DK1300_ch01

7/25/2019 DK1300_ch01

1/84

1

Adsorption and Ion-ExchangeProperties of Engineered

Activated Carbons andCarbonaceous Materials

Michael Streat, Danish J. Malik, and Basudeb Saha

Loughborough University, Loughborough, Leicestershire, United Kingdom

I. INTRODUCTION

The importance of environmental pollution control has increased significantly inrecent years. Environmentalists are primarily concerned with the presence of heavymetals, pesticides, herbicides, chlorinated hydrocarbons, and radionuclides ingroundwater, surface water, drinking water, and aqueous effluents due to their hightoxicity and impact on human and aquatic life.

Several techniques have been developed and used to remove and=or recover awide range of micropollutants from water and a variety of industrial effluents.

Adsorption using activated carbon is well established for the removal of organic

molecules from aqueous solution but to a much lesser extent for the removal oftoxic heavy metals. Of course, polymeric ion-exchange resins are applied for watertreatment and for trace metal removal under extreme conditions, ranging fromhighly acidic to highly alkaline solutions. The high sorption capacity of thesematerials, usually greater than that of carbonaceous adsorbents, and good selectivitytoward metal ions render them attractive candidates for a wide range of appli-cations. However, the use of ion-exchange resins involves significant capital andoperating costs. Activated carbons are generally a cheaper alternative because theyare derived from renewable natural materials although the operating costs remain

a significant factor.

Copyright 2004 by Taylor & Francis Group, LLC

7/25/2019 DK1300_ch01

2/84

The removal of toxic metal ions from dilute or concentrated solutions hasreceived considerable attention in the last few decades. In recent years, stringentstatutory regulations were introduced to reduce the discharge of toxic metals tolow levels at the source, particularly from plating shops and other metal processingindustries. The Environmental Protection Agency (EPA) and the EuropeanCommunity (Directive 98=78=EC [1]) have highlighted the most common heavymetals arising in residual water, and the maximum admissible concentrations aregiven in Table 1. Cadmium and mercury are two of the most toxic metals presentin the aqueous environment; hence, their maximum permissible concentrations indrinking water have been set at 5 and 1 mg=L (ppb), respectively, and this presents a

particularly challenging problem. The threshold limits given in Table 1 are achievedpartially by minimization and recycling of existing resources. With the increasingdemand for cleaner water, attention has been focused on improvements to existingtreatments and the development of new techniques and materials.

The removal of metal contaminants from effluent streams has the advantageof reducing the cost of waste disposal. In most cases, the treatment of wastewatergives rise to secondary effluents. Efficiency of such processes can be improved byrecycling treated water and=or by metal recovery. General methods applied tothe removal of metals include ion exchange, precipitation, coagulation, flocculation,

evaporation, and membrane processes. By using ion exchange or adsorption, most

Table 1 Maximum Admissible Concentrations of Undesirableand Toxic Metals in Water

Substance

Maximum admissible

concentration (mg=L) Comments

Arsenic 10Antimony 10Cadmium 5Chromium 50Copper (3,000) Guide level after standing

12 h at point of consumptionCyanides 50Iron 50Lead 50 In running water

Mercury 1Nickel 50Selenium 10Zinc (5,000) Guide level after standing

12 h at point of consumption


7/25/2019 DK1300_ch01

3/84

of the water can be recycled without the need for further treatment. In some cases,the metal can also be recovered in a useful form.

A variety of materials have been investigated for the removal of metals frommetallurgical effluents. Conventional activated carbons are used extensively in watertreatment for removal of color, odor, and organic contaminants [2,3]. Thesecarbonaceous materials possess the potential for removal of inorganic species fromeffluent streams. Activated carbons have high porosities and high surface areas andare prepared from readily available carbonaceous precursors such as coal, wood,coconut shells, and agricultural wastes. These precursors are normally exposed toa number of different activation methods in an effort to achieve an activated carbonwith the most favorable properties for a particular application. The texture ofactivated carbons can be adapted to suit the situation by adequate choice of theactivation procedure.

Removal of metals by conventional activated carbons has been studied by anumber of authors (46). In general, ordinary activated carbons possess a largesurface area but have a relatively low capacity for metal ions. Modified activatedcarbons have been examined as alternatives to conventional polymeric ion-exchangeresins. By far the most widely developed large-scale application of activated carbonin hydrometallurgy is the recovery of gold from dilute cyanide leach solutions. Anextensive review of this process is given by Bailey [7]. Tai and Streat [8] reported theion-exchange reactivity of oxidized carbon for the removal of copper, zinc, andnickel from solution.

In this chapter, we discuss the preparation, properties, and metal sorptionperformance of a range of as-received and oxidized samples of granular and fibrousactivated carbon that were either prepared or modified in our laboratory. Sampleswere evaluated for the removal of trace toxic metal ions from aqueous solutions.Batch and column experiments were performed to elucidate the relationshipbetween sorptive performance and the physical and chemical structure of thesematerials.

The EU and UK national standard for any individual pesticide in drinkingwater at the point of supply is 0.1 mg=L (0.1 ppb), with a maximum of 0.5mg=Lfor all detected compounds. In 1990, the most frequently detected pesticides in

UK drinking water supplies were atrazine, simazine, isoproturon, diuron, chloroto-luron, and mecoprop. Atrazine is one of the most difficult herbicides to removefrom potable water supplies and as a consequence was prohibited from nonagricul-tural uses in England and Wales in 1993. Since that time there has been an increasein the use of alternative herbicides for both agricultural and nonagriculturalpurposes. Imazapyr and triclopyr are two of a group of four herbicides (benazolin,bentazone, imazapyr, and triclopyr) that have been identified as alternatives. Theseherbicides are more soluble in water than atrazine and therefore also constitute apotential pollution hazard. The average concentration of the pesticides in source

waters is generally below the legal limit of 0.1 mg=L. However, seasonal variations


7/25/2019 DK1300_ch01

4/84

in the use of pesticides results in concentrations that significantly exceed the limit,so that the source cannot be used for drinking water production. Data provided bythe UK Environment Agency show the maximum concentration of each herbicideunder investigation during sampling in 1995, 1996, and 1997 (see Table 2).

The adsorption of herbicides onto activated carbon is also discussed in detailin this chapter, and the findings are compared with data for a set of novel hyper-cross-linked polymer phases that offer an alternative approach for the treatmentof potable waters.

II. PREPARATION AND PROPERTIES OF

ACTIVATED CARBON

The extraordinary ability of carbon to combine with itself and other chemicalelements in different ways is the basis of organic chemistry and life itself [10]. Asa consequence, there is a rich diversity of structural forms of solid carbon becauseit can exist as any of several allotropes. It is found abundantly in nature as coal oras natural graphite and also in much less abundant form as diamond. Engineered

carbons can take many forms, e.g., coke, graphite, carbon and graphite fiber, carbonfibrecarbon composite, carbon monoliths, glassy carbon, carbon black, carbonfilm, and diamond-like film.

The principal reasons that engineered carbons find extensive use as adsorbentsare their porous and highly developed internal surface area and the complex natureof their surface chemical structure. In this chapter, we review the synthesis, struc-ture, and adsorption properties of engineered carbons with reference to the bodyof work carried out in our research laboratory. We also present techniques thatwe have employed to modify the surface properties of engineered carbons by

introducing heteroatoms such as oxygen, sulfur, nitrogen, and phosphorus. Active

Table 2 Maximum Concentration of Herbicides in UK Water during 1995, 1996,and 1997

1995 1996 1997

Maximumconc. (mg=L)

Maximumconc. (mg=L)

Maximumconc. (mg=L)

Atrazine 1.37 1.8 5.42Benazolin 0.093 0.525 0.331Bentazone 1.12 0.423 1.84Imazapyr 0.058 0.074 < 0.04Triclopyr 0.122 0.988 0.16

Source: Ref. 9.


7/25/2019 DK1300_ch01

5/84

carbon fibers (unfunctionalized and functionalized) as adsorbent materials are alsoof considerable interest because they offer a kinetic advantage over granular acti-vated carbons (typical dimension 0.45 mm) owing to their smaller diameters(620 mm).

We do not apologize for omitting other exciting areas such as the advancesin the synthesis of carbon nanotubes, nanocones, and multiwalled carbon spheresand in the production of carbon films (employing physical vapor deposition andchemical vapor deposition techniques) including amorphous as well as diamondfilms. These topics, although of immense interest, are beyond the scope of thischapter. In particular, the cost of the materials currently limits their potentialapplication in the field of environmental remediation. However, we recognize theimportance of these new areas although the potential application of fullerenesand carbon nanotubes has not yet been fully realized.

The majority of engineered carbons discussed in this chapter have graphitic ordisordered graphitic microstructures. Also, for most engineered carbon materials theoriginating precursor is organic and the materials arise from heat treatment of theprecursor in inert atmospheres (carbonization). A selection of technically importantcarbons arising from solid, liquid, and gaseous organic precursors is presented inTable 3.

Engineered carbons are the product of the carbonization process, i.e., ofpyrolysis of the carbon-containing material, conducted in the absence of air (usuallyin a nitrogen, argon, or similar oxygen-free atmosphere). Depending on the choice

of the starting precursor material, the char obtained as a result of the carbonizationprocess may be virtually inactive as regards adsorption, with a specific surfacearea of several square meters in gram. On the other hand, we have shown that car-bonization of porous polyaddition=polycondensation polymers results in retentionof the porosity of the original polymer structure in the final char and hence in reten-tion of the highly developed surface area. The inactive char obtained in the previousinstance requires activation to convert the char into an adsorbent of high porosity(pore volume at least 0.2 cm3=g, although values as high as 1 cm3=ghave been recorded). During the activation process, the carbonaceous material issubjected to selective thermal treatment under suitable conditions, which results

in gasification of loosely bound pyrolysis products and thereby in the opening upof innumerable pores, fissures, and cracks. The surface area occupied by poresper unit mass of the material increases significantly. The activation process maybe carried out by impregnation with chemicals (primarily dehydrating agents suchas zinc chloride and phosphoric acid) or activation by steam, air, or carbon dioxide.The content of volatile substances (dehydration products and cyclization products

The purge gas is usually passed through an oxygen trap to remove any traces of oxygen. Also, the

high-temperature furnaces are sealed to prevent any air from leaking into the furnace chamber.


7/25/2019 DK1300_ch01

6/84

such as polyphenolic-type compounds) in these materials is an important parameterof their susceptibility to undergo activation. Materials containing a small propor-tion of such compounds will not be amenable to significant surface activation.

Activation with air is selective but involves the risk of external surface burn-off of the carbon particles. Consequently, steam and carbon dioxide are thepreferred activation agents. Activation with carbon dioxide and steam requires high

temperatures to increase the rate of the gasification reaction (temperature-dependent).Thus, temperatures in the range 8001000C are employed. Carbonization andactivation are carried out in specialized apparatus such as fluidized bed ovens,rotary ovens, and multilevel ovens. On a laboratory scale, the maintenance ofrotary ovens precludes their use. Blockage of the rotary seals with pyrolysisproducts requires frequent cleaning of the system. Shallow-bed alumina boatsin horizontal sealed-tube furnaces or muffle furnaces were employed in ourlaboratory to carry out carbonization. During activation, the mass of the carbo-naceous material decreases significantly (decrease in the yield of final carbon).

A consequence of this is an increase in the porosity of the carbon.

Table 3 Starting Precursors for Carbons

Primary Secondary Example of carbons

Hydrocarbongases

Pyrocarbons, carbon blacks,vapor-grown carbon fibers,matrix carbona

Petroleum-derived Delayed coke, calcined cokePetroleum pitch Needle coke, carbon fibers,

binder and matrix carbona

Mesophase pitch Mesocarbon microbeads,carbon fibers

Coals Semi-coke, calcined cokeCoal chars Activated carbonsCoal tar pitch Premium cokes, carbon fibers,

binder and matrix carbonsa

Mesophase pitch Mesocarbon microbeads,carbon fibers

Polymers Polyaddition andpolycondensationtypes of resins

PAN-based carbon fibers;glassy carbons, binder andmatrix carbons, carbon beads;graphite films and monoliths

Biomass Coconut shells, apricot,olive and peachstones, pits

Activated carbons

aPrecursor as a binder in granular carbons and graphites and in carboncarbon composites.

Source: Adapted from Ref. 10.


7/25/2019 DK1300_ch01

7/84

An informed choice regarding the suitability of the engineered carbon for aparticular application must be based on the following considerations: size andcomposition of the granules, specific surface area of the pores, pore volume, andpore volume distribution along with the character and chemical structure of thecarbon surface. Carbon granules are available as irregular shapes, spherical beads(polymer-derived carbons), and cylinders (usually extruded granules). The proper-ties of the carbon required for a particular application may be tailored by choice ofthe carbon precursor and the activation procedure (including time and conditionsof activation). The total number of pores and their volume distribution dependmainly on the nature of the raw material used and on the physicochemical para-meters of the carbonization and activation process.

Commercially, raw materials used to manufacture active carbons includewood (sawdust), charcoal, peat, peat coke, and certain types of hard and brown coal

(including the semi-coke of brown coal). Coconut shells are used to prepare activecarbons with a large volume of small pores (micropores); these exhibit a highadsorption capacity. Increasingly, polymer-derived carbons are attracting attentionin the literature [11]. Their principal advantage lies in the ability to control the porestructure of the carbon precursor during the polymer synthesis process. A reviewarticle by Kyotani [12] is recommended reading. The pore structure present inthe polymer precursor may then be retained in the final pyrolyzed carbon. Also,the final carbons are free of additional impurities.

An important property of carbons relates to their prospective use in many

fields where other absorbents would be unacceptable, e.g., in aggressive media. Thisproperty is related to some structural features of carbon and depends primarily onthe carbon precursor, the heat treatment temperature, and the presence of elementswithin the carbon matrix. Acid stability of carbons increases with increase in heattreatment temperatures as the proportion of the more stable bonds increases, whilethe more perfect carbon structure prevents acid diffusion into the carbon matrix.Impurities in the carbon precursor may produce a marked effect on the formationof ordered carbon structures; the presence of additives has been found to influencethe chemical stability of carbons with respect to the action of some corrosive liquids.

Engineered carbonaceous adsorbents adsorb components from liquid or

gaseous media by physical forces or chemical interaction via an ion-exchange=coordination mechanism. The porous structure and the chemical nature of thecarbon surface are significantly related to its crystalline constitution. The graphite-like microcrystalline structure is the main building block of engineered carbonadsorbents. The crystal structure of graphite consists of parallel layers of condensedregular hexagonal rings spaced 0.335 nm apart. The in-plane CC distanceis 0.142 nm, intermediate between Csp3Csp3 and Csp2Csp2 bond lengths,0.153 nm and 0.132 nm, respectively. The elementary microcrystalline structurein active carbons resembles that of graphite. The large difference between the

in-plane CC distance, 0.142 nm, and the interlayer distance of 0.335 nm results


7/25/2019 DK1300_ch01

8/84

from differences in the types of chemical bonding. The large interlayer spacing maylead us to believe that the contribution from p-bond overlap is negligible, andthe usual assumption has been that interlayer potentials are of the van der Waalstype [10]. Tomanek et al. [13] provide evidence (by obtaining scanning probemicroscopic images of a graphite surface) that there may be some p-orbital inter-action between planes.

The early use of X-ray diffraction and subsequently electron and neutrondiffraction, high resolution electron microscopy, and atomic force and scanningtunneling microscopy enabled researchers to improve their understanding of thestructure of engineered carbons and graphites. Concepts of graphitizing andnongraphitizing carbons arising from work by Franklin [14] have been used todistinguish low-temperature carbons. Graphitizing carbons develop a three-dimen-sional graphitic structure at heat treatment temperatures above 2000C, whereas

nongraphitizing carbons cannot be transformed into a graphitic carbon solely byheat treatment (in an inert atmosphere) at temperatures up to 3500 K.

The principal difference in the development of graphitizing and nongraphi-tizing carbons is that in the former, crystallinity is established due to layeringtogether of high molecular weight polyaromatic hydrocarbon-type molecules. Thisformation of a liquid crystal phase occurs from within the fluid phase duringcarbonization. The pregraphitic structure arising due to the coalesced mesophasecan be developed into graphite upon high-temperature treatment. Pyrolysis ofaromatic petroleum, coal tar pitches, and model polyaromatic compounds such

as polyvinyl chloride (C2H2Cl2) carbonize to graphitable, anisotropic cokes [15].Carbonization of precursors of nongraphitizing carbons does not involve the forma-tion of the mesophase. Cross-linked polymeric precursors such as sulfonated styrenedivinylbenzene resins or phenolic resins or other nongraphitizing precursors thatundergo cross-linking reactions during the early stages of carbonization, e.g.,polyacrylonitrile resins, are parent materials for nongraphitizing carbons.

The porosity in carbons arises from imperfections of the graphitic lamellaebonded together to create the three-dimensional network. Consequently, thepacking density of highly porous carbons with isotropic physical properties canbe low, 1.0g=cm3. The greater the order of the lamellae stacked parallel to each

other over relatively larger dimensions of micrometers, the higher the density,approaching 2.0 g=cm3 (the density of pure graphite is 2.267 g=cm3). Such materi-als have little or no useful porosity with sizes comparable to molecular dimensions,i.e., nanometers.

Pore structure imparts widely different properties to engineered carbons.For example, the introduction of micro- and mesopores is important because thisdetermines the ability to adsorb large amounts of various types of molecules fromgaseous or liquid streams. The control of pore size and pore size distributionremains an important technological goal given the complexity of the carbon struc-

ture. The parent materials used as precursors in our work have included coconut


7/25/2019 DK1300_ch01

9/84

shell, wood, coal, fruit stones (e.g. apricot stones), sawdust, and polyaddition andpolycondensation types of resins. Other variables include the diversity of feedstock(especially for natural materials), carbonization conditions including rates of heat-ing, final heat treatment temperature, soak time, and ambient gases. This hasresulted in the synthesis of a wide range of porous carbons possessing differentphysicochemical properties.

Naturally occurring materials are three-dimensional polymeric networkscomposed of lignin and cellulose. During pyrolysis (heat treatment of material to< 700C) major changes in structure occur, whereas at higher carbonization tem-peratures (> 700C) development of the aromatic carbon occurs. The cellulosicstructure loses small molecules such as water and carbon dioxide as volatiles on pyr-olysis, together with a complex mixture of aliphatic acids, carbonyls, alcohols, etc.The pyrolysis products do not evolve at a single decomposition temperature but

over a range of temperatures. Removal of the small volatile molecules from themacromolecular network results in structural vacancies. The chemically reactive lat-tice quickly recombines around the vacancies, consequently, no melting is observed.

A new lattice with higher C=H and C=O ratios results, because there is preferentialloss of hydrogen and oxygen. A schematic representation of a cellulosic-type precur-sor and an intermediate structure between cellulose and an aromatic carbon adaptedfrom Byrne and March [16] is shown inFig. 1.

The newly created lattice, richer in carbon than the original cellulosiclattice, possesses considerable strain energy and is not in thermodynamic equili-

brium. The more stable state is the graphite-like structure that can be achievedby increasing the heat treatment temperature. At higher temperatures, furthercyclization reactions take place as the carbon atoms readjust and approximate tosix-membered ring systems. With increasing heat treatment temperature, a macro-molecular structure composed of small clusters of six-membered carbon ring sys-tems in defective nonlamellar configuration arises. At higher temperatures,further removal of hydrogen, oxygen, nitrogen, sulfur, and phosphorus occurs,resulting in a more carbonaceous and aromatic network. The defective nature ofthe clusters of ring systems and the random bonding together of these clustersresults in a porous open network structure in which the density of the carbon is con-

siderably less than in graphite. The spaces between the lamellar clusters bondedtogether constitute the microporosity of the carbons.Figure 2presents surface areaand pore size distribution of a porous carbon, designated KAU ini, derived fromapricot stones. The measured BET surface area of this carbon is about1800 m2=g, and its pore volume is 0.75 cm3=g.

The macromolecular structure of the precursor impinges on the structure ofthe final carbon; i.e., different natural precursors have different macromolecularnetworks. The pyrolysis reactions resulting in release of volatile components andrearrangement of the macromolecular structure into isotropic lamellar networks

create different carbon lattices. Different matrices evolve depending on the number


7/25/2019 DK1300_ch01

10/84

Figure 1 (a) A cellulosic-type precursor. (b) An intermediate structure between celluloseand an aromatic carbon.

Figure 2 Pore structure of KAU carbon.


7/25/2019 DK1300_ch01

11/84

and strength of bonding of the lamellar clusters and the space between them. Theresulting carbons therefore display different types of adsorption behavior.

We have carried out work on carbonization of unfunctionalized and function-alized ion-exchange resins in an effort to control the precursor properties, inclu-ding the elemental composition, degree of cross-linking, and macromolecularnetwork structure. Precursors used include styrene=divinylbenzene copolymers,polyacrylonitrile=divinylbenzene copolymers, and other polyaddition and polycon-densation types of resins (see Fig. 3 for precursor structures). Our approach has

been to prepare or buy resins with well-characterized structure that allow us tocontrol the design of the pore structure of the resulting carbons.In a recent study in our laboratory, we carried out carbonization of a sulfo-

nated styrene=divinylbenzene copolymer (designated MN500HS). The decompo-sition ratio of the cation-exchange resin after pyrolysis was found to be only50wt%, whereas that of the unfunctionalized styrene=divinylbenzene copolymerwas about 90 wt%. The functional sulfonic acid groups were shown from X-rayphotoelectron spectroscopic (XPS) studies to convert to sulfonyl and sulfur bridgesbetween the base polymers during pyrolysis (Fig. 4shows a hypothetical transforma-tion occurring during carbonization). This resulted in thermal stabilization of the

base polymer.Thermogravimetric analysis (TGA) of MN500HS was carried out following

the temperature profile: 20800C at a heating rate of 0.5 K=min and isothermalheating at 400C and 800C for 1h. The TGA curve of MN500HS showed a19% weight loss between about 200C and 400C. At a temperature close to400C, rapid thermal decomposition of the MN500HS resin started to occur, withanother 12%loss of weight between 400C and 450C. Increasing the temperatureto beyond 450C resulted in a further steady loss in weight, with a final yield of 60%.

Pyrolysis of MN500HS was carried out to investigate the influence of sulfo-

nic groups on the thermal stability of the resin during the carbonization process.

Figure 3 (a) Cross-linked polystyrene with sulfonic acid groups. (b) Polyacrylonitrile=divinylbenzene copolymer.


7/25/2019 DK1300_ch01

12/84

Such an event would enable opening of the three-dimensional structure of thepolymer, which may result in this structure being retained in the final carbon.

Thermal treatment of MN500HS at discrete heat treatment temperatures(HTT) of 350C, 400C, and 800C was carried out, and sulfur elemental analysisof the samples is reported in Table 4. Pyrolysis of MN500HS showed that rapidthermal decomposition of the resin started to occur at a temperature close to350C, resulting in a 34% weight loss between 200C and 350C. Increasingthe temperature beyond 400C resulted in a further loss of weight. A final yield

Table 4 Thermal Stability and Elemental Analysis Data for Furnace-Treated Samples

Heat treatmenttemperature Loss of weight

Sulfur content

Sample (C) (wt%) (mmol=g) Wt%

MN500HS 3.2 10.1MN500HS-HT350 350 34 1.2 3.7MN500HS-HT400 400 n.a. 1.0 3.1MN500HS-C1 800 45 0.5 1.5

Figure 4 Hypothesis of changes in chemical structure of a sulfonated styrene divinyl-benzene resin (MN500HS) that occur during carbonization.


7/25/2019 DK1300_ch01

13/84

of about 55 wt% was obtained. Carbonization of nonsulfonated styrene=divinylbenzene (DVB) polymer resulted in a significant weight loss (around90%) as reported by Matsuda and Funabashi [17].

The sulfur present within the resin was converted into SO2 gas at about300C (the furnace offgases were passed through a mass spectrometer). However,about 37 wt% of the original sulfur present in the resin remained in the residueat 350C (sample designated as MN500HS-HT350), and 15 wt% of the originalafter an HTT of 800C.

Sulfonation of the hypercross-linked sulfonated styrene=DVB copolymerfollowed by carbonization yields a significantly higher weight yield of final carbon.By comparison, the unfunctionalized hypercross-linked styrene=DVB resin yieldeda char that corresponded to less than 10% of the original polymer weight.

Carbonization of the sulfonated polymer results in retention of the bimodal

pore structure (micro- and mesoporosity) in the final mesoporous carbonMN500HS-C1 (see Fig. 5). This is also reflected in the high carbon yield (lowerweight loss) for the pyrolyzed sulfonated polymer as discussed above. Carbo-nization of MN500HS results in a moderate increase in the measured BET surfacearea value for sample MN500HS-C1. The density functional theory (DFT) surfacearea attributed to pores greater than 20 A remains unchanged at 12 m2=g; however,

Figure 5 Pore stucture of () polymer MN500HS and () carbon MN500HS-C1.


7/25/2019 DK1300_ch01

14/84

in the carbon MN500HS-C1, most of the contribution comes from pores largerthan 200 A .

The valence of the sulfur in the thermally treated polymer samples discussedabove was investigated using XPS. This was done to elucidate how the sulfonic acidgroups of the original polymer stabilize the structure so that the final carbon retainsthe pore structure of the polymeric precursor.

The change in the sulfur chemistry may be monitored by S 2pXPS spectra.Figures 6adshow the XPS spectral data. Although the total sulfur content of thesurface decreases as the HTT increases, differences in sulfur valence were observedat different HTTs.

The binding energy of the sulfur 2p peak was 168.8 eV for the polymersample MN500HS before pyrolysis, which equaled the value for benzenesulfonicacid (C6H5SO3H) analyzed as a reference (see Fig. 6a). A new peak at 164.1 eV

was detected in the spectra after pyrolysis at 350

C (see Fig. 6b). This was similarto the value of 164.0 eV for phenyl sulfide polymer, (C6H5S)n. The spectra forMN500HS-HT350 also contained a small peak at a binding energy similar to thatof the original S species. A hypothesized structural transition is shown in Fig. 4.Diphenyl sulfone, (C6H5)2SO2, was also analyzed as a reference for the intermedi-ate species. However, this reference material gave a binding energy very similar tothat of benzenesulfonic acid (168.7 eV), and therefore the intermediate species wasnot identified, although this may be a step in the cascade of transformations occur-ring during the pyrolysis process. Figure 6c suggests that at a pyrolysis temperature

of 400

C none of the original sulfonic acid groups are evident.Apart from the sulfur 2pspectra, surface compositional data for the sampleswere also obtained by XPS (seeTable 5). Sulfur present on the external surface ofthe carbon beads follows a trend similar to that obtained for the bulk sulfur elemen-tal data discussed earlier (seeTable 4). However, XPS surface data show a reductionin sulfur content from 4 at% to about 1 at% as the polymer sample is heated to400C. The S 2pspectra show no oxidized sulfur at 400C (around 1 at% is thelimit of detection of the instrument) and higher, but oxygen is still present inMN500HS-HTT400 and MN500HS-C1. This residual surface oxygen may beattributed to surface oxidation of the carbon upon exposure to the atmosphere when

the samples are removed from the furnace. Within the furnace, the samples are inan inert nitrogen atmosphere.

A series of activated carbon samples were prepared from MN500HS usingCO2 as the activation agent at various temperatures to obtain different levels ofburn-off. These samples were designated MN500HS-C1-(A1A6). Activation ofMN500HS-C1 using carbon dioxide as the activating agent results in a progressiveincrease in burn-off and hence reduction in the yield of the final char (seeTable 6).

Activation of MN500HS-C1 at 800C for 0.5 h resulted in a carbon yieldof 53%, whereas the yield of carbon fell to 44%for the MN500HS-C1-A4 sample

activated for 5 h. Increasing the activation temperature had a significant impact on


7/25/2019 DK1300_ch01

15/84

Figure 6 (a) XPS data for polymer MN500HS sample. (b) XPS data for MN500HS-HT350 sample. (c) XPS data for MN500HS-HT400 sample. (d) XPS data forMN500HS-C1 sample.


7/25/2019 DK1300_ch01

16/84

Figure 6 (continued)


7/25/2019 DK1300_ch01

17/84

the degree of burn-off, with the MN500HS-C1-A6 sample yielding a charcorresponding to 30% of the original starting polymer weight. Surface area valuesmeasured for the activated MN500HS samples show an increase with degree ofburn-off. MN500HS-C1-A1 displays a BET surface area of 515 m2=g compared with1182 m2=g measured for MN500HS-C1-A6. Pore volume in MN500HS-C1-A1was 0.5 cm3=g, and this increased to about 1 cm3=g in MN500HS-C1-A6.

The effect of activation on the mechanical strength of the resultingMN500HS carbons is demonstrated by single-bead crush tests; the results aresummarized in Table 6. MN500HS-C1 shows the highest mechanical strengthvalue (1.7 kg force). Upon activation, the activated carbon samples show an increase

in the BET surface area values as the degree of activation increases; for samplesA1A5, surface area increased from 515 m2=g for sample A1 to 782m2=g forsample A5. The activated carbon MN500HS-C1-A6 sample shows a sudden dropin mechanical strength as the BET surface area value approached 1200 m2=g.

Table 5 Surface Composition of MN500HS Samples During Pyrolysis

Sample

Surface composition excluding H (at %)

S C O

Polymer A 4.2 74.4 21.4HT 350C 1.3 86.5 12.2HT 400C 1.2 89.3 9.5HT 800C 1.1 96.2 2.7

Source: Ref. 20, with permission from Elsevier.

Table 6 Activation Conditions and Selected Physical Properties of Activated Carbons

Sample Activationtemp. (C) Activationtime (h) Yield(wt%)

BET

surface area(m2=g)

Pore

volume(cm3=g) Mechanicalforce (kg)

MN500HS 414 0.37 0.8MN500HS-C1 55 473 0.48 1.7MN500HS-C1-A1 800 0.5 53 515 0.50 1.3MN500HS-C1-A2 800 1 49 535 0.49 1.3MN500HS-C1-A3 800 3 46 618 0.54 1.4MN500HS-C1-A4 800 5 44 698 0.57 1.2MN500HS-C1-A5 850 3 40 782 0.67 1.2MN500HS-C1-A6 900 3 30 1182 1.04 0.7


7/25/2019 DK1300_ch01

18/84

The yield of the carbon was around 30 wt%. These results suggest that mechanicallyrobust activated carbons with BET surface area values in the range of 800 m2=g canbe prepared using sulfonated styrene-divinylbenzene resins.

We also carried out investigations aimed at controlling the microporosity ofcarbons by using ion-exchange resins as carbon precursors. We employed sphericalpolystyrene-based resins with sulfonic acid groups as ion-exchangeable sites. Theresin was converted to the H, Na, Cs, Cu2, Co2, Fe2 forms and thencarbonized at 800900C. Figure 7 shows the surface area distributions of theresultant carbons prepared by carbonizing sulfonic acid resin in different ionicforms. Although the starting resin was the same before cation exchange, carbonsresulting from the resin in various metal forms gave different microporous andmesoporous structures. The average micropore diameter of the carbons varied inthe range of 0.30.5 nm, depending on the type of cation used. This implies that

the micropore size can be tuned by changing the type of cation. From a practicalpoint of view, this method is promising because waste ion-exchange resins can beconverted to microporous molecular sieving carbons [12].

Metal sulfonate groups in the resins decompose at higher temperatures thansulfonic acid groups in H form. Miura and coworkers [18] carried out similar

Figure 7 Incremental pore area distribution of carbons.


7/25/2019 DK1300_ch01

19/84

studies and explained that the metal form of the resin holds a larger amount ofsulfur as metal sulfide and=or other forms even at high heat treatment temperatures.The metal sulfides or remaining metal may play a role as a supporting pillarbetween carbon lamellae. Such pillars may prevent further shrinkage of microporesduring high-temperature treatment.

Whereas micropores are essential for the sorption of small molecules, largerorganic molecules such as dyes, herbicides, and pesticides require the presence ofmesopores. We employed polymers with a bimodal pore size distribution to preparecarbons containing a significant pore volume contribution from mesopores. Weretained the bimodal pore structure of the precursor in the final carbons bythermally stabilizing the polymer precursor, e.g., due to the presence of sulfonicgroups in polystyrene or oxidation of polyacrylonitrile (see Fig. 8).

Oxidation of polyacrylonitrile is crucial to successfully carbonize this

material. The oxidation starts off the cyclization reaction between the nitrogen ofthe nitrile group and the carbon of a neighboring nitrile group; this requires thepresence of a nucleophile stronger than a nitrile nitrogen atom. Oxidation of asmall proportion of nitrile groups to carboxylic groups can be carried out by usingwet oxidation techniques or air at elevated temperatures. Carboxylic groups can inturn initiate cyclization of acrylonitrile at elevated heat treatment temperatures,

Figure 8 Incremental pore volume distribution of acrylonitrile=divinylbenzene copolymers.


7/25/2019 DK1300_ch01

20/84

leading to pyridinic type structures. The effect of the oxidation step on the yield ofthe material obtained by acrylonitrite=divinylbenzene copolymer (AN=DVB) car-bonization is illustrated in Table 7. Despite the presence of a cross-linker (DVB),the carbonization yield of the AN=DVB sample without oxidative thermal stabiliza-tion was only about 10 wt%, whereas when the sample polymer was oxidized at300C a yield of about 45 wt%carbon was obtained. The weight loss at tempera-tures in the range of 250350C can be attributed to dehydration and to loss ofnitrogen due to formation of ammonia and hydrogen cyanide.

III. CHARACTERIZATION OF THE SURFACE

CHEMICAL GROUPS IN ENGINEERED CARBONS

The pore structure and surface area of engineered carbons determines the physicalcharacteristics of the adsorbent. However, the surface chemical structure affectsinteraction with polar and nonpolar molecules due to the presence of chemicallyreactive functional groups. Active sites, e.g., edges, dislocations, and discontinu-ities, determine the reactivity of the carbon surface. Adsorption is a complex inter-play between the chemical and porous surface structures of the carbon. Importantfactors include the nature and relative amounts of surface functional groups, thesurface area and pore size distribution, as well as the characteristics of the adsor-bate molecule, e.g., size and nature of the cation=anion, polarity and chemical

structure of the molecule, molecular and ionic dimensions, etc. The followingdiscussion highlights the importance of precise characterization of the surfaceand structure of engineered carbons in order to tailor adsorbents for specificapplications.

The chemical nature of the surface of porous carbons is strongly dependenton the type, quantity, and bonding of various heteroatoms in the structure, par-ticularly oxygen [19]. The heteroatoms may be distributed at random within thecarbon matrix, concentrated at the exposed surface of carbons, present at disloca-tions in the microcrystalline structure, or a combination of all three. Surface-bound

heteroatoms may be considered analogs of substituted aromatic compounds.

Table 7 Acrylonitrile=DivinylbenzeneDerived Carbon Characteristics

CarbonYield

(%)BET surfacearea (m2 g)

Pore volume(cm3=g)

AN=DVB-oxid250C-C1 35 187 0.39AN=DVB-oxid300C-C1 46 429 0.59AN=DVB-oxid350C-C1 41 439 0.60


7/25/2019 DK1300_ch01

21/84

Heteroatoms present in the core of the carbon matrix may be unreactive due to theirinaccessibility and their mode of combination with local atoms.

Much of our work has focused on the identification and characterization ofoxygen-containing functional groups. We have also developed methods of selec-tively incorporating or destroying particular types of groups in order to improvethe selectivity and=or adsorptive capacity of the engineered carbons with respectto a specific metal ion or group of ions. We have employed a variety of gaseousoxidizing agents such as oxygen, ozone, air, water vapor, carbon dioxide, and nitro-gen oxides to increase the quantity of oxygen on the surface of the carbon. Alter-natively, we have employed oxidizing solutions, e.g., nitric acid, a mixture of nitricand sulfuric acids, and hydrogen peroxide. Some work has been carried outemploying electrochemical oxidation in the presence of electrolytes such as KCl,KNO3, NH4HCO3, and HNO3. The number and distribution of oxygen-

containing surface functional groups and reactivity can differ over a fairly widerange depending on the type of oxidizing agent used and on the experimentalconditions.

A series of carbons oxidized using hot air and nitric acid oxidation treatmentswere prepared. The carbons were derived from apricot stones (designated as KAU)and sulfonated styrene=divinylbenzene copolymer (designated as CKC). Thematerials were used to study the influence of oxidation conditions on the acidicsurface functional group content of the resulting materials. Table 8 presents theconditions of oxidation and some properties of the oxidized carbons. (Note: The

table also presents materials that are discussed later in the chapter.) Preparationof carbons and other details may be found in separate publications [20,50,51].Oxygen surface compounds are usually divided into two main categories:

functional groups that undergo neutralization by bases (having acidic nature)and, conversely, basic groups that can be neutralized by acids. Commonly charac-terized acidic groups are exemplified schematically inFig. 9.

Basic groups are far less well characterized. Researchers have proposed analogsof the chromene- [21] or pyrone-like structures illustrated schematically inFig. 10.Puri [22] and others have argued that there is no independent proof of theexistence of chromene-type structures present on the surface of carbons. Pyrolyzed

carbons behave as polycondensed aromatic hydrocarbons (known to be Lewis bases,i.e., electron donors) and can preferably exchange H ions to OH ions from wateror aqueous solutions [23].

A model of a fragment of oxidized carbon surface illustrating the generalchemical character of the active carbon surface is shown in Fig. 11. The positionof these groups in close proximity to one another may influence their acidity owingto direct interactions with adjacent groups of the same or other type or to electro-static interactions with the lamellar matrix. For example, the acidity increases bymore than an order of magnitude from benzoic to m-phthalic and salicylic acids

(see Fig. 12 and Table 9). We found that surface carboxyl groups of oxidized


7/25/2019 DK1300_ch01

22/84

carbons exhibit even lower acid dissociation constants. This may be related to thefact that the surface groups are connected to a p-conjugated condensed system ofgraphite-like planes. The number of conjugated benzene rings and the positionsof the groups will also influence their acidity.

Owing to the special electrophysical properties of carbons (namely the pre-

sence of delocalized p-electrons that are relatively easily transferred in a conjugated

Table 8 Conditions of Oxidation and Some Properties of the Oxidized CarbonsDiscussed in This Chapter

Carbon

Oxidizing

agent

Exposure

time (h)

Temperature

(

C)

Na capacity

(mmol=g)

KAU-ini Unoxidized N.A. N.A. 0.4KAUN-1.7 HNO3 2 9095 1.7KAUN-2.0 HNO3 5 9095 2.0KAUN-2.6 HNO3 10 9095 2.6KAUN-2.9 HNO3 15 9095 2.9KAUA-0.8 Air 2 410 0.8KAUA-1.3 Air 3 430 1.3KAUA-1.8 Air 5 450 1.8KAU-125 Electrochemical 1 25 1.3

KAU-525 Electrochemical 5 25 3.1F400 Unoxidized 0.1F400(ox) HNO3 15 9095 1.6F400-Ox9h-WWa HNO3 9 9095 2.4F400-Ox9h-AWb HNO3 9 9095 2.5F400-Ox9h-HTc HNO3 9 9095 2.1F400-Ox24h-WW HNO3 24 9095 3.1F400 Aox Air 24 420 2.1BGP unoxidized Unoxidized 0.3BGP-OxII-AWd HNO3 9 9095 2.2

BGP Aox Air 24 420 0.8CKC HNO3 15 9095 2.2CKC-325 Electrochemical 3 25 2.4

N.A. not applicable.aSample was washed with distilled water until no change in the pH could be detected.bSample was washed with 0.1 M NaOH to remove alkali-soluble products during the oxidation reaction

(e.g., humic compounds).cAnother treatment to remove humic acids produced during the oxidation was to heat the samples to

320C for 12 h under vacuum. This removed most of the alkali-leachable material, and no trace of these

substances could be detected when the heat-treated carbon was contacted with akali solution.d

Following the oxidation process, the oxidized carbon was washed with 0.1 M NaOH to remove labilehumic compounds. The samples were then washed with 0.1 M HNO3, then with water until neutral pH

was attained.


7/25/2019 DK1300_ch01

23/84

system of aromatic bonds), the carbonaceous adsorbents possess significant electricalconductivity. The properties of functional groups attached to a carbon surfaceshould diverge considerably in comparison to those of groups attached to a noncon-ductive polymeric matrix (ion-exchange resins). The strengths of real protonogenicsurface functional groups (primarily carboxylic groups) in carbons should be higherthan of those in resins, because the negative charge usually builds up close to the

Figure 9 Principal types of oxygen-containing surface acidic groups: (a) carboxyl,(b) phenolic, (c) quinonic, (d) normal lactone, (e) fluorescein-type lactone, (f) anhydride ori-ginating from carboxyl groups. (Redrawn from Ref. 19.)

Figure 10 Hypothesized chromene groups on basic carbons readily oxidized at roomtemperature in the presence of acid to the corresponding benzopyrylium structures (carbo-nium) with the adsorption of the anion of the acid and liberation of hydrogen peroxide.

(From Ref. 23.)


7/25/2019 DK1300_ch01

24/84

Figure 11 Hypothetical fragment of an oxidized carbon surface. (Adapted from Ref. 19.)

Figure 12 Structures of selected aromatic acids. (From Ref. 26.) (a) 1-Benzenecarboxylicacid, pKa(1) 4.19; (b) 1-hydroxybenzoic acid, pKa2.93 and 12.37; (c) 5-hydroxybenzoic acid,pKa4.08 and 9.78; (d) 4-hydroxybenzoic acid, pKa4.61 and 9.31; (e) 2,4-dihydroxybenzoicacid, pKa3.22; (f) 2,6-dihydroxybenzoic acid, pKa 1.22; (g) naphthalene-1-carboxylic acid,pKa3.6; (h) naphthalene-2-carboxylic acid, pKa4.16; (i) anthracene-1-carboxylic acid, pKa3.68; (j) anthracene-2-carboxylic acid, pKa 4.17; (k) anthracen-9-carboxylic acid, pKa

3.65.


7/25/2019 DK1300_ch01

25/84

surface of the oxidized carbons. This is demonstrated by the fact that cation

exchange on many oxidized carbons, as opposed to typical carboxylic resins, beginsat pH values around 1 [24], i.e., the acidity of carboxylic groups on the oxidizedcarbon surface is 10 times as strong as that of the carboxylic resin. Delocalizationof the 2pelectrons of functional group oxygen into the p-conjugated system resultsin a decrease in the effective negative charge present on the surface oxygen. This willresult in increased mobility of the associated protons, thereby resulting in lower pKavalues. Radovic et al. [25], in their comprehensive review article (recommendedreading) on interactions of inorganic solutes with active carbon surface, highlightthe key molecular features of the edges and basal planes of graphene layers.Delocalized p-electrons may act as Lewis bases in aqueous solution, giving rise to

a positive surface charge (notice the positive zeta potential values at low solutionpH inFig. 13). Graphene layers may also interact with aromatic solutes by ppinteractions.

Figures13 and14show the electrophoretic mobility (zeta potential) versuspH plots for KAU carbons oxidized by air and nitric acid. An important para-meter used to characterize the electrokinetic behavior of a solid=liquid interfaceis the point of zero zeta potential. The pH value at this point is often calledthe isoelectric point (IEP) of the interface. A shift in the pH IEP of KAUini tolower values is observed as the degree of surface oxidation (sorbent capacity)

increases (see Table 10). The isoelectric point for some acid-oxidized carbons

Table 9 Dissociation Constants for Selected AromaticAcids in Water

Acid pKa1 pKa2

o-Phthalic 2.89 5.51m-Phthalic 3.54 4.60

p-Phthalic 3.51 4.82Benzoic 4.19 Salicylic 2.93 12.37m-Oxybenzoic 4.08 9.78

p-Oxybenzoic 4.61 9.312,4-Dioxybenzoic 3.22 2,6-Dioxybenzoic 1.22 Naphthalene-2-carboxylic acid 4.16

Naphthalene-1-carboxylic acid 3.69 Anthracene-9-carboxylic acid 3.65 Anthracene-1-carboxylic acid 3.68 Anthracene-2-carboxylic acid 4.17

Source: Ref. 26.


7/25/2019 DK1300_ch01

26/84

Figure 13 Electrophoretic mobility curves for KAU carbons oxidized by air. (From Ref.20, with permission from Elsevier.)

Figure 14 Electrophoretic mobility curves for KAU carbons oxidized by nitric acid.

(From Ref. 20, with permission from Elsevier.)


7/25/2019 DK1300_ch01

27/84

was found to be as low as 1.1 (see extrapolated data for KAU-2.9 in Table 10). Arather low isoelectric point (pHIEP2.5) for the KAUini carbon was detected (seeFig. 13). This indicates that exposure of carbon to atmosphere (during storage)results in some surface oxidations of the carbon granules.

All carbons are the so-called L-carbons [21] (i.e., pHIEP< 7). The zeta poten-tial curves for different acid-oxidized carbons fall steeply with increasing pH untilpH 5 and then start to level off. Dissociation of relatively strong carboxylic surfacegroups is probably responsible for this effect. Indeed, the dissociation of this type ofgroup occurs at pH values between 2 and 6, thereby enhancing the negativity of the

surface (refer to Table 10 for dissociation constants). In air-oxidized carbons,weaker surface functional groups are more prevalent than in acid-oxidized carbons,as discussed earlier. Their dissociation begins at higher pH values (above 6). There-fore, their zeta potential curves change more gradually in comparison to those ofnitric acidoxidized samples.

The quantity of oxygen combined with the carbon surface in the form offunctional groups not only increases via oxidation, it can also be decreased as aresult of thermal decomposition in an inert atmosphere. We have shown that thedistribution of oxygenated functional groups may be changed by selectively destroy-ing thermally unstable carboxylate groups at temperatures of 400700C to yield

CO2, CO, H2O, or H2.The concept of oxygen present on the surface of carbon as organic func-

tional groups suggested by Garten and Weiss [21] was further developed in the1960s by Boehm [27] and Donnet [28]. Evidence regarding the existence of differ-ent types of functional groups was obtained by means of organic chemical detectionand titration experiments. More recently, the study of surface groups has also beenundertaken by spectroscopic methods, chiefly Fourier transform infrared (FT-IR),

X-ray photoelectron spectroscopy (XPS), and Nulcear magnetic resonance spectro-scopy (NMR) (1113). We have employed a number of techniques (Boehms

titration, pH titration, and zeta potential measurements) to characterize the changes

Table 10 Electrochemical Properties and Dissociation Constants of the Adsorbents

Sorbent PZC IEPa PZC IEP pKa1 pKa2 pKa3

KAUini 9.9 2.5 7.4 9.97KAUA-0.8 6.2 2.3 3.9 9 10.7KAUA-1.3 3.7 1.8 1.9 6 8.5 10.85KAUA-1.8 3 1.7 1.3 4.7 7.0 9.8KAUN-1.7 3.1 1.5 1.6 3.6 6.5 9.7KAUN-2.0 2.8 1.4 1.4 4 6.9 10.2KAUN-2.6 2.5 1.3

a 1.2 3 6.75 9.5KAUN-2.9 2.1 1.1

a 1 4.1 6.5 10

aIEP values were obtained by extrapolation.


7/25/2019 DK1300_ch01

28/84

in carbon surface chemistry that take place during different oxidation treatments.This method (due to Boehm [27]) is based on the fact that surface acidic groupsundergo a differential neutralization depending on their acid dissociation pKa.Thus, carboxyl groups (with low pKa) are neutralized by sodium hydrogen carbo-nate; carboxyl and lactone groups are neutralized by sodium carbonate; whereascarboxyl, lactone and phenolic groups are neutralized by sodium hydroxide. Differ-ential titration results presented in Table 11 suggest that all the carbonaceoussorbents possess oxygen functionalities in the form of carboxylic, lactonic, andphenolic groups. The concentration of carboxyl-type groups substantially increases

after oxidation. In comparison with other oxygen-containing groups on the carbonsurface, e.g., phenolic, lactonic, quinone, carboxyl groups have the lowest pKavalues. Thus, the number of carboxyl-type groups neutralized by NaHCO3corre-sponds to about 17% of the total number of acidic groups neutralized by NaOH(in the case of untreated active carbon KAUini), whereas the concentration ofcarboxyl-type groups corresponds to almost 50% for KAUN-2.9 after nitric acidoxidation (see Table 11).

Phenolic groups are more prevalent at low degrees of surface oxidation (this istrue for both acid- and air-oxidized samples; compare KAUN-1.7 with KAUN-2.9

and KAUA-0.8 with KAUA-1.8 in Table 11). The total acidity can be increased by a

Table 11 Concentration and Distribution of Surface Functional Groups in Carbons

CarbonCarboxylic groups

(meq=g)Lactones(meq=g)

Phenolicgroups (meq=g)

Total acidicgroupsa (meq=g)

KAUini 0.078 0.000 0.370 0.448Percentb 17.41 0 82.59 100KAUN-1.7 0.842 0.196 0.664 1.701Percentb 49.5 11.52 39.04 100KAUN-2.0 0.959 0.393 0.702 2.054Percentb 46.69 19.13 34.18 100KAUN-2.6 1.275 0.548 0.781 2.604Percentb 48.96 21.04 29.99 100KAUN-2.9 1.430 0.650 0.795 2.875Percentb 49.74 22.61 27.65 100

KAUA-0.8 0.130 0.320 0.362 0.811Percentb 16.03 39.46 44.64 100KAUA-1.3 0.370 0.479 0.457 1.306Percentb 28.33 36.68 34.99 100KAUA-1.8 0.540 0.560 0.680 1.779Percentb 30.35 31.48 38.22 100

aNeutralized by NaOH.bPercent in comparison to the total acid capacity (as neutralized by NaOH).


7/25/2019 DK1300_ch01

29/84

factor of 5 compared to that of the original KAUini for carbons oxidized with nitricacid. The increase in total acid capacity is about twice that of the starting materialfor the most severely air-oxidized carbon. Progressive oxidation of carbon KAUwith nitric acid creates a greater quantity of relatively strong carboxylic surfacegroups (KAUini

7/25/2019 DK1300_ch01

30/84

distinguishing difference between the carbon samples oxidized in hot air and thoseoxidized in nitric acid. The titration curves of the former set of samples display lessalkali neutralization capability between the crossover point (point of zero charge;pHPZC) and a pH around 6. Carboxyl groups are primarily responsible for Na

uptake at these pH values. The lower Na uptake value observed for air-oxidizedcarbons can be attributed to the decomposition of carboxyl groups during thehigh-temperature oxidation process. The gradient of the proton-binding curvesshows a shallow decay over the pH range 36. The curves for acid-oxidized samplesare steeper in this pH range as a result of their greater ion-exchange capacity. Mironov

and Taushkanov [29] suggested that carboxylic groups in carbons dissociate in thepH interval 36 and phenolic groups dissociate above that pH range. This is inagreement with our pH titration results. The ion exchange on the carbon samplesbegins at lower pH values (pH 2 and lower) [30,31].

As pH increases, weaker functional groups progressively participate in theion-exchange process. The crossover point with the pH axis on the titration curvesis the point where anion and cation exchange are in equilibrium. This point is con-sidered to be a point of zero charge (PZC). As the degree of oxidation increases, thecrossover point occurs at lower pH values (seeTable 10). Thus, the crossover point

for KAUini is located at about pH 10 whereas for KAUN-2.9 it is at about pH 2.

Figure 16 Proton-binding curves for KAU carbons oxidized with nitric acid.


7/25/2019 DK1300_ch01

31/84

For all carbons, pHIEP< pHPZCdue to preferential (diffusion-controlled) ambientairinduced surface oxidation (e.g., carboxyl group generation) on the externalsurfaces of carbon particles [25]. The smaller the difference between pHIEP andpHPZC, the more homogeneous the distribution of surface functional groups. Thisis illustrated inTable 10; shorter oxidation times (both in concentrated nitric acidand in air) introduce acidic functional groups primarily on the external surface ofthe carbon particles. This phenomenon is dramatically illustrated by the differencebetween pHIEP and pHPZC for KAUini (7.4 pH units).

All the carbons that were studied exhibit surface protonogenic groups withrather different dissociation constants that vary between 102 and 1012. Thesechanges in surface acidity suggest that cation exchange on active carbons beginsat lower pH values compared to conventional carboxyl-type resins with pKavaluesbetween 3 and 5.

It is difficult to obtain a precise description of the nature and chemical make-up of the functional groups solely on the basis of these methods. Spectroscopictechniques such as FT-IR and XPS can provide independent evidence regardingthe chemical constitution of the surface groups present in engineered carbons.Differences in thermal stability of particular kinds of surface functional groupsdue to the varying energies necessary to split particular carbon bonds have beenemployed in our laboratory to elucidate the functional groups present on the carbonsurface. CC bonds between skeletal carbon and carboxylic groups require a smallerenergy to split the bond compared to skeletal carbon quinone or phenolic group

bonds. When a carbon sample is heated in an inert atmosphere, particular surfacecompounds decompose at different characteristic temperatures, yielding thefollowing products [19]:

1. Carbon dioxide from decomposition of carboxylic and lactone groups, inthe range of approximately 200C to 700800C

2. Carbon monoxide from decomposition of quinone, phenol, and ethergroups, in the range of approximately 500C to about 1000C

3. Water from decomposition of phenolic groups, in the range from200300C to 400500C

4. Molecular hydrogen from recombination of hydrogen atoms liberatedas a result of splitting of CH and OH bonds, in the range above500700C

A combination of evolved gas analysis using mass spectrometry or gaschromatography and conventional thermogravimetric analysis provides further infor-mation on the functional groups. Unfortunately, interpretation of the results is ratherdifficult. Interaction of groups with each other or with other groups in the vicinityleads to multistep thermal decomposition, so the limits of the temperature ranges

characteristic of particular processes become rather diffuse.


7/25/2019 DK1300_ch01

32/84

IV. ACTIVATED CARBON FIBERS AND

WOVEN CLOTHS

We also investigated the properties of activated carbon fibers in the form of wovencloths (designated ACC) for potential applications in the field of liquid- and gas-phase separations. ACC can be prepared by carbonization of nitrogen-containingpolymers, e.g., polyacrylonitrile fibers [32]. Natural cellulosic precursors may alsobe used. Active carbon fiberbased adsorbents possessing surface areas of over1000m2=g have been prepared and evaluated. An activated carbon fiber cloth,TC-66 C supplied by KoTHmex (Taiwan), was recently evaluated in our labora-tory. The specifications of the material (as supplied by KoTHmex) are given inTable 12.

Our work with activated carbon fibers has involved surface modification of

TC-66 C. We have employed low-temperature electrochemical oxidation to modifythe surface of the carbon fibers. The oxidation reactions are mostly anode-mediated;thus the carbon fibers undergo electrochemical modification by serving as the anode[32,33]. Oxidation of fibers was brought about by the evolution of oxygen andchlorine at the anode. The electrolyte composition (usually KNO3, NH4HCO3,HNO3, or KCl) may be varied; current density (usually chosen to be within0.510 A=m2) and the potential 220 V along with the duration of the processare the variables that may be tuned to obtain materials with desirable surfacefunctionality.

Active carbon fibers may selectively adsorb various components from liquidor gaseous media by physical forces (adsorbents) or chemical interactions via anion-exchange and=or complexation mechanism (carbon fiber ion exchangers). Car-bon fiber ion exchangers with a variety of acidic and basic surface functional groupsincluding carboxylic, sulfuric, phosphate, and amino groups can be prepared. Inprinciple, oxidation employing hot air, nitric acid, sulfuric acid, phosphoric acid,

Table 12 Specification of Activated Carbon Fibers (KoTHmexTC-66 C)

Specification TC-66 C

Surface area (m2=g) 10001100Total pore volume (cm3=g) 0.50.6

Average pore diam. (A ) 1920Fiber weight (g=m2) 95105Fiber thickness (mm) 0.40.5Fiber width (cm) 1002Texture PlainDecomposition temp (C) > 500


7/25/2019 DK1300_ch01

33/84

and electrochemical techniques, etc., introduce cation-exchange functional groupsin a manner similar to those discussed for granular materials. Spectroscopic dataobtained in our laboratory combined with potentiometric titration studies providestrong evidence of the presence of various acid groups that are expected to givecarbon fibers weak acidic polyfunctional ion-exchange properties.

KoTHmex TC-66 C was modified using nitric acid, ozone, and electro-chemical oxidation [33] to enhance cation sorption capacity. BET surface areameasurement for the unoxidized sample gave a value of 973 m2=g. This was reducedafter 3 h of nitric acid oxidation to 730 m2=g (sample designated as TC-66-C--acid-ox) and 627 m2=g for the electrochemically oxidized sample (sample desig-

nated as TC-66-C-elec-ox). The reduction in surface area can be attributed toerosion and blockage of pores by degradation products produced during chemicalreaction (for changes in pore structure due to surface oxidation see Fig. 17).

The concentrations and types of oxygen-containing groups on the surface ofunoxidized carbon fiber TC-66 C and the series of oxidized samples are reported inTable 13. The concentration of carboxyl groups increased from about 0.4 mmol=gin TC-66-C-unox to 2.4 mmol=g in TC-66-C-elec-ox. The enhancement ofoxygen-containing groups in particular, carboxyl groups will significantly increasethe cation ion-exchange capacity at near-neutral pH because these groups have a

pKaaround 3 [29].

Figure 17 Pore size distribution of TC-66 C activated carbon fiber samples.


7/25/2019 DK1300_ch01

34/84

In terms of total concentration of acidic surface functional groups (as titratedby NaOH), oxidized samples showed the following trend: electrochemical treat-ment > ozone treatment > acid treatment. Electrochemical oxidation was aneffective technique to incorporate carboxyl-type acidic groups on the surface ofactive carbon fiber. The point of zero charge for TC-66-C-unox changed from

4.2 to 2.8 for TC-66-C-ozone-ox, 2.4 for TC-66-C-acid-ox, and 2.3 for TC-66-C-elec-ox samples. The surfaces of activated carbon fibers show a net negativecharge (i.e., are cation exchangers) at pH values greater than the point of zero chargeand a net positive charge (i.e., are anion exchangers) at pH values less than the pointof zero charge. The characteristic feature of the oxidized carbons is the wide range ofionization constants within each type of group (in particular, carboxyl and phenolicgroups). This may be due to different energy states on the surface where identicalfunctional groups are situated [32]. The ionic groups introduced have differentchemical reactivities that depend on various factors such as whether they arepositioned at the edges or on the planes of the crystalline structures, included in

aromatic or aliphatic structures, the spacing between these functional groups, orwhether they happen to be in the pores with different adsorption potentials [32].

The ion-exchange capacity of ion exchangers based on carbon fibers dependson a wide variety of variables including the type of carbon fiber, the treatmentmethod, and choice of treatment conditions, including type of oxidizing agent,length of treatment, temperature of treatment, etc. For example, as the temperatureof nitric acid solution increases, an increase in the concentration of the functionalgroups responsible for cation exchange is observed. Prolonged duration of oxidationmay result in significant structural changes that may translate to a decrease in

strength of the fiber core. Destruction of the surface layers also takes place, and

Table 13 Oxygen-Containing Groups on As-Received and Modified Carbon FiberTC-66 C

TC-66C

Carboxylic groups

(mmol=g)

Lactones

(mmol=g)

Phenolic groups

(mmol=g)

Total acidic groups

(mmol=g)

Unox 0.4 0.4 0.2 1.0Percenta 40 40 20

Acid-ox 0.8 1.6 0.5 2.9Percenta 28 55 17Ozone-ox 1.5 2.5 1.3 5.3

28 47 25Elec-ox 2.4 2.2 1.6 6.2Percenta 39 35 26

a

Percent in comparison to the total acid capacity (as neutralized by NaOH).Source: Ref. 34.


7/25/2019 DK1300_ch01

35/84

flaws in the fiber volume are exposed. This makes the fibers disintegrate into minorcomponents in the form of longitudinal strands of fibrils.

The presence of bound nitrogen in the composition of the fiber (arising fromthe polymer precursor) gives the material anion-exchange capacity. Carbonizationof nitrogen-containing organic polymer fibers, e.g., polyacrylonitrile, results inthe formation of amino groups (CNH2) and imino groups (CNH) involvedin anion exchange [35].

V. SORPTION OF TRACE METALS ONTO

ACTIVATED CARBON

Previous publications have discussed the adsorption of heavy metal ions fromaqueous solutions by carbonaceous materials [3640]. This topic is of great practicalinterest not only in the field of water treatment and metal removal but also in thepreparation of carbon-supported catalysts where the catalyst precursor is initiallydissolved in water prior to sorption onto the carbon support. It is widely acceptedthat surface acidic functional groups are responsible for metal ion binding [41].Chemical oxidation is commonly used to introduce these functional groups ontothe surface of carbons. The sorptive capacity and selectivity of oxidized carbons varyfor different metal ions, and higher valence metal ions are usually preferred to thoseof lower valence [42,43]. It has also been observed that selectivity differs even withina series of metals with the same valence. Among divalent metals, Cu2 is generally

the most preferred ion. However, the reasons for the higher affinity of oxidizedcarbons toward this particular metal ion have not been clearly identified.

We have studied a series of unoxidized and oxidized activated carbons for theadsorption of d-block metal ions such as Cu2, Ni2, Co2, Zn2, Mn2, andCd2 from aqueous solutions. A list of the source adsorbent materials is given inTable 14.

Table 14 Materials Used in Metal Adsorption Studies (Unoxidized and Oxidized)

Designation Detail

Granular adsorbentsChemviron F400 Coal-based activated carbonKAU Apricot stone-based activated carbonCKC Polystyrene divinylbenzenebased activated carbonCeca BGP Wood-based activated carbon

Activated carbon clothKoTHmex TC-66 C Polyacylonitrile-based activated carbon woven fiber

Ion-exchange resinPurolite C-104 Polymethyl methacrylate copolymer


7/25/2019 DK1300_ch01

36/84

The diversity of acidbase surface properties of adsorbents is reported inTables15 and16. All granular carbons evaluated in the study exhibited a negativesurface charge in the pH range studied (pH < 5). As inferred from the isoelectricpoint (IEP) and point of zero charge (PZC) values, the extent of surface chargevaries with the degree of surface oxidation. Carbons with greater surface oxidationpossess a greater surface negative charge. Dissociation of surface functional groupsmay be characterized by four discrete dissociation constants (pKa) based on Boehms

titration. Surface oxidation generates a distribution of surface functional groups.The positions of these different surface groups in close proximity to one anothermay influence their acidity and hence impact on the mechanism by which hydratedmetal ions are sequestered from solution. For example, the acidity increases by morethan an order of magnitude from benzoic to m-phthalic and salicylic acids. Thesurface carboxyl groups of oxidized carbons exhibit even lower dissociationconstants. This may be related to the fact that the surface groups are connectedto ap-conjugated condensed system of graphite-like planes. The number of conju-gated benzene rings and the positioning of the groups will also influence their

acidity. For example, carboxylic acids derived from naphthalene and anthracene

Table 15 Concentration and Distribution of Surface Functional Groups in Carbons

Carbon

Carboxylicgroups

(meq=g)

Lactones

(meq=g)

Phenolicgroups

(meq=g)

Total(non-carbonyl)

(meq=g)

Carbonylgroups

(meq=g)

Totalcapacity

(meq=g)

F400 0.047 0.073 0.003 0.123 0.235 0.358Percenta 38.21 59.35 2.44 100F400(ox) 0.719 0.439 0.427 1.586 1.356 2.941Percenta 45.33 27.68 26.92 100KAUini 0.078 0 0.370 0.448 0.628 1.076Percenta 17.41 0 82.59 100KAU-1.8 0.540 0.560 0.680 1.779 0.664 2.443Percenta 30.35 31.48 38.22 100KAU-2.9 1.430 0.650 0.795 2.875 2.531 5.405

Percenta 49.74 22.61 27.65 100KAU-1-25 0.580 0.320 0.380 1.280 1.170 2.450Percenta 45.30 25.00 29.70 100KAU-5-25 1.580 0.665 0.830 3.075 1.390 4.465Percenta 51.40 21.60 27.00 100CKC 1.149 0.585 0.462 2.160 1.056 3.216Percenta 53.19 27.08 21.39 100CKC-3-25 1.366 0.479 0.540 2.385 1.615 4.000Percenta 57.27 20.09 22.64 100

a

Percent in comparison to the total noncarbonyl capacity.


7/25/2019 DK1300_ch01

37/84

possess different dissociation constants depending on the location of the carboxylicfunctional group on the aromatic ring [44].

The results of metal sorption studies with F400-, KAU-, and CKC-basedadsorbents revealed that the uptake and selectivity toward Cu2 ions were greaterthan for Ni2, Co2, Zn2, and Mn2 (refer toFig. 18). Variation in the uptakeof the other metal ions was also detected; e.g., Co2 and Zn2 were less preferred

than Ni

2

, and Mn

2

was the least favored ion. Figure 18 clearly indicates that thecomplex stability=selectivity trend remains independent of the method and extentof adsorbent oxidation. It is also independent of the type of carbon precursor,the porous structure, and the type of adsorbent for all adsorptive materials investi-gated. Correlations of metal uptake as a function of individual groups, e.g.,carboxylic or lactonic, or sum total of noncarbonyl groups did not show a lineartrend. Based on the metal sorption data discussed above, the relative sorption affi-nity of metal ions can be described as follows (the arrangement of metals in thesequence is the same as in the periodic table): Mn2

7/25/2019 DK1300_ch01

38/84

configuration of the metal ion, is the other factor responsible for the variablecomplex stability [46,47]. As a rule, the greater LFSE is associated with the morestable complex. Consideration in terms of the ionic radius or the LFSE shows thatboth factors predict that the maximum stabilities should be associated with com-plexes of Ni2 rather than those of Cu2. This anomaly is a consequence of thestabilizing influence of the JahnTeller distortion [47] that results in stronger bind-ing of the four ligands in the plane of the tetragonally distorted Cu2 complex. Thetypical pattern of JahnTeller distortions, observed in Cu2 complexes, involves theformation of four shorter bonds and two trans bonds that are considerably longer

than the remaining four (tetragonal distortion). The reason for the JahnTeller dis-tortion is that the ninth electron in copper is placed into a set of the egorbitals insuch a way as to produce an asymmetric electron population (i.e., two in oneorbital and one in the other). This distortion is possible for any electronic con-figuration with asymmetry of this kind. The electron population of the egorbitalsis symmetric in Ni2 (i.e., one electron in each orbital), and therefore this ion doesnot exhibit any distortion.

Many Cu2 complexes are known to have either four short bonds and twolong bonds or two short and four long bonds. The outcome is that the JahnTeller

distortion of Cu2

compounds yields shorter and stronger metalligand bonds

Figure 18 Stability constants of Me2-adsorbent surface complexes.


7/25/2019 DK1300_ch01

39/84

(stronger complexes) than might be expected on the basis of the isotropic ionicradius of Cu2.The correlation between low molecular weight metal complexes with

carboxylic acids, hydroxy acids, and other compounds (as described in the IrvingWilliams series) with Mez OC bonds can only be remote to the surface com-plexes between Mez and adsorbent. However, there is good correlation betweenthe stability of such complexes [48] and the metal sorption by adsorptive materialsstudied (refer toFigs. 18and 19).

The formation of metal surface complexes on the oxidized carbon due tocooperative action seems quite likely, approximate calculation of oxygenated func-

tional group density per unit area for F400(ox) yields a value of 0.02 functionalgroup per angstrom squared. Given that the oxidized carbons evaluated in thepresent study possess a large proportion of pores in the region 1020 A and thediameter of the hydrated metal ions is approximately 8 A [49], it is reasonable toassume a cooperative binding mechanism (seeFig. 20).

In summary, the results show that these adsorbents exhibit an ability toremove metal ions from aqueous solutions with varying affinity in the orderMn2

7/25/2019 DK1300_ch01

40/84

of adsorbent precursor, porous structure, and type of adsorbent for all the materialsinvestigated. The higher preference of adsorbents for Cu2 is a consequence of thefact that this ion often forms distorted and hence more stable octahedral complexesdue to the asymmetric electronic structure. The generalization of metal sorptivebehavior using the IrvingWilliams approach leads to a novel way of understanding

metal sorption by active carbons and other carbonaceous adsorbents.We performed minicolumn experiments to determine breakthrough charac-teristics and regeneration performance. Adsorbents were used in as-received andoxidized form for the removal of copper, nickel, zinc, and cadmium from aqueoussolution. A brief summary of materials (including surface oxidation conditions)used for obtaining breakthrough column data is provided in Table 8. Physicaland chemical characterization of all samples has been described elsewhere[50,51]. Activated carbon samples in hydrogen form were packed into a mini-column (Isolute SPE columns of nominal capacity 6 mL, supplied by Jones Chroma-tography Ltd, UK) fitted with 20 mm polyethylene frits as bed supports. A known

quantity of carbon was contacted with distilled water before being placed in aminicolumn. A 1 mM solution of chloride salt of each metal was passed throughthe column to generate the breakthrough curves. All experiments were conductedusing a solution at a pH of 4.7. A flow rate of 510 bed volumes per hour (BV=h)was maintained during these experiments. A series of experiments were conductedusing a feed containing 1 mM of each of the metals under investigation.

Samples of BGP (a coal-derived carbon) were tested using solutions contain-ing target metals to determine overall performance criteria. This consisted of testingsamples of BGP against solutions containing copper, nickel, zinc, or cadmium. All

breakthrough capacities were calculated based on the total amount of metal

Figure 20 Postulated complexation reaction between copper(II) and oxidized carbonsurface.


7/25/2019 DK1300_ch01

41/84

removed before 5%or 50%of the feed concentration of metals was detected in theoutlet. This was to give an indication of uptake performance of each material for thefour metals used. The results of the minicolumn experiments are summarized inTable 17.

An unoxidized sample of BGP was used to generate a breakthrough curve tocompare copper sorption capacity with that of oxidized samples. The copper break-

through curves are presented in Fig. 21. Results show that there was significantimprovement in copper sorption capacity on oxidation of the original sample.The copper sorption capacities were increased by a factor of up to 100 by oxidizingthe original material. This change can be attributed to oxygen surface groups intro-duced onto the surface of Ceca BGP. The 5% copper breakthrough capacities ofBGP OxII and BGP AOx were 0.35 and 0.04 mmol=g, respectively. Over100 BV of 1 mM solution was passed before any copper could be detected leavingthe column using the acid-oxidized sample, BGP OxII.

Figure 22 shows breakthrough curves generated by using three different1 mM metal solutions (Ni2, Zn2, and Cd2). At a feed pH of 4.7, BGP OxII

had a 5% breakthrough capacity of 0.16 mmol=g for cadmium, 0.14 mmol=g fornickel, and 0.13 mmol=g for zinc. The breakthrough performance of the BGPOxII was very similar for all three metals, with 40 BV passing through the columnbefore any metals could be detected in the solution leaving the column.

BGP AOx was also used to generate breakthrough curves for removal ofnickel, zinc, and cadmium. Results for these experiments can be seen in Fig. 23.It shows that the metal sorption capacity of air-oxidized samples was much less thanthat of acid-oxidized samples for a feed concentration of 1 mM and pH of 4.7.Breakthrough capacities at 5%were 0.02 mmol=g for nickel, and for other samples

the 5%breakthrough was instantaneous. The low capacity of the BGP AOx sample

Table 17 Uptake Capacity of Carbons from 1 mM Metal Solutions (pH 4.7)

Breakthrough capacity (mmol=g)

Samplea Metal pH 5% 50%

BGP OxII Cu 4.7 0.35 0.42BGP AOx Cu 4.7 0.04 0.05BGP unoxidized Cu 4.7 0.004BGP OxII Zn 4.7 0.13 0.16BGP AOx Zn 4.7 0.13BGP OxII Ni 4.7 0.14 0.20BGP AOx Ni 4.7 0.03 0.06BGP OxII Cd 4.7 0.16 0.21BGP AOx Cd 4.7 0.02

aSamples: AOx, air-oxidized; OxII, alkali-washed, acid-oxidized for 9 h.


7/25/2019 DK1300_ch01

42/84

Figure 21 Copper breakthrough curves for samples of BGP. Feed concentration 1 mM copper, feed pH 4.7, flow rate 5 BV=h.

Figure 22 Breakthrough curves for nickel, zinc, and cadmium using BGP OxII sample.

Feed concentration 1 mM Ni, Zn, or Cd; feed pH 4.7; flow rate 5 BV=h.


7/25/2019 DK1300_ch01

43/84

can be partially attributed to the high reaction temperature during the oxidation,which rose above 723 K. This would have resulted in desorption of some oxygensurface groups. Very poor breakthrough performance was obtained for the BGP

AOx sample with the exception of the nickel experiment, when 10 BV was treatedbefore any metal was detected in the outlet.

A combined feed containing all four metalscopper, zinc, nickel, and cad-miumwas passed through the column containing 1 g of BGP OxII (seeFig. 24).The feed concentration of all metals was 0.25 mM, and the feed pH was 4.7.These indicate the selectivity series Cu2>Ni2,Cd2>Zn2. The increase in the

concentration above the feed was due to chromatographic elution. A solution con-taining copper, zinc, nickel, and cadmium was used to obtain the selectivity ofBGP AOx. The results shown in Fig. 25 indicate a selectivity order of Cu2>Ni2>Zn2,Cd2. These studies confirm that both materials are highly selectivetoward copper in the presence of other metal ions.

The results of the minicolumn experiments for various samples of F400 aresummarized in Table 18. The results for uptake of copper, nickel, zinc, andcadmium can be seen inFigs. 2629, respectively. These results indicate that thecapacity of unoxidized material was significantly less than that of the modified

samples. Figure 26 shows that the breakthrough of copper occurs instantly for the

Figure 23 Breakthrough curves for nickel, zinc, and cadmium using BGP AOx sample.Feed concentration 1 mM Ni, Zn, or Cd; feed pH 4.7; flow rate 5 BV=h.


7/25/2019 DK1300_ch01

44/84

F400 unoxidized sample. Breakthrough profiles for F400 AOx and OX9h HT sam-ples were found to be similar. Any functional groups that were destroyed at 593 Kduring heat treatment would also have been removed during air oxidation at 693 K.

The results of breakthrough experiments for Ox9h-WW, Ox9h-AW, andOx24h samples of F400 are shown inFig. 26.There was no improvement in break-through capacity as a result of further oxidation between 9 h and 24 h. Any func-tional groups introduced after an additional 15 h oxidation are counterbalancedby breakage and losses due to attrition.

The results for nickel are different, with a wider distribution of bed volumespassed before breakthrough. The breakthrough for the Ox9h-AW sample occurredafter 90 BV, whereas the Ox9h-WW sample treated over 120 BV before any nickelcould be detected at the outlet of the column. This means that the humic-typesubstances generated as by-products are removed during alkali washing had a sub-stantial nickel capacity. This effect was not as pronounced for copper because thechemical behavior of copper makes it more amenable to removal by any numberof functional groups and also by surface adsorption. Zinc and copper breakthroughexperiments were similar, as shown in Fig. 28, although breakthrough volumes

for zinc were lower. The best result was obtained for the Ox9h-WW, with over

Figure 24 Breakthrough curve using a feed solution containing copper, nickel, zinc, andcadmium with a sample of BGP OxII. Feed concentration 0.25 mM each of Cu, Ni, Zn,and Cd; feed pH 4.7; flow rate 5 BV=h.


7/25/2019 DK1300_ch01

45/84

100 BV of feed solution treated before breakthrough was detected.Figure 29showsminicolumn results for cadmium. The results are similar to those for zinc andnickel, with similar breakthrough volumes treated before cadmium is detected forOx9h-WW, Ox9h-AW, and Ox24h samples.

The lowest capacity at 50% breakthrough was 0.01 mmol=g for cadmiumusing an unoxidized sample. The highest capacity at 50% breakthrough was0.69 mmol=g for copper using an F400-Ox9h-WW sample. All the other sampleshad an uptake capacity between these two extremes. A 60-fold increase in copper

uptake was obtained by acid oxidation compared to the unoxidized sample.Experiments were conducted to determine the ion-exchange performance aftermultiple sorptiondesorption cycles. Samples of F400-Ox9h-WW and Ox9h-

AW were used for four cycles and regenerated with 0.1 M HCl after each sorptionrun. The results show that there was no significant difference in breakthrough per-formance between as-received samples and the samples after four sorptiondesorptioncycles (seeFig. 30).

Selectivity experiments were conducted using a solution containing 1 mMeach of copper, nickel, zinc, and cadmium. Three materials were selected for

these experiments: Ox9h-WW, Ox9h-AW, and AOx. The results can be seen in

Figure 25 Breakthrough curve using a feed solution containing copper, nickel, zinc, andcadmium with a sample of BGP AOx. Feed concentration 0.25 mM each of Cu, Ni, Zn,and Cd; feed pH 4.7; flow rate 5 BV=h.


7/25/2019 DK1300_ch01

46/84

Figs. 3133, which indicate that the selectivity order is Cd2

7/25/2019 DK1300_ch01

47/84

Figure 26 Copper breakthrough experiments for samples of F400. Feed concentration1 mM, flow rate 10BV=h, feed pH 4.7. (From Ref. 51, with permission from The Insti-tution of Chemical Engineers.)

Figure 27 Nickel breakthrough experiments for F400 samples. Feed concentration1 mM, flow rate 10BV=h, feed pH 4.7. (From Ref. 51, with permission from The Insti-

tution of Chemical Engineers.)


7/25/2019 DK1300_ch01

48/84

Figure 28 Zinc breakthrough experiments using F400 samples. Feed concentration1 mM, flow rate 10BV=h, feed pH 4.7. (From Ref. 51, with permission from The Insti-tution of Chemical Engineers.)

Figure 29 Cadmium breakthrough experiments using samples of F400. Feed con-centration 1 mM, flow rate 10BV=h, feed pH 4

Documents

DK1300_ch01