5 Physical–Chemical Properties and Reactivity of Cyanide in Water and Soil David A. Dzombak, Rajat S. Ghosh, and Thomas C. Young CONTENTS 5.1 Free Cyanide 58 5.1.1 Cyanide Ion Bonding 58 5.1.2 HCN Formation and Dissociation 58 5.1.3 HCN Volatilization 60 5.1.4 Free Cyanide Adsorption to Soil and Sediment 61 5.1.5 Free Cyanide Oxidation 62 5.1.6 Free Cyanide Hydrolysis 64 5.2 Metal Cyanides: Aqueous Species 65 5.2.1 Weak Metal–Cyanide Complexes 65 5.2.1.1 Formation 65 5.2.1.2 Dissociation 67 5.2.1.3 Adsorption on Soil and Sediment 68 5.2.1.4 Oxidation 71 5.2.2 Strong Metal–Cyanide Complexes 73 5.2.2.1 Formation 73 5.2.2.2 Dissociation 75 5.2.2.3 Adsorption on Soil and Sediment 76 5.2.2.4 Oxidation–Reduction 78 5.3 Metal–Cyanides: Solid Phase Compounds 79 5.3.1 Simple Cyanide Solids 80 5.3.2 Alkali or Alkaline Earth Metal–Metal Cyanide Complex Solids 80 5.3.3 Other Metal–Metal Cyanide Complex Solids 80 5.4 Cyanate 82 5.5 Thiocyanate 84 5.6 Organocyanides 86 5.7 Summary and Conclusions 88 References 88 The reactivity, fate, and toxicity of cyanide in water and soil is highly dependent on the chemical exist. The simplest form of soluble cyanide is the negatively charged cyanide ion, CN − , which is composed of a carbon atom triple bonded to a nitrogen atom (–C≡N). The nature of this triple bond controls the reactivity of the cyanide anion, including complexation with other metal cations, 57 © 2006 by Taylor & Francis Group, LLC speciation of the cyanide. As outlined in Chapter 2, many different soluble and solid forms of cyanide 58 Cyanide in Water and Soil formation of molecular hydrogen cyanide (HCN), oxidation of cyanide to cyanate, and adsorption onto clays and other soil components. In environmental systems, wastewaters, and wastes, cyanide usually is found in free and com- plexed forms, asHCN and as metal–cyanidecomplexes. Because of a reactiveelectronicarrangement, cyanide anions can readily form metal–cyanide complexes with most metal cations. Most of these complexes exist as soluble species, but many, particularly iron-cyanide complexes, can react further with metal cations to form stable cyanide solids. The soluble and solid phase cyanide species that this chapter, the specific physical–chemical properties and reactivity characteristics of the differ- ent chemical forms of cyanide are presented. Included are examinations of the nature of bonding in and with the cyano group and free cyanide speciation; the properties and reactivities of soluble metal–cyanide complexes; the properties and reactivities of metal–cyanide complex solids; and the properties and reactivities of cyanate, thiocyanate, and organocyanide compounds. 5.1 FREE CYANIDE 5.1.1 C YANIDE ION BONDING Free cyanide consists of the cyanide anion, CN − , and molecular hydrogen cyanide, HCN, both existing as water soluble entities. The cyanide ion acts as a monodentate ligand with the carbon acting as the donor atom, and also as an ambidentate ligand acting as a donor at both ends of the ion [1]. Several structural factors govern the reactivity of free cyanide. The triple bonded structure of a cyanide anion is comprised of a sigma bond, two π bonds, and two empty bonding orbitals [2]. The “s” and the “p” orbitals are filled with maximum number of electrons, while the “d” and “f” orbitals are empty. This configuration allows for a number of bonding arrangements. Since halogens also have filled “s” and “p” orbitals, the behavior of the cyanide anion is similar to that of halogens [3]. The cyanide ion is considered a pseudo-halide in that it can form π-acceptor covalent bonds with transition metals [3]. It may also share electrons at the triple bond with the Group VI elements oxygen and sulfur, forming cyanate, CNO − , or thiocyanate, SCN − [3], or may act as a strong nucleophile in reactions with organic molecules, for example, nucleophilic addition reactions with aldehydes and ketones to form cyanohydrins [4]. The cyanide ion readily forms neutral compounds or anionic complexes with most major metal cations. The partially or wholly filled “d” orbitals of transition series metals can form covalent bonds with the empty anti-bonding orbitals of the cyanide ion. This involves acceptance of electron density into π orbitals of the carbon atom. The cyanide ion is a strong σ donor, which is responsible for the high stability of some of the metal–cyanide complexes [3]. 5.1.2 HCN FORMATION AND DISSOCIATION The cyanide anion protonates in water to form hydrocyanic acid, HCN, the most toxic form of a for HCN dissociation reaction is 9.24 at 25 ◦ C [5]. Thus, at pH greater than 9.24, cyanide anion dominates free cyanide speciation, while soluble HCN is the dominant species under acidic to neutral pH conditions (pH < 9.24). The free cyanide dissociation reaction is as follows: HCN = H + +CN − ,pK a = 9.24 at 25 ◦ C, I = 0 (5.1) − species as a function of pH for a simple aqueous solution at 25 ◦ C. The temperature dependence of the equilibrium constant governing the species © 2006 by Taylor & Francis Group, LLC occur most often in water and soil are outlined in Chapter 2 and examined in more detail here. In cyanide (see Chapters 13 and 14). The pK Figure 5.1 shows the distribution of HCN and CN Physical–Chemical Properties and Reactivity 59 567 89 1110 pH Ionization fraction ([HCN]/CN T , [CN-]/ CN T ) CN – HCN 1.0 0.8 0.6 0.4 0.2 0 FIGURE 5.1 Free cyanide species distribution as a function of pH at 25 ◦ C(pK a = 9.24 for HCN dissociation at T = 25 ◦ C, I = 0). distribution of free cyanide can be calculated via the van’t Hoff equation: ln(K 2 /K 1 ) = (H r,25C /R)[1/T 1 −1/T 2 ] (5.2) where H r,25C is the standard enthalpy change of reaction at 25 ◦ C(298 K), R is the molar gas constant (8.314 ×10 −3 kJ mol −1 K −1 ), T 1 is the reference temperature (298 K), and T 2 is the temperature of interest in K. The standard enthalpy change for the reaction given in Equation (5.1) is 146 kJ mol −1 , as calculated using the thermodynamic data compiled in Stumm and Morgan [6]. Substitution of this value, and assuming it is approximately constant for the temperature range 5 to 30 ◦ C, enables calculation of the temperature dependence of the acidity constant in Equation (5.1): K T = exp[1.756 ×10 4 K(3.356 × 10 −3 K −1 −T −1 ) − 21.28] (5.3) where K T is the equilibrium constant for HCN dissociation at the temperature T (K) of interest. Combining Equation (5.3) with the mass action equation for the reaction in Equation (5.1), and the mass balance equation for free cyanide (molar concentrations in [ ]), TOTCN =[HCN]+[CN − ] (5.4) yields the following expression for the species distribution fractions for HCN and CN − : α HCN =[HCN]/TOTCN ={H + }/[{H + }+exp[1.756 × 10 4 K(3.356 × 10 −3 K −1 −T −1 ) − 21.28]] (5.5) α CN =[CN − ]/TOTCN = 1 −α HCN (5.6) where {H + } is the hydrogen ion activity, 10 −pH . The species distribution fraction for HCN, α HCN , ◦ © 2006 by Taylor & Francis Group, LLC is presented in Figure 5.2 for temperatures between 5 and 30 C (278 and 303 K), and zero ionic 60 Cyanide in Water and Soil 67891011 pH [HCN]/CN T T = 30°C T = 25°C T =20°C T =15°C T =10°C T =5°C 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 FIGURE 5.2 Ionization fraction for HCN as a function of pH and temperature (I = 0). TABLE 5.1 Literature Values of Henry’s Law Constant (K H,HCN ) for HCN Temp. ( ◦ C) K H (atm L mol −1 ) K H (mg L −1 a /mg L −1 w ) Reference 25 0.122 0.005 Bodek et al. [7] Not given 0.073 0.003 Doudoroff [80] Not given 0.104 to 0.114 0.0043 to 0.0047 Smith and Mudder [2] 25 0.115 0.0047 Avedesian [81] strength (I). As is evident in Figure 5.2, temperature has a significant effect on free cyanide spe- cies distribution. As temperature decreases, dissociation of HCN decreases, extending the species dominance of HCN to higher pH values. 5.1.3 HCN VOLATILIZATION Hydrogen cyanide has a very low boiling point (25.7 ◦ C) and thus is volatile in water under environmental conditions. The equilibrium air–water partitioning of HCN can be described by Henry’s Law: P HCN = K H,HCN [HCN] (5.7) where P HCN is the partial pressure of HCN gas, atm, K H,HCN the Henry’s Law constant, atm L mol −1 , and [HCN] the equilibrium aqueous phase concentration of HCN, mol L −1 . Table 5.1 lists reported values of Henry’s Law constant for HCN. Henry’s Law constants with units relevant to Equation (5.1) are provided, along with dimensionless analogs corresponding to an equilibrium partitioning expression in which both the aqueous and gas phase concentrations are expressed in the same mass concentration units. Note that the Henry’s Law constant is a function of temperature. Thereare variousempirical relationships that expressHenry’s Lawconstantas a function © 2006 by Taylor & Francis Group, LLC Physical–Chemical Properties and Reactivity 61 of temperature. One such relationship, reported by Bodek et al. [7], is as follows: log K H,HCN =−1272.9/T + 6.238 (5.8) where K H,HCN is the Henry’s Law constant, mm Hg/M and T the Temperature, K. Equation (5.8) is reported to be valid for HCN concentrations ranging from 0.01 to 0.5 M and temperatures from 20 to 95 ◦ C. 5.1.4 FREE CYANIDE ADSORPTION TO SOIL AND SEDIMENT Free cyanide (CN − , HCN) adsorbs weakly on soils and sediment. The cyanide anion can be retained by soils with anion exchange capacity, but in the pH range 4 to 9 of interest for most soils, HCN is the dominant form of cyanide and CN − concentrations are very low. HCN adsorbs weakly or not all to inorganic soil components such as iron oxide [8], aluminum oxide, clay, and sand [9]. However, HCN has been shown to adsorb significantly to soils with appreciable organic carbon content. The magnitude of cyanide adsorption onto soils tested by Chatwin et al. [10] showed excellent correlation with organic carbon content. Higgins and Dzombak [9] further demonstrated the interaction of HCN with organic carbon in experiments with activated carbon and freshwater sediment. They developed an expression relating sorbed HCN concentration, C S , to aqueous phase concentration, C w , through an organic carbon normalized distribution coefficient K oc (=K d /f oc ). C S = K oc C w f oc = (6.5 L/g s )C w f oc (5.9) where C S is in µg/g s , C W is in µg/L, and f oc is the fraction of organic carbon in the adsorbent. The experiments upon which this linear relationship is based all involved low concentrations of free cyanide in water (<150 µg/L), which is typical for total cyanide concentrations encountered in environmental contamination scenarios. Adsorption capacities were not determined in the experi- ments with activated carbon and sediment. Literature data on free cyanide adsorption onto activated carbon have shown an adsorption capacity of about 1 to 2 mg of free cyanide per gram of carbon, while similar tests performed with soil organic carbon have revealed an adsorption capacity of 0.5 mg of free cyanide per gram of carbon [11]. Batch and column tests performed by Alesii and Fuller [12] with various soils yielded significant removal of free cyanide at near-neutral pH values. Soil constituents included kaolin clay, chlorite, gibbsite clay, and iron and aluminum oxides. Based on the laboratory results discussed earlier, it is unlikely that these inorganic constituents would adsorb free cyanide to an appreciable extent. As the soils used in the experiments by Alesii and Fuller were not sterilized and hence biologically active, it is more likely that the free cyanide was removed from the system via biodegradation. Dzombak and Morel [13] estimated equilibrium surface complexation constants for the adsorp- tion of CN − , CNO − , and SCN − on hydrous ferric oxide based on correlations of acidity constants and surface complexation constants fitted to adsorption data for other inorganic ions. The surface complexation reactions and the estimated surface complexation constants for those reactions are −4 M solutions of these ions in hydrous ferric oxide suspensions with TOTFe = 10 −3 M and ionic strength of 0.01 M are shown in face complexation constants, the predictions provide some idea of the expected adsorption behavior based on what has been observed with other inorganic ions. Available data for free cyanide adsorp- tion on mineral surfaces, however, indicates that the free cyanide adsorption in Figure 5.3 is likely to be substantially overpredicted. Free cyanide has been observed to exhibit little to no adsorption on mineral surfaces, including the crystalline iron oxide goethite, across a range of pH [8,9]. © 2006 by Taylor & Francis Group, LLC Figure 5.3. While the accuracy of these predictions is uncertain due to the estimated nature of the sur- given in Table 5.2. Predicted adsorption versus pH curves for 10 62 Cyanide in Water and Soil TABLE 5.2 Estimated Surface Complexation Reactions and Constants for Adsorption of CN − , CNO − , and SCN − on Hydrous Ferric Oxide Adsorbing log K 2 log K 3 species, A − (25 ◦ C, I = 0) a (25 ◦ C, I = 0) b CN − 13.0 5.7 CNO − 8.9 1.8 SCN − 7.0 0.1 a SC reaction: ≡FeOH 0 +A − +H + =≡FeA 0 +H 2 O; K 2 . b SC reaction: ≡FeOH 0 +A − =≡FeOHA − ; K 3 . Source: Data from Dzombak, D.A. and Morel, F.M.M., Surface Complexation Modeling: Hydrous Ferric Oxide, Wiley-Interscience, New York, NY, 1990 (Table 10.10). pH Percent adsorbed CN – CNO – – 0 10 20 30 40 50 60 70 80 45678910 FIGURE 5.3 Predicted adsorption of 10 −4 MCN − , CNO − , and SCN − on hydrous ferric oxide as a function of pH. Predictions made using surface complexation constants of Dzombak, D.A. and Morel, F.M., Surface ComplexationModeling: Hydrous Ferric Oxide, 1990; seeTable 5.2. TOTFe = 0.001M, I = 0.1M. Adsorption of CN − is likely overpredicted. 5.1.5 FREE CYANIDE OXIDATION Free cyanide can be oxidized to cyanate, CNO − , or hydrogen cyanate, HCNO, depending on the pH [14]: CN − +H 2 O = CNO − +2H + +2e − (5.10) HCN +H 2 O = HCNO +2H + +2e − (5.11) © 2006 by Taylor & Francis Group, LLC SCN Physical–Chemical Properties and Reactivity 63 Cyanate is protonated only at low pH, as the pK a is 3.45 [5]: HOCN = H + +CNO − ,pK a = 3.45 at 25 ◦ C, I = 0 (5.12) In the oxidation reactions of Equations (5.10) and (5.11), the oxidation state of carbon in CN − is +2, while in CNO − , the oxidation state of carbon is +4. Free cyanidecanbeoxidizedreadily by strong oxidants such as chlorine, hypochlorite, ozone, and hydrogen peroxide [15–18]. Under neutral to alkaline conditions, the end product iscyanate (CNO − ), which is relatively nontoxic. The oxidative conversion of CN − to CNO − in alkaline chlorination is often exploited for rapid treatment of free cyanide in water. The general reaction chemistry for alkaline chlorination is as follows [17]: CN − +Cl 2 → CNCl + Cl − (5.13) CNCl +2NaOH → CNO − +2Na + +2Cl − +H 2 O (5.14) 3Cl 2 +2CNO − +6NaOH → 2HCO − 3 +N 2 +6Cl − +6Na + +2H 2 O (5.15) As indicated in Equation (5.13), cyanide ion is first converted to cyanogen chloride, CNCl, a highly toxic species. Under alkaline conditions, the CNCl reacts rapidly with OH − to form CNO − , and upon further chlorination the cyanate decomposes to form the completely benign products bicarbonate, HCO − 3 and elemental nitrogen, N 2 . In the last step, Equation (5.15), the nitrogen is oxidized, moving from an oxidation state of −3 to zero. Gurol and Bremen [19] studied ozonation of free cyanide. It was found that the ozone molecule, O 3 , reacts primarily with the cyanide ion; its reaction with HCN is negligible. Further, it was determined that the presence of free cyanide promotes the formation of free radicals (OH • ,HO 2 • ), and that free radical reactions as well as direct reaction of the free cyanide with ozone contribute to the overall oxidative destruction of the cyanide. Hence, there are numerous initiators and pathways involved in the oxidation of free cyanide by ozone. Some of the reactions identified by Gurol and Bremen [19] as involved with the ozonation of free cyanide are as follows: O 3 +CN − → CNO − +O 2 (5.16) HCN +OH • → HCNO +HO 2 • (5.17) CN − +OH • → CNO − +HO 2 • (5.18) CN − +OH • → CN • +OH − (5.19) CN • +CN • → (CN) 2 (5.20) (CN) 2 +2OH − → CNO − +CN − +H 2 O (5.21) The direct reaction of molecular ozone with the cyanide ion is indicated in Equation (5.16). Other reactions of ozone with water, specifically OH − , yield the superoxide radical O 2 − • , which reacts further with ozone to give the hydroxyl radical OH • . The oxidation of cyanide by ozone is rapid and pH dependent [19]. Solutions of several mM of free cyanide were oxidized within 5 to 30 min by ozone, with faster rates at higher pH values where more of the free cyanide was in the form of CN − . The end product of ozonation of free cyanide is cyanate. The cyanate is further oxidized by ozone, but since this is a relatively slow reaction cyanate accumulates in solution until free cyanide is oxidized completely [20]. Free cyanide in the environment is oxidized rapidly by aerobic bacteria, for which it can serve the environment is usually of secondary concern. © 2006 by Taylor & Francis Group, LLC as an energy source, as discussed in Chapters 6 and 23. Thus, abiotic oxidation of free cyanide in 64 Cyanide in Water and Soil Chatwin et al. [10] detected cyanate in effluent from saturated soil columns through which aqueous solutions containing free cyanide were passed. It was hypothesized that the free cyanide was oxidized to cyanate on the surfaces of clay components of the soil, and that the process was enhanced with the addition of copper and nickel to the system. However, since the soils studied were microbiologically active, microbial degradation was likely to have had an equal or possibly greater role in the conversion of the cyanide, a factor not addressed by Chatwin et al. [10]. Based on other work showing very limited to no interaction of free cyanide with mineral surfaces [8], it appears unlikely that abiotic oxidation of free cyanide on mineral surfaces will occur appreciably in natural systems. Free cyanide can also react with and be oxidized by various forms of sulfur, especially poly- sulfides and thiosulfate (S 2 O 2− 3 ), to form thiocyanate (SCN − ). In neutral to alkaline solutions, both polysulfides and thiosulfate are products of oxidation of sulfide. The reactions of polysulfide and thiosulfate with the cyanide ion are as follows [2,21]: S x S 2− +CN − → S x−1 S 2− +SCN − (5.22) S 2 O 2− 3 +CN − → SO 2− 3 +SCN − (5.23) For thiocyanate, the oxidation states of the S, C, and N are −1, +3, and −3, respectively. In the reaction of polysulfide with free cyanide (Equation 5.22), one polysulfide sulfur atom is reduced from its oxidation state 0 to −1, while the cyanide carbon atom is oxidized from +2to+3 [21]. In the reaction of thiosulfate with free cyanide (Equation 5.23), one thiosulfate sulfur atom changes from oxidation state +2to+4, while the other thiosulfate sulfur atom is reduced from the +2to the −1 oxidation state [21]. The rate of thiocyanate formation through reaction of polysulfide and free cyanide is approximately three orders of magnitude greater than through reaction of free cyanide and thiosulfate, depending on pH [21]. Thus, in systems with equal amounts of polysulfide and thiosulfate present, the reaction of free cyanide with polysulfide will be the dominant thiocyanate formation route. The formation of polysulfide through oxidation of sulfide occurs at a slow rate, however, so available polysulfide is often limited [21]. Reaction of thiosulfate with free cyanide thus governs the formation of thiocyanate in many systems. 5.1.6 FREE CYANIDE HYDROLYSIS As discussed in Section 5.1.2, the cyanide ion reacts with water (H + ) to form HCN, with the protonated species HCN being the dominant form of free cyanide at pH values less than 9.24 at ◦ cyanide with water. Free cyanide can react with molecular water under alkaline conditions and high temperature to yield formate and ammonia: CN − +2H 2 O → HCOO − +NH 3 (5.24) The reaction proceeds at appreciableratesonly at high temperatures, and at fast rates at high temperat- ure and pressure, for example, temperatures in the rangeof 165–180 ◦ C and pressuresof 100–150 psig lysis is very slow at room temperature, increasing about threefold for every 10 ◦ C rise in temperature. At lower pH values, HCN can also be hydrolyzed, yielding formic acid and ammonia [2]: HCN +2H 2 O → HCOOH + NH 3 (5.25) Under acidic conditions the reaction is also very slow. © 2006 by Taylor & Francis Group, LLC [22]; seeChapters20and 22. Wiegand and Tremelling[23]showed that the rate of free cyanide hydro- 25 C (Figure 5.1). At ambient temperatures, this protonation reaction is the primary reaction of free Physical–Chemical Properties and Reactivity 65 Alkaline hydrolysis has been exploited for treatment of free and complexed cyanide in waste- waters and sludges [22,24]. Alkaline conditions assure that any free cyanide remains dissolved in the form of CN − during the treatment process. Temperatures in the range of 140 to 275 ◦ C, and pressures up to 900 psig are employed in the alkaline hydrolysis process. This treatment process is discussed 5.2 METAL CYANIDES: AQUEOUS SPECIES Many metals form aqueous complexes with cyanide ion by means of π bonding, which occurs when the participating metal atom donates one or more electrons to CN − , which serves as an elec- tron accepting ligand. These soluble metal–cyanide complexes, represented by a general formula [M(CN) x ] y− , where, M signifies a metal cation, can be classified into weak and strong metal–cyanide complexes, depending upon the strength of the metal–cyanide bonding. Use of vibrational spec- troscopy reveals different electronic structures of [M(CN) x ] y− complexes [1]. Depending on the different modes of vibration, a [M(CN) x ] y− species can exist in tetrahedral, square planar, or octahedral forms. These common electronic structures are shown in Figure 5.4. 5.2.1 WEAK METAL–CYANIDE COMPLEXES 5.2.1.1 Formation The cyanide anion can form weak metal–cyanide complexes with many transition metals, the most common among them being cadmium, zinc, silver, copper, nickel, and mercury. Most of these metals fall in Groups IB, IIB, and VIIIB of the periodic table. The metal–cyanide bonds in these complexes are mostly arranged in tetrahedralorsquareplanarformswithrelativelyweakbondingenergy existing between the heavy metal atom and the cyanide ligand as compared to the strong cyanide complexes with iron, cobalt, and platinum. Because weakly-bonded metal–cyanide complexes dissociate under weakly acidic pH conditions (4 < pH < 6), they are commonly termed weak-acid-dissociable (WAD) complexes [15]. Formation data determined by direct thermodynamic measurements are available for complexes of nickel(II),copper(I),silver(I), zinc(II), cadmium(II),andmercury(II)[1,5]. Equilibrium orstability constants have been calculatedfrom standard thermodynamic data for a broadrange of metal–cyanide Sehmel [5] for formation of weak metal–cyanide complexes. For comparable reaction stoichiometry, the higher the value of the formation equilibrium constant (K), the greater is the energy of formation and the stability of the metal–cyanide complex. [Fe(CN) 6 ] 3– (ferricyanide) Ni 2+ C C C C N N N N Ni 2+ C C C C N N N N [Ni(CN) 4 ] 2– (tetracyanonickelate) Fe (III) N N C C N N C Fe (III) C N N C C FIGURE 5.4 Common electronic structures of metal–cyanide complexes. © 2006 by Taylor & Francis Group, LLC in more detail in Chapter 22. complexes, however [5]. Table 5.3 lists the measured and calculated stability constants compiled by 66 Cyanide in Water and Soil TABLE 5.3 Equilibrium Constants for Formation of Selected Weak Metal–Cyanide Complexes Reaction log K (at25 ◦ C, I = 0) Ag + +CN − +H 2 O = AgCN(OH) − −0.56 Ag + +2CN − = Ag(CN) − 2 20.38 Ag + +3CN − = Ag(CN) 2− 3 21.40 Ag + +2OCN − = Ag(OCN) − 2 5.00 Cd 2+ +CN − = CdCN − 5.32 Cd 2+ +2CN − = Cd(CN) 0 2 10.37 Cd 2+ +3CN − = Cd(CN) − 3 14.83 Cd 2+ +4CN − = Cd(CN) 2− 4 18.29 Cu + +2CN − = Cu(CN) − 2 24.03 Cu + +3CN − = Cu(CN) 2− 3 28.65 Cu + +4CN − = Cu(CN) 3− 4 30.35 Ni 2+ +2CN − = Ni(CN) 0 2 14.59 Ni 2+ +3CN − = Ni(CN) − 3 22.63 Ni 2+ +4CN − = Ni(CN) 2− 4 30.13 Ni 2+ +H + +4CN − = NiH(CN) − 4 36.75 Ni 2+ +2H + +4CN − = NiH 2 (CN) 0 4 41.46 Ni 2+ +3H + +4CN − = NiH 3 (CN) + 4 43.95 Zn 2+ +2CN − = Zn(CN) 0 2 11.07 Zn 2+ +3CN − = Zn(CN) − 3 16.05 Zn 2+ +4CN − = Zn(CN) 2− 4 16.72 Hg(OH) 0 2 +2H + +CN − = HgCN + +2H 2 O 24.17 Hg(OH) 0 2 +2H + +2CN − = Hg(CN) 0 2 +2H 2 O 40.65 Hg(OH) 0 2 +2H + +3CN − = Hg(CN) − 3 +2H 2 O 44.40 Hg(OH) 0 2 +2H + +4CN − = Hg(CN) 2− 4 +2H 2 O 47.41 Hg(OH) 0 2 +2H + +2CN − +Cl − = Hg(CN) 2 Cl − +2H 2 O 40.37 Hg(OH) 0 2 +2H + +3CN − +Cl − = Hg(CN) 3 Cl 2− +2H 2 O 43.83 Hg(OH) 0 2 +2H + +3CN − +Br − = Hg(CN) 3 Br 2− +2H 2 O 44.94 Source: Data from Sehmel, G.A., Cyanide and antimony thermodynamic database for the aqueous species and solids for the EPA-MINTEQ geochemical code, PNL-6835, Pacific Northwest Laboratory, Richland, WA, 1989, (Table 5). The equilibrium constants compiled by Sehmel [5] were selected and included in Table 5.3 rather than those reported in some other compilations, for example, the work of Beck [25] and Martell et al. (1993), because the constants reported by Sehmel were calculated in a consistent manner using the most current thermodynamic data from the U.S. National Bureau of Standards [26,27]. The constants reported in Beck [25] were calculated using older (1952) NBS thermodynamic data [28]. As shown by Gilgore-Schnorr and Dzombak [29], the key difference is in the value used for the partial molar entropy, S o , of the cyanide ion CN − , for which the 1952 [28] value of 28.2 cal K −1 mol −1 was revised in 1965 [27] to 22.5 cal K −1 mol −1 , a value retained in the 1982 (and most current) thermodynamic data compilation [26]. The work of Sehmel [5] was performed for the USEPA, which incorporated the metal–cyanide complexation constants in the thermodynamic database of the general chemical © 2006 by Taylor & Francis Group, LLC [...]... 6CN− = NH5 Fe(CN)2− 4 6 Na+ + Fe2+ + 6CN− = NaFe(CN)3− 6 2Na+ + Fe2+ + 6CN− = Na2 Fe(CN)2− 6 Na+ + Fe2+ + H+ + 6CN− = NaHFe(CN)2− 6 Sr 2+ + Fe3+ + 6CN− = SrFe(CN)− 6 Tl+ + Fe2+ + 6CN− = TlFe(CN)3− 6 log K (at 25 C, I = 0) 49.40 55 .44 49.69 55 .47 51 .00 52 .71 45. 61 50 .00 52 . 45 52.63 56 .98 52 .31 50 .22 48.12 48.98 51 .47 47.69 48 .53 51 .22 55 .39 49.43 48.07 48.87 51 .40 47.99 48.74 51 .43 55 .62 48. 75 Source:... by oxidizing agents such as chlorine or ozone The more strongly bonded complexes in the WAD category, such as nickel, silver, and mercury cyanide complexes, oxidize more slowly [ 15] The more weakly-bonded complexes, including those of cadmium, copper, and zinc, decompose rapidly in the presence of oxidizing agents As discussed in Section 5. 1 .5, and in more detail in Chapter 20, alkaline chlorination... ferro- and ferricyanide, Broderius and Smith [ 45] estimated mid-day half-lives (for mid-summer at the latitude and climatic conditions of St Paul, MN) for 25 to 100 µg/L concentrations of these species to be 18 and 64 min, respectively Photolytic degradation of ferro- and ferricyanide follows approximately first-order kinetics [ 45, 49], at least initially, but the rate slows as free cyanide accumulates in. .. mass law expressions in Table 5. 10 may be combined with appropriate mass law expressions from Tables 5. 3 and 5. 4 to determine solubility products corresponding to reactions written with metal cyanide species as dissolution products 5. 4 CYANATE As discussed in Section 5. 1 .5 and shown in Equations (5. 10) and (5. 11), free cyanide can be oxidized to form cyanate, CNO− , or, depending on the pH, its protonated... in turn regulates the cyanide speciation Figure 5. 11 presents a species predominance diagram for dissolved cyanide species in a system in equilibrium with hydrous ferric oxide The diagram was calculated with MINEQL+ [32] using the reactions and equilibrium constants in Equation (5. 1) and Table 5. 4, and in the MINEQL+ thermodynamic database for the iron dissolution, hydrolysis, and redox reactions In. .. F.M.M., Surface Complexation Modeling: Hydrous Ferric Oxide, Wiley-Interscience, New York, NY, 1990 14 Bard, A.J., Parsons, R., and Jordan, J., Standard Potentials in Aqueous Solutions, Marcel Dekker, Inc., New York, 19 85 15 APHA/AWWA/WEF, Method 450 0-CN: Cyanide, in Standard Methods for the Examination of Water and Wastewater, 20th ed., Clesceri, L.S., Greenberg, A.E., and Eaton, A.D., eds., American... of cyanide in aluminum potlining leachate using ozone as an oxidizing agent, M.S thesis, Clarkson University, Potsdam, NY, 1996 19 Gurol, M.D and Bremen, W.M., Kinetics and mechanism of ozonation of free cyanide species in water, Environ Sci Technol., 19, 804, 19 85 20 Gurol, M.D and Holden, T.E., The effect of copper and iron complexation on removal of cyanide by ozone, Ind Eng Chem Res., 27, 1 157 ,... relation to water quality criteria, Report 98-HHE -5 , Water Environment Research Foundation, Alexandria, VA, 2003 51 Laszlo, Z and Dombi, A., Oxidation of [Fe(CN)6 ]4− and reduction of [Fe(CN)6 ]3− in VUV-irradiated aqueous solutions, Chemosphere, 46, 491, 2002 52 Paschka, M.G., Ghosh, R.S., and Dzombak, D.A., Potential water- quality effects from iron cyanide anticaking agents in road salt, Water Environ... dissociation in natural waters is dependent on various environmental factors, including free cyanide content of the solution, sunlight intensity, temperature, turbidity, and depth of the water column [ 45, 50] In many surface waters, significant photolysis will occur only in the top 50 to 100 cm of the water column where sunlight intensity is sufficient, providing opportunity for dilution of any free cyanide. .. importance in water and soil systems Where cyanide has been introduced in water and soil systems, these © 2006 by Taylor & Francis Group, LLC Physical–Chemical Properties and Reactivity 81 TABLE 5. 5 Some Simple Cyanide Solids and Solubility Information Aqueous solubilityd (g/100 g H2 O) Temp (◦ C) 1.7a Qualitative solubilitya,c 15 Formula Physical forma Barium cyanide Cadmium cyanide Calcium cyanide Cyanogen . discussed in Section 5. 1 .5, and in more detail in Chapter 20, alkaline chlorination is the 72 Cyanide in Water and Soil Cd +2 = 2.08 ϫ 10 – 5 M Cu +2 = 1.92 ϫ 10 – 5 M Zn +2 =2.41ϫ 10 – 5 M α. 5 Physical–Chemical Properties and Reactivity of Cyanide in Water and Soil David A. Dzombak, Rajat S. Ghosh, and Thomas C. Young CONTENTS 5. 1 Free Cyanide 58 5. 1.1 Cyanide Ion Bonding 58 5. 1.2. governing the species © 2006 by Taylor & Francis Group, LLC occur most often in water and soil are outlined in Chapter 2 and examined in more detail here. In cyanide (see Chapters 13 and 14).