ1.In the As–O stretching vibration region(700–950cmÀ1),the infrared spectra of acidic sorption samples (pH3,4)exhibited a strong,well-resolved single band at similar position to that of poorly crystalline ferric arse-nate.This was indicative that a ferric arsenate surface precipitate formed when arsenate adsorbed on nitrate–fer-rihydrite at acidic pH,similar to As(V)adsorption onTable1Arsenic equilibrium concentration(mg/L)and sorption density(mol-As/mol-Fe)for the adsorption of arsenate on ferrihydrite synthesized from sulfate and nitrate medium at pH3–8and initial Fe/As molar ratio of2,4and8pH Fe/As=2Fe/As=4Fe/As=8Sulfate Nitrate Sulfate Nitrate Sulfate3198(0.49)318(0.46) 1.0(0.25) 1.9(0.25)<0.02(0.125) 4321(0.46)471(0.43) 1.9(0.25) 3.4(0.25)0.2(0.125) 5440(0.44)615(0.41) 3.6(0.25)14.8(0.25)0.2(0.125) 6498(0.43)788(0.38) 5.0(0.25)29.6(0.25)0.2(0.125) 7702(0.40)1070(0.32)15.8(0.25)100(0.24)0.6(0.125) 81178(0.30)1439(0.25)141(0.24)178(0.23) 2.1(0.125)The numbers in bracket are sorption densities,i.e.As/Fe molar ratio of the solids.1648Y.Jia et al./Geochimica et Cosmochimica Acta71(2007)1643–1654sulfate–ferrihydrite.As pH increased,the single As–O stretching vibration band split into two strong bands at $806–810cmÀ1and$878cmÀ1as observed for the sul-fate–ferrihydrite samples(see Fig.1)and were attributed to the As–O stretching vibration of bidentate binuclear complexed arsenate ions.The formation of a ferric arse-nate surface precipitate at acidic pH was also evidenced by the change of the$3190cmÀ1O–H band with pH.The spectrum of ferrihydrite displayed strong NO3Àbands at1250–1500cmÀ1.These bands disappeared after adsorption of arsenate indicating that nitrate ions previ-ously adsorbed during ferrihydrite synthesis were dis-placed by arsenate ions.3.2.Eﬀect of coverage density on the nature of adsorbed arsenateThe eﬀect of coverage density on the nature of adsorbed arsenate on ferrihydrite was evaluated.At acidic pH(i.e. pH3,4),the As–O stretching vibration region(700–950cmÀ1)was dominated by a strong well-resolved band (Fig.5).As the Fe/As molar ratio increased from2to8(i.e.the adsorption density As/Fe decreasing from0.49to0.125)for the pH3sorption systems(see Table1),the intensity of the As–O stretching vibration band decreased markedly.The As–O stretching vibration peak was located at similar position.The pH4arsenate–ferrihydrite sorptionThe nature of adsorbed arsenate on ferrihydrite1649samples with various arsenate sorption densities exhibited very similar infrared spectra to that of pH3samples.Poor-ly crystalline ferric arsenate was the major surface arsenate species.The nature of surface arsenate species was not sig-niﬁcantly inﬂuenced by arsenate sorption density and equi-librium concentration.At pH6,the major As–O stretching vibration peak was located at$822cmÀ1for all the initial Fe/As molar ratios used,i.e.2,4and8.A visible shoulder was also emerging at $878cmÀ1on the infrared spectra.A bidentate complex of arsenate was becoming detectable in addition to the major surface species of ferric arsenate.When pH increased to8,all infrared spectra exhib-ited two bands at808and878cmÀ1irrespective of the initial Fe/As molar ratios indicating that bidentate com-plexes were the dominant surface arsenate species.The equilibrium concentration of arsenic ranged from0.028 to15.8mM for the pH8sorption systems and the sorp-tion density(As/Fe)ranged from0.125to0.31.Waych-unas et al.(1993)conducted EXAFS analysis on samples with sorption density of0.001–0.1and concluded that arsenate was sorbed via bidentate surface complexation. Goldberg and Johnston(2001)used FTIR to character-ize the arsenate–iron oxide sorption solids synthesized at pH5and9with equilibrium concentration of0.1–1.0mM and suspension density of4g/L($8g/L for the present work).The obtained results showed similar As–O band to the present study,i.e.two bands located at$817–824and$854–861cmÀ1.The infrared spectrum obtained by Roddick-Lanzilotta et al.(2002)for the sorption system pH2.6and arsenic equilibrium concen-tration of$0.5mM displayed a well-resolved band $825,which is similar to the single As–O band of the acidic samples in this work.Carlson et al.(2002)pro-posed the formation of ferric hydroxyarsenate in Fe/As=2.6–3.2sorption solids prepared from solutions of pH3.Therefore,by comparing the present study with other works,it is proposed that the pH plays the most impor-tant role in controlling the nature of the surface arsenate species sorbed on ferrihydrite,whereas other conditions (e.g.coverage density,suspension density etc.)play less important factors.The fate of sulfate in synthetic ferrihydrite before and after adsorption of arsenate can also be monitored on the infrared spectra.As indicated by the bands between950 and1250cmÀ1,more sulfate was incorporated into the fer-rihydrite at acidic pH(3,4)than neutral to alkaline pH (6,8)(Fig.5).After adsorption of arsenate at pH3,4 and Fe/As=4,8,there was still measurable amount of sul-fate remaining in the ferrihydrite,while at Fe/As=2,all sulfate has been desorbed.It was proposed that sulfate ions were adsorbed on goethite as both inner-sphere and outer-sphere surface complexes at acidic pH(Peak et al.,1999). Adsorption of arsenate on the ferrihydrite involved ligand exchange with previously adsorbed sulfate ions.With increasing pH,the adsorbed sulfate ions on ferrihydrite were more easily displaceable by arsenate.At pH6,sulfate bands were visible only on the Fe/As=8infrared spectrum while no sulfate was detectable by FTIR in the pH8sam-ples.This is in good agreement with a previous study that reported the quantitative analysis of the sorption solids (Jia and Demopoulos,2005).It was suggested that surface precipitation of phosphate and arsenate on goethite occurred almost simultaneously (Zhao and Stanforth,2001;Ler and Stanforth,2003).The process of surface precipitation may involve slow dissolution of ferrihydrite,ternary complexation of Fe3+ and subsequent precipitating of arsenate(Ler and Stan-forth,2003).The process was depicted schematically in Fig.6.It was not surprising that arsenate surface precipitate formed at such a high initial Fe/As molar ratio of8where arsenic equilibrium concentration is below the detection limit(<0.02mg/L)in this work.There possibly also existed tridentate complex structure that formed initially upon contacting of the arsenate ions with ferrihy-drite given the amorphous nature of the latter.The saturation state with respect to poorly crystalline ferric arsenate in acidic media was estimated in a previous study(Jia et al.,2006).When the p K sp(ferrihydrite)=39 (Langmuir,1997)and p K sp(ferric arsenate)=22.89 (Mahoney,2002)were taken in the calculation,the log IAPð½Fe3þ ½AsO43À was obviously lower than log K sp (ferric arsenate)for the pH3,Fe/As=4and Fe/As=8 sorption systems,indicating that the systems were undersat-urated with respect to poorly crystalline ferric arsenate.For。

骨炭的应用和研究进展

骨炭的应用和研究进展黄丽娟;李海朝【摘要】骨炭作为废骨利用的一个重要途径,如果得到科学发展,对我国的畜牧业、农业的可持续发展具有实际意义.骨炭是一类结构特殊、性能优良、应用广泛的传统材料,但还没有引起足够的重视.把废弃的家禽、家畜的骨头制作成骨炭是无疑是一种很好的方法,不但可以节约资源,还有利于保护环境.【期刊名称】《应用化工》【年(卷),期】2018(047)009【总页数】5页(P2015-2019)【关键词】骨;骨炭;吸附;环境【作者】黄丽娟;李海朝【作者单位】青海民族大学化学化工学院,青海西宁 810007;青海民族大学化学化工学院,青海西宁 810007【正文语种】中文【中图分类】TQ424.2众所周知，中国有13亿多的人口，所消耗的各种肉食的量是十分庞大的，因而产生的各种牲畜，家禽的废骨的数量也是十分巨大惊人的。

我国对家畜、家禽废弃骨头的利用率不足1%[1]，因而充分利用废弃的家畜、家禽的骨头是亟待解决的问题。

一般骨头占生物体的一定比例，猪为12.9%，牛为 20.5%，羊为32.2%(羊的骨比例是取一级羊 24.3%，二级羊31.8%，瘦羊40.5%的平均值)，家禽为30%～45%[2-3]。

以2015年为例，猪、牛、羊、家禽的产肉量分别为5 486.5，700.1，440.8，1 826.3万t[4]，进而估算出当年猪骨产量为812余万t，牛骨产量为180余万t，羊骨209余万t，家禽骨782～1 264余万t。

这些骨头的处理便成了一大难题，处理的不好就会污染环境，浪费资源。

骨的组成复杂，兽骨和禽骨组成结构、化学成分也存在差异，对合理利用也产生了挑战。

目前废骨利用的方式主要有将骨炭制成骨泥(先高温软化，再高倍粉碎)、骨胶(高压蒸煮提出骨髓，再加热溶解漂白，然后在真空蒸发罐蒸浓，最后干燥后出成品)、骨油(将骨粉碎，加水煮沸至骨成粉状，取出锅面上的油浓缩成骨油)、骨粉(提完骨油剩下的骨粉加水，少量碳酸钾加热24～30 h，后晒干粉碎过筛成骨粉)、骨脂(将骨头放入高压锅内蒸煮10 h，然后取出上层液盛于水浴锅内反复加热，后提取上层淡黄色、透明、无水分杂质的骨脂)、骨灰(把畜禽骨燃烧直至成灰)、晒骨(将骨头刮去内筋浸泡在生石灰水中20～30 d，然后取出暴晒最后加以摩擦修整成晒骨)和骨炭(将骨头干馏取得油脂，再加高温使其碳化，取出粉碎得骨炭)，这些产品在食品、化工、生活用品制造、医药、农业肥料等方面具有广泛用途[5-8]。

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Journal of Hazardous Materials 144(2007)522–531Adsorption of arsenic from aqueous solution onsynthetic hydrous stannic oxideBiswaranjan Manna,Uday Chand Ghosh ∗Department of Chemistry,Presidency College,86/1College Street,Kolkata 700073,India Received 25July 2006;received in revised form 22October 2006;accepted 23October 2006Available online 27October 2006AbstractThe hydrated stannic oxide (HSO)was synthesized and arsenic adsorption behaviour is reported.HSO is found to be amorphous,and stable thermally up to 700◦C.The adsorption of As(III)is much higher than As(V)in the drinking water pH (6.5–8.5)range.The time required for reaching equilibrium is 4.0and 3.0h,respectively for As(III)and As(V).The adsorption kinetic data obtained at pH 7.0(±0.1)and temperature 27(±1)◦C follow the pseudo-second-order kinetic model best (R 2>0.98).The analyzes of isotherm adsorption data by two parameter isotherm model equations show the order to obey:Langmuir >Freundlich >Temkin for As(III),and Langmuir >Temkin >Freundlich for As(V).The monolayer adsorption capacities (mg/g)obtained for As(III)and As(V)are 15.85and 4.30,respectively.Excepting phosphate,other anions studied show no adverse effect on adsorption of As(III)onto HSO.A ﬁxed bed HSO packed column (internal diameter 0.70cm,bed height 3.7cm and particle size 0.14–0.29mm)generates 2400and 450BV of potable water (As <0.01mg/L),respectively,for As(III)and As(V)from arsenic spiked (1.0mg/L)water samples at pH 7.0(±0.1),which indicated that HSO can be used as an efﬁcient scavenger for As(III)from the contaminated water.©2006Elsevier B.V .All rights reserved.Keywords:Adsorption;Aqueous solution;Arsenic;Hydrous stannic oxide1.IntroductionOccurrence of arsenic in ground water much exceeding the tolerance limit (10␮g/L)is a global problem [1],and posed an ever increasing degree of health hazard.The Bengal delta basin [West Bengal in India and Bangladesh]has become infested with this menace,and in some pockets of this region it has assumed a life-threatening proportion causing deaths of a good number of inhabitants.The arsenic content in ground waters as reported [1]is wide in range (West Bengal in India:<10–3200␮g/L,and Bangladesh:<0.5–2500␮g/L).The cause of accumulation of arsenic in ground water in this delta region is an anoxic environ-ment around its alluvial deposits of geogenic arsenic pyrites and iron oxyhydroxide with adsorbed arsenic undergoing microbial reduction [2–5].The aquifers thus become rich in this reduced As(III)along with Fe(II).The reported ratios of As(III)/As (Total)at a depth of 30–40m in these aquifers [5]are in the range of∗Corresponding author.Tel.:+913322413893.E-mail address:ucg@ (U.C.Ghosh).0.6–0.9.This high ratio is a matter of great concern since As(III)has much greater combining afﬁnity with the thiol (–SH)group of proteins.Treatment of arsenic-contaminated ground water by sur-face adsorption method has been found to be simple and cost effective suitable for catering to the needs of this peo-ple of rural areas of third world countries like India and Bangladesh.The adsorbents studied so far are many and include amorphous iron hydroxide [6],hydrous ferric oxide [7],granular ferric hydroxide [8],ferrihydrite [9],red mud [10],activated alumina [11–13],iron oxide coated poly-meric materials [14],iron oxide coated sand [15],Fe(III)–Si binary oxide [16],iron oxide impregnated activated alumina [17],blast furnace slug [18],iron–cerium bimetal oxide [19],iron-coated sponge [20],nano-scale zerovalent iron [21–23],sulfate modiﬁed iron oxide-coated sand [24]and hydrous fer-ric oxide incorporated into naturally occurring porous diatomite [25].In our laboratory,we have used crystalline hydrous ferric oxide [26],crystalline hydrous titanium oxide [27],granular hydrous zirconium oxide [28]and iron(III)–tin(IV)binary mixed oxide [29]for arsenic removal.A pilot-scale run0304-3894/$–see front matter ©2006Elsevier B.V .All rights reserved.doi:10.1016/j.jhazmat.2006.10.066B.Manna,U.C.Ghosh/Journal of Hazardous Materials144(2007)522–531523Nomenclaturea1,a2,b1,b2constants of power function and simple Elovich equationA,B Temkin constantsb Langmuir constantBV bed volumeC0initial adsorbate concentration(mg/L)C e equilibrium adsorbate concentration(mg/L)C t the concentration(mg/L)of adsorbate at any time,tK0constantK2pseudo-second-order rate constantK ad pseudo-ﬁrst-order rate constantK f Freundlich constantM mass of HSO added(g/L)n Freundlich constantpH ZPC zero point surface charge pHp K a negative logarithm of dissociation constantq e uptake capacity at equilibrium(mg/g HSO)q t uptake capacity at time t(mg/g HSO)R L separation factorR2correlation coefﬁcientV volume of solution(mL)Greek symbolsαconstantθ0Langmuir constant(mg/g)for removal of excess iron and arsenic from contaminated water using manganese dioxide and crystalline hydrous ferric oxide packed columns has also been reported[30]from our laboratory.Literature survey revealed that there are some adsorbents which have greater afﬁnity for As(V)than for As(III)both in the acid and drinking water pH range;but hydrous oxides of tetrava-lent metals clearly showed the opposite property as reported from our laboratory[28].Since the underground aquifers at the depth of30–40m in the Bengal delta basin region are reported[5]to contain arsenic mainly as As(III),the objective of the present work is to develop adsorbents for preferen-tial removal of As(III)from water,and offers still a frontier area of research work.As(III)exists as As(OH)3[p K a1=9.2] in the pH range of potable water and is a softer Lewis base than the anionic As(V).Sn(IV)[4s24p64d10]is a softer Lewis acid than either Fe(III)or Ti(IV).So hydrated stannic oxide (HSO)is expected to have greater afﬁnity for As(III)than for As(V).With this view,the present paper reports the synthesis, characterization and systematic arsenic adsorption behaviour of hydrous stannic oxide(HSO).Finally,aﬁxed bed column adsorption test on arsenic removal from water on a lab-bench scale has also been reported.2.Materials and methods2.1.Preparation of adsorbentHydrous stannic oxide(HSO)was prepared by neutraliz-ing0.1M sodium stannate with drop-wise addition of0.1M ammonium chloride solution.The white gel-like precipitate of stannic hydroxide formed was aged for ten days with mother liquor.Decanting the mother liquor,thereafter,the precipitate was washed with de-ionized water till alkali free.Theﬁltered white mass was dried in an air oven at50–60◦C.The dried mass when treated with cold water was broken intoﬁne particles,and sieved to collect the particles ranged between0.14and0.29mm.2.2.Arsenic solutionsA standard stock As(III)solution(1000mg/L)was prepared by dissolving0.1320g of As2O3(99.9%Aldrich,USA)in l0mL4%(w/v)NaOH,acidiﬁed with2.0mL concentrated HCl and then diluted to100mL with arsenic free de-ionized water.The working solutions of As(III)of desired concentra-tions were made by diluting the stock with0.2%(v/v)HCl.The stock solution was prepared freshly after every three days and frozen to prevent oxidation.The standard stock As(V)solution (1000mg/L)was also prepared by dissolving4.1601g of sodium arsenate(Na2HAsO4,7H2O,E.Mark,Germany)with arsenic free de-ionized water in to a1000mL volumetricﬂask.The working solutions of As(V)of a required concentration were made by diluting the stock with de-ionized arsenic free water and frozen.2.3.Analytical methodsTotal arsenic as well as As(III)species was analyzed by hydride generation atomic absorption spectrophotome-ter(Perkin-Elmer-3100)and UV–vis spectrophotometer (HITACHI model-3210).Total dissolved inorganic arsenic in water sample was determined by adding HCl(32%,v/v)and KI (10%,w/v)where As(V)was reduced to trivalent state,which was converted to arsine with3%(w/v)NaBH4.The arsine gas was carried over to theﬂame of atomic absorption spectropho-tometer,and absorbance value was noted at193.7nm against blank,and compared with standard for arsenic concentration. As(III)in water sample was determined by selective reduction with3%NaBH4to volatile arsine at pH4.8in acetate buffer.This method allows rapid determination of arsenic species down to 1␮g/L with±4%accuracy[31].In spectrophotometric method, the generated arsine gas was adsorbed in silver dietyl dithiocar-bamate(SDDC)(G.R.,E.Marck)solution in chloroform solvent, and absorbance was measured at535nm against blank and com-pared with the standard curve[31].Some common ions,viz., phosphate,sulfate,nitrate,chloride,ﬂuoride and bicarbonate in water sample were estimated by ion-chromatography(IC)using standard method[31].Stannic ion in solution was determined colorimetrically.Here,a deﬁnite volume(40.0mL)sample solu-tion was concentrated to10.0mL,and used for developing color in dithiol method[32].524 B.Manna,U.C.Ghosh/Journal of Hazardous Materials144(2007)522–531X-ray diffraction(XRD)analysis of the synthetic oxide was conducted by the powder method with Philips PW1830gen-erator.Thermo gravimetric(TG)and differential thermal(DT) analyzes of the adsorbent were recorded using a seterm ana-lyzer in argon atmosphere at a heating rate of20◦C/min over a temperature range of30–1000◦C.Point of zero charge(pH zpc) was determined from electrophoretic mobility(EM)of the oxide using a Zeta-Meter.The electrophoretic mobility of solid sor-bent suspensions containing0.02%solid in0.01(M)NaCl was determined at various pH.Surface area of the sorbent particles was measured by Brunauer–Emmett–Teller(BET)method by adsorption of N2gas at liquid N2temperatures using Micromer-ities high-speed surface analyzer.2.4.Batch experiments2.4.1.pH effectThe effect of pH in a range from2to10on arsenic adsorp-tion was conducted by mechanical agitation(250–260rpm)at room temperature,27(±1)◦C.Here,0.1g HSO(particle size: 0.14–0.29mm)was added with50.0mL aliquots of As(III)and As(V)solutions(5.0and10.0mg As/L)taking separately in well capped100mL polythene bottles.The required initial pH(pH i) of test solutions were adjusted using0.1(M)HCl and/or0.1(M) NaOH.Agitating the reaction mixtures for5.0h for both As(III) and As(V),the solutions wereﬁltered using0.45␮m membrane ﬁlter;and analyzed for residual arsenic and equilibrium solution pH(pH f).2.4.2.Adsorption kinetic studiesAdsorption kinetic experiments were performed by batch method at pH i7.0(±0.1)and room temperature,27(±1)◦C. Here,50.0mL solution of arsenic[concentrations(mg/L):5.0 and10.0]was taken in to100mL polythene bottles with0.1g HSO,and agitated(250–260rpm)using a shaker.A pH elec-trode was inserted into the solutions10min before withdrawing of samples for pH measurement,and0.1(M)NaOH or0.1(M) HCl was added,if required,for pH adjustment(∼7.0)during experimental run.The reaction mixtures including bottle were withdrawn at a time interval of15min from starting till the equi-librium reached.Theﬁltered sample solutions were analyzed for arsenic.2.4.3.Adsorption isothermAdsorption isotherm experiments were conducted at room temperature,27(±1)◦C and at pH7.0(±0.1)by batch adsorption procedure as described above.Here,six different concentrations ranging from1.0to10.0mg/L were used sep-arately each for As(III)and As(V).The agitation time and speed used were5.0h and250–260rpm,respectively,for equilibrium. The solution pH was adjusted twice in the middle(after1.0and 4.0h of agitation)using0.1M HCl and/or0.1M NaOH solution. Similar intermediate pH adjustment was made by Jang et al.[25]for arsenic adsorption isotherm experiment using250rpm shaking speed.Residual arsenic was determined in theﬁltered sample.2.4.4.Column testThe glass columns(internal diameter:0.70cm)were packed with HSO(particle size:0.14–0.29mm,pH∼7.0)up to a height of3.7cm.The laboratory tap water of Presidency College(PC) was spiked with As(III)and As(V)separately up to a level of 1.0mg/L.The tap water used for sample preparation was the ground water of PC.The tap water composition[27]was found to be almost the representative of theﬁeld samples of arsenic-contaminated belt of West Bengal in India.The arsenic spiked PC tap water samples were passed with a downﬂow(ﬂow rate:∼100mL/h;contact time:∼35s)direction to evaluate column efﬁciency for As(III)and As(V)removal.3.Results and discussions3.1.Physico-chemical characteristics of HSOThe X-ray diffraction analysis showed that the synthesized hydrous stannic oxide(HSO)was of amorphous variety.The thermal stability of the hydrous oxide is shown in Fig.1.Thermo gravimetric(TG)analysis recorded8.26%and6.81%weight loss in drying temperature range below190and190–700◦C, respectively.The weight loss below and above190◦C corre-sponds to dehydration and phase transition of the oxide.The calculated weight loss of4.22%on drying up to110◦C was attributed to the irreversible loss of physically adsorbed water molecules,and that would be the moisture content of the hydrous oxide.This is conﬁrmed from a broad endoergic DTA peak (Fig.1)at110◦C.The differential thermal analysis(DTA) showed no exoergic peak(Fig.1),so the chemical reactions,viz., polymerizations or crystallizations did not take place on drying up to700◦C.Electrophoretic mobility studies of the oxide sus-pensions at various pH values showed that pH ZPC of the synthetic oxide was4.6to6.4.The wide range of pH zpc value of HSO is presumably owing to the lower basicity of tin(IV).Thus,HSO could be a surface-active oxide.The anion and cation exchange properties of HSO were also reported[33]at pH below4.6and above6.4,respectively.Fig.1.Thermo gravimetric and differential thermal analysis of HSO.B.Manna,U.C.Ghosh/Journal of Hazardous Materials144(2007)522–531525Fig.2.(a)Effect of initial pH on arsenic adsorption by HSO at27(±1)◦C and (b)the plot of equilibrium pH(pH f)vs.initial pH(pH i).Some determined physico-chemical characteristics of HSO are:nature:amorphous;metal content:61.60%;moisture con-tent:4.22%;bulk density:1.38g/cm3;particle size:0.14to 0.29mm;surface area:140.8m2/g;pH zpc:4.6–6.4;solubility (pH6.5–8.0)=0.02mg/L.3.2.Effect of pHThe effect of initial pH(pH i)on adsorption of arsenic was determined by loading5.0mg As/g of adsorbent in the pH i range of2.0–10.0and at temperature27(±1)◦C.Fig.2a demonstrates the efﬁciency of arsenic removal,and shows the increase in adsorption capacity(q e,mg/g)for As(III)with increasing pH i from2.0to6.0,that remained almost same in the pH i range6.0to8.0while that for As(V)decreased with increasing pH i from2.0to10.0.In general,the adsorp-tion of As(III)is less pH dependent,while that of As(V) is highly pH dependent.The determined q e-value for As(V) was signiﬁcantly greater than As(III)at pH i2.0,while that for As(III)was clearly greater than As(V)at pH i>3.0.Results (Fig.2a)show that the percentage of As(V)adsorption was found to be about12.0to13.0%of As(III)at pH i10.0. Additionally,the equilibrium solution pH(pH f)was found to increase(Fig.2b)somewhat more for As(V)adsorption onto the HSO than the increase found for As(III)in the pH i-range of 2.0–8.0.The lower adsorption capacity for As(III)than As(V)of HSOat pH i2.0is due to the electrostatic hindrance between posi-tive surface sites of the adsorbent[pH zpc=4.6to6.4]and H␦+center of OH groups in As(OH)3[p K a1=9.2].For As(V),thepositive surface sites of the adsorbent attracts the anionic adsor-bate species[p K a1=2.6and p K a2=6.4of H3AsO4],and theadsorption capacity is high.With increasing pH i,the positivecharge density of adsorbent surface decreases and,therefore,the adsorption of non-ionic As(III)increases due to decreasein electrostatic hindrance while that of As(V)decreases due todecreasing electrostatic attraction.At pH i∼6.0,HSO surfaces are to be negative and the As(V)anion being repelled due tolike nature of charge,and the adsorption capacity for As(V)is about50%less than that obtained at pH i5.0.In contrast,that decrease for As(III)adsorption was not found,becausethe earlier noted adsorptive force persists up to pH i8.0.AtpH i>8.0,the less adsorption capacity of adsorbate ions are obvi-ously due to(i)increase in electrostatic hindrance between likecharges on adsorbent and adsorbate,and(ii)increase in compet-ing OH−concentration for sorption sites.Thus,the adsorptionof As(III)and As(V)onto HSO takes place with differentmechanisms.The high value of As(III)adsorption data in the pH i rangeof4.0–8.0resembled closely to our earlier studies[26,28].The present observation on As(III)adsorption is similar tothe literature available reports made by other workers usingdifferent adsorbents[6,14,16].However,the pH i value for max-imum As(V)adsorption differ from the reports made by others[11,16,26].The lower pH i value for high arsenic adsorptioncompared to other literature available data may be due to thepresence of4d10conﬁguration in valence shell of the metal ionin HSO.The mechanisms for As(III)and As(V)adsorption aresimilar to that reported by others.In addition,As(V)adsorptionat pH i≤6.0can be attributed by ligand(or anion)exchange[33] phenomenon.This supports the increase in pH of equilibrium solution.Sn–OH(s)+−O–AsO(OH)2(aq)+H+Sn+–−OAsO(OH)2(s)+H2O(1)The adsorption capacity for As(V)was very high at pH i 2.0despite some solubilization loss of the HSO(0.03mg/L) at that pH i,because the anion-exchange property increases with decreasing pH.The high adsorption capacity of HSO for As(III) may be due to soft Lewis acid–base type interaction.,where SnO2·x H2O(Sn4+,a4d10species)functions as soft Lewis acid and As(OH)3a soft Lewis base,giving a surface complex. Other workers,however,have earlier reported the probability of surface complex formation without specifying the stabilizing interaction.SnOH(s)+As(OH)3 HO–Sn–As(OH)3(s)(2)where SnOH(s)stands for solid HSO in noted two adsorption reactions(Eqs.(1)and(2))for mechanisms.526 B.Manna,U.C.Ghosh/Journal of Hazardous Materials144(2007)522–531Fig.3.Effect of drying temperature of adsorbent on As(III)adsorption at pH 7.0(±0.1).3.3.Effect of drying temperatureThe synthetic HSO was dried for an hour at the temper-atures25,50,100,150and200◦C separately,and used for As(III)adsorption.Results(Fig.3)show a small increase in adsorption capacity of As(III)when the drying temperature was increased from25to50◦C,and that,thereafter,decreased slowly with increasing drying temperature of adsorbent.This indicates the increase of surface active sites for As(III)adsorption with increasing drying temperature from25to50◦C which is due to the loss of physically attached(H-bonded)water molecules from the surface and,increases availability of surface active OH groups for arsenic adsorption.The decrease in adsorption capac-ity with increasing temperature for drying above50◦C is due to the loss of surface active OH sites as water which agrees well with the loss of weight in thermal analysis(Fig.1).Simi-lar results were reported for As(III)adsorption using crystalline hydrous titanium oxide[27].In contrast,crystalline hydrous fer-ric oxide showed an increase in arsenic adsorption capacity with increasing temperature for drying up to300◦C[26].Thus,it can be suggested that HSO,dried at50◦C,could be used for As(III) removal.3.4.Kinetic modelingFig.4a and b show the time dependent arsenic adsorption data at pH7.0(±0.1)and temperature27(±1)◦C.For the ini-tial arsenic loads2.5and5.0mg per g adsorbent,the adsorption capacities(mg/g)nearly obtained were1.80and3.33,2.30and 3.96,and2.49and4.26for As(III);and those were0.47and 0.74,0.56and1.00,and0.68and1.09for As(V),respectively, at1,2and4h of contact.The results show that the initial rate of adsorption was fairly rapid for both As(III)and As(V).∼70% of the adsorbed amount took place in an hour of agitation and, the remaining∼30%adsorption took place taking additional3.0 and2.0h.That is,the time required for reaching adsorption equi-librium was4h for As(III),while that was3h for As(V).Results show(Fig.4a and b)that the adsorption of As(V)onto HSO is faster than As(III)which is similar to the observation reported by Raven et al.[9].However,the equilibrium time obtainedin Fig.4.Adsorption kinetic data for:(a)As(III)at pH7.0(±0.1)and at27(±1)◦C and(b)As(V)at pH7.0(±0.1)and at27(±1)◦C.the present case is found to be less compared to the studies made using some other oxides[26,27].The data shown in Fig.4a and b are used for analysis taking the linear form of the following kinetic model equations: The pseudo-ﬁrst-order equation[34]:log(q e−q t)=log q e−K ad2.303t(3)where q e and q t are the adsorption capacity(mg/g)at equilibrium and at time,t,respectively;and K ad the pseudo-ﬁrst-order rate constant(time−1).The pseudo-second-order equation[35]:tt=12e+1et(4)where both q e and q t have the same meaning as above.K2is the pseudo-second-order rate constant(g mg−1time−1).The Bhattacharya–Venkobachar equation[36]:ln(1−U(t))=K B t(5) where U(t)=C0−C t/C0−C e,where C0,C t and C e are the con-centrations(mg/L)at t=0,t and equilibrium,respectively.The power function equation[37]:log q t=log a1+b1log t(6) where a1and b1are some constants,q t the adsorption capacity (mg/g)at time,t.The Elovich equation[38]:q t=βln(αβ)+βln t=a2+b2ln t(7)B.Manna,U.C.Ghosh/Journal of Hazardous Materials144(2007)522–531527Fig.5.Pseudo-second-order plot of:(a)As(III)adsorption at pH7.0(±0.1)and at27(±1)◦C and(b)As(V)adsorption on HSO at pH7.0(±0.1)and at27 (±1)◦C.whereαandβhave usual signiﬁcation of the model equation, and a2=βln(αβ)and b2=β.The estimated kinetic models and related parameters with linear regression coefﬁcient(R2)are shown in Table1.The present kinetic adsorption data for both As(III)and As(V)on HSO describe the pseudo-second-order model(Fig.5a and b)best(0.98<R2<1.0)while,those also describe the pseudo-ﬁrst-order model well(0.97<R2<0.98).Other models tested describe the present kinetic data,in general,poorly.The sim-ple Elovich equation explained the kinetic data fairly well. The regression coefﬁcient(R2)values obtained for As(III) are0.9582and0.9919,and those for As(V)are0.9932and 0.9210,respectively for the studied initial arsenic loads2.5and 5.0mg/g.The Bhattacharya–Venkobachar[36]kinetic model equation describes the kinetic data well(R2=∼0.98)obtained for2.5mg/g of As(III)while the power function kinetic model equation describes the kinetic data well(R2=∼0.98)obtained 5.0mg/g of As(III).However,no systematic order is found for ﬁtting the kinetic data with analyzed model equations.As the present kinetic data describe the pseudo-second-order model best,thus the arsenic adsorption onto HSO obeyed the pseudo-second kinetics.3.5.Intra-particle(pore)diffusionThe adsorption process onto an adsorbent follows three steps, viz.,ﬁlm diffusion,macro and micro-pore diffusions phenom-ena.The slowest of the three-step controls the overall rate pro-cess.Generally,intra-particle(pore)diffusion is the rate-limiting step in batch process,while for a continuousﬂow systemﬁlm diffusion is more likely to be the rate-limiting step.The intra-particle(pore)diffusion rate constant(k i)can be deﬁned as the gradient of the plot of sorption density versus square root of time [39,40].The plot of q t versus t1/2was found to be linear for a wide range of contact time for arsenic sorption indicating the presence of external mass transfer(Fig.6).It shows three separate regions: initial part of the curve is attributed to mass transfer[41](slope k1)taking place with boundary layer(ﬁlm)diffusion,while theﬁnal linear portion(s)of the curve indicate intra-particle diffusion(slope k2and k3).The values of k2and k3indicate that the pores are micropore and intra-particle diffusional resis-tance is due the micro pores only which is more pronounced for large As(V)compared to small As(III)species.This is evident from smaller values of k2and k3for As(V).Moreover,pore sorption of As(III)and As(V)are concentration dependent andTable1Adsorption kinetic model equations with different parameters evaluated on adsorption of arsenic by HSO at pH7.0(±0.1)and at27(±1)◦CAdsorption kinetic model equations Parameters Initial As(III)loads(mg/g of adsorbent)Initial As(V)loads (mg/g of adsorbent)2.5 5.0 2.55.0Pseudo-ﬁrst order:log(q e−q t)=log q e−K ad/2.303t K ad×102(min−1) 1.36 1.49 1.75 3.09 R20.97820.98220.97190.9723Pseudo-second-order:t/q t=1/K2q2e+t/q e K2×102(g/(mg min)−1) 1.35 1.16 2.63 1.20 R20.99900.99910.99600.9804Bhattacharya–Venkobachar:ln(1−U(t))=K B t;U(t)=C0−C t/C0−C e K B(min−1)×102 1.50 1.360.20 1.88 R20.98110.87980.85220.9559Power function:log q t=log a1+b1log t a10.4533 1.34780.05990.0432 b10.32840.21160.47760.6286 R20.90650.98110.91570.9098Simple Elovich: q t=a2+b2ln t a2−0.41810.4402−0.2635−0.7072 b20.54670.70670.17800.3383 R20.95820.99190.99320.9210528 B.Manna,U.C.Ghosh /Journal of Hazardous Materials 144(2007)522–531Fig.6.Intra-particle diffusion rate constants (k i ,mg g −1min −1/2)of arsenic adsorption at pH 7.0(±0.1)and at 27(±1)◦C.the rate of pore diffusion increases with increasing concentra-tion.The Bangham’s equation,furthermore,as suggested by Aha-roni and Ungarish [42](Eq.(8))has been applied for verifying the pore-controlled adsorption phenomena,log log C 0C 0−q t m =log K 0m 2.303V+αlog t (8)where C 0is the initial As(III)concentration (mg/L),V the volume of solution (mL),m the mass of adsorbent (g/L),q t the amount of adsorbate retained in solid at time t ,αand K 0are con-stant.The plots of log(log C 0)/(C 0−q t m )versus log t (Fig.7a and b)obtained for As(III)and As(V)adsorption data show better correlation coefﬁcient values for As(III)(R 2=0.97–0.99)than for As(V)(R 2=0.73–0.75).This supports the previous conclusion that the adsorption kinetics for the As(III)is more pore-diffusion controlled than As(V).This is presumably due to larger size of As(V)than As(III)in aqueous solution that hindered to enter into the pore.3.6.Isotherm envelopeFig.8a and b show the equilibrium isotherm data for arsenic at pH 7.0(±0.1)and temperature 27(±1)◦C.The following linear isotherm model equations are used for analyzing the equilibrium adsorption data of As(III)onto HSO:Freundlich equation [43]:q e =k f C 1/ne .The linear from of the equation is:log q e =log K f +1nlog C e (9)Langmuir equation [44]:q e =(θ0bC e )/(1+bC e ).The linear from of the equation applied for isotherm data analysis is:1q e =1θ0bC e +1θ0(10)Fig.7.Pore diffusion controlled adsorption kinetics of:(a)As(III)at pH 7.0(±0.1)and at 27(±1)◦C and (b)As(V)at pH 7.0(±0.1)and at 27(±1)◦C.The Temkin isotherm equation applied for isotherm analysis in the following form [45]:q e =A +B log C e(11)Fig.8.Equilibrium isotherm data for:(a)As(III)adsorption at pH 7.0(±0.1)and 27(±1)◦C and (b)As(V)adsorption at pH 7.0(±0.1)and 27(±1)◦C.B.Manna,U.C.Ghosh/Journal of Hazardous Materials144(2007)522–531529Table2Estimated isotherm parameters for As(III)-sorption onto HSO at pH7.0(±0.1) and at27(±1)◦CEstimated isotherm parametersAs(III)As(V)LangmuirR20.98500.9707θ0(mg/g)15.850 4.303b(L/mg)0.27270.0488FreundlichR20.97200.8935K f 3.5480.1521n 1.569 1.0739TemkinR20.80690.9553A 4.2472−0.928B 4.9862 2.4739where q e is the amount adsorbed at equilibrium(mg/g)and C e the equilibrium concentration of adsorbate in solution(mg/L), K f and n are the Freundlich constants corresponding to adsorp-tion capacity and adsorption intensity,respectively;θ0and b are Langmuir constants related to maximum monolayer capac-ity and energy of adsorption,respectively;A and B are Temkin constants.The isotherm parameters and linear regression coefﬁcient (R2)values obtained from the linear plots are shown in Table2. The best-ﬁt plots are shown in Fig.8a and b.The results (Table2)obtained from the isotherm analysis indicated that the Langmuir model describes the As(III)and As(V)adsorp-tion isotherm data onto HSO best.The R2values were of0.9850 and0.9707for As(III)and As(V),respectively.The order ofﬁt-ting of the present data with the two parameter isotherm models is:Langmuir>Freundlich>Temkin for As(III),and Lang-muir>Temkin>Freundlich for As(V).The Langmuir isotherm model describes the adsorption data for As(III)better than for As(V).Similar order on theﬁtting of the As(III)adsorption data was reported by Zeng[16]and Raven et al.[9]at pH6.5and 9.2,respectively.The betterﬁt of As(III)adsorption data with the Langmuir isotherm was reported earlier[27,28].The different trend found in describing the adsorption data with two-parameter isotherm models suggests different adsorption mechanisms for As(III)and As(V)adsorption onto HSO(Fig.9).The evaluated values for mono-layer adsorption capacity(θ0) and adsorption energy(b)from the Langmuir isotherm plots show(Table2)that the values for As(III)are much higher than for As(V).The higher values of the Freundlich and Temkin con-stants for As(III)compared to As(V)also support the fact that HSO has much greater afﬁnity for the lower valent species at pH 7.0(±0.1).This agreed well to the results(Fig.2a)obtained in the studies of pH effect.Despite well-ﬁt of adsorption data with used three different isotherm models,yet it can not high light the mechanis-tic implications.However,the dimensionless parameter,R L [=1/(1+C0b)][46],where b is the Langmuir constant and C0 is the initial adsorbate concentration,can predict whetherthe ngmuir isotherm plot of:(a)As(III)at pH7.0(±0.1)and at27(±1)◦C and(b)As(V)at pH7.0(±0.1)and at27(±1)◦C.adsorption is favorable or unfavorable.The calculated R L-values for a range of initial concentrations(1.0–20.0mg/L)of adsor-bates are found to lie between0and1.0,and indicate the favorable adsorption of both arsenic species onto HSO at pH 7.0(±0.1).The R L-values for As(III)are found always lower than those for As(V)suggested the more adsorption favorability of the reduced arsenic compared to that of the oxidized species.3.7.Adverse effect of some anionsThe adverse effect of some anions(commonly found in groundwater)on As(III)adsorption onto the adsorbent were studied in the pH range5.0–8.0.The studied sulfate,nitrate, chloride andﬂuoride ions show no signiﬁcant interference while phosphate ion shows(Fig.10)notable effect on As(III)-adsorption,particularly at higher phosphate concentrations.The arsenic(III)removal was found to decrease from87.86to79.30%Fig.10.Effect of phosphate on As(III)adsorption by HSO.。