CYANIDE in WATER and SOIL: Chemistry, Risk, and Management - Chapter 5 pps

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
Physical–Chemical Properties and Reactivity 67
0
60
40
20
100
80
34 56 78 9
pH
Percent Zn species
CN
T
= 2.00 ϫ 10
– 3
M
I = 0.1 M
Zn(CN)
4
+2
Zn(CN)
3
–
Zn(CN)
1
+
P
CO
2

=10
– 3.5
atm
Zn
T
= 2.00 ϫ 10
– 5
M
Zn
+2
Zn(CN)
2
0
FIGURE 5.5 Calculated aqueous speciation of zinc(II) in the presence of excess cyanide. Zn
T
= 10
−4.70
M,
CN
T
= 10
−2.70
M,P
CO
2
= 10
−3.5
atm, I = 0.1 M.(Source: Theis, T.L. and West, M.L., Environ. Technol.
Lett., 7, 309, 1986.)
equilibrium program MINTEQ[30,31]. Several other general chemical equilibrium programsemploy

the MINTEQ thermodynamic database, such as MINEQL+ [32].
An example of zinc speciation in aqueoussolution calculated by Theis and West [8] is presented in
Figure 5.5. (Note: Equilibrium constants from Sehmel [5] were not used in calculating the speciation
plots ofFigure5.5; the plots are forillustrativepurposes.) In the system shown, thetotalamountof free
cyanide added to the system (2 ×10
−3
M) is in excess of the amount of zinc present (2 ×10
−5
M).
Thus, cyanide species dominate the speciation of zinc above pH 6.2. At lower pH values, H
+
outcompetes Zn
2+
for complexation with CN
−
as H
+
becomes more abundant and HCN forms.
Soluble zinc hydroxide complexes also form, at higher pH, but in this system CN
−
outcompetes
OH
−
in the pH range shown. Concentrations of dissolved zinc hydroxide species are very small and
their influence on zinc speciation is not significant. As is evident in Figure 5.5, different zinc cyanide
species predominate in different pH regions, with Zn(CN)
0
2
the dominant form from pH 6.2 to 6.7,
and Zn(CN)

2−
4
being most abundant above pH 6.7.
The kineticsofmetal–cyanide complexformationcanbe slow[1,33]. For example, Broderius [33]
showed that Ni(CN)
2−
4
formation took three days to reach equilibrium (0.5 to 6.5 ppm CN
T
), while
formation of copper cyanide complexes took over 100 days to reach equilibrium. The slow rate of
formation of metal–cyanide complexes, and the potential for oxidation of many weak metal–cyanide
complexes by atmospheric oxygen, makes it difficult to measure the equilibrium constants for these
complexes [1].
5.2.1.2 Dissociation
The dissociation properties of weak metal–cyanide complexes in aqueous solutions depend on their
stability constants, pH, temperature and the redox potential of the solution. In general, metal–cyanide
complex dissociation may be described by
M(CN)
n−
x
= M
+
+xCN
−
(5.26)
© 2006 by Taylor & Francis Group, LLC
68 Cyanide in Water and Soil
where M
+

is a monovalent metal cation, x is the number of cyanide groups, and n is the ionic charge
of the metal–cyanide complex.
metal–cyanide complexes in solution to dissociate. The higher the value for K for reactions with the
same stoichiometry, the more limited the dissociation into free cyanide (CN
−
).
Because of the labile nature of weak metal–cyanide complexes, mildly acidic conditions (pH ≈ 4
to 6) can result in the dissociation of many of these complexes, especially nickel and zinc. For this
reason, the weak-acid-dissociable cyanide analytical method [15] employs a pH 4.5 buffer while
distilling the aqueous sample at 125
◦
C for 1.5 h to break down these complexes completely and
5.2.1.3 Adsorption on Soil and Sediment
Weak metal–cyanide complexes can adsorb on common soil and sediment components such as
iron, aluminum, silicon, and manganese oxides, and clays, which in most systems will inhibit their
aqueous transport [8,12,34]. However, complexation of metals by cyanide can also serve to hold them
in solution, inhibiting their adsorption and retention. The enhancement or inhibition of adsorption
depends on the metal–cyanide species, the adsorbent, and the solution conditions.
Theis and Richter [34] studied the adsorption ofthe predominant nickel–cyanide anion Ni(CN)
2−
4
on silicon dioxide, SiO
2
(s) and the iron oxide goethite, FeOOH(s). Batch adsorption experiments
were conducted in 0.01 M NaClO
4
aqueous solutions containing 10
−4.77
M total nickel (Ni
T

) and
amounts of total free cyanide (CN
T
)of10
−5
,10
−4
, and 10
−3
M. Calculated plots of the equilibrium
distribution of nickel species as a function of pH in aqueous solution with no solids present are
the speciation plots of Figure 5.6; the plots are for illustrative purposes.) As may be seen there,
Ni(CN)
2−
4
is predicted to dominate nickel speciation at pH < 5.5 in the system with CN
T
=
10
−4
M, and at pH < 4.5 in the system with CN
T
= 10
−3
M. A speciation diagram for the
system with CN
T
= 10
−5
M is not shown, but with Ni

T
in excess of CN
T
in this system Ni
2+
and NiOH
+
are the dominant forms of nickel and a maximum of about 10% of the nickel becomes
bound to cyanide. Results of the batch adsorption experiments conducted with SiO
2
(s) are shown in
T
= 10
−5
M is very similar to
the adsorption of nickel on SiO
2
(s) in the absence of cyanide. For the series of batch experiments
with CN
T
= 10
−4
and 10
−3
M, however, adsorption of nickel is inhibited. These data indicate
that the nickel–cyanide species Ni(CN)
2−
4
, which dominates nickel speciation in both systems,
has no affinity for the SiO

2
(s) surface. The SiO
2
(s) surface is negatively charged for pH > 2, so
electrostatic repulsion of the negative Ni(CN)
2−
4
species is in part responsible for the absence of
surface binding. For the most part, however, it is the presence of cyanide that inhibits adsorption
FeOOH(s). Once again, the adsorption of nickel observed in the system with CN
T
= 10
−5
M is very
similar to the adsorption of nickel on FeOOH(s) inthe absence of cyanide. In the systems withCN
T
=
10
−4
and 10
−3
M, adsorption of nickel is enhanced at lower pH values, and inhibited at higher pH
values. The FeOOH(s) surface is positively charged up to about pH 6, or even higher, so electrostatic
attraction of Ni(CN)
2−
4
explains in part its adsorption at lower pH values. Electrostatic attraction
alone is notsufficient to explain the extent of removal observed, however. Throughsurface interaction
modeling, Theis and Richter [34] demonstrated that Ni(CN)
2−

4
must bond at specific surface sites
on goethite, in surface complexation reactions that involve high free energies of interaction. They
proposed the formation of a goethite–cyanide–metal surface complex via a surface complexation
reaction:
≡Fe
2
(OH)
2+
2
+Ni(CN)
2−
4
+2H
+
=≡Fe
2
–(CN)
2
–Ni–(CN)
2
+2H
2
O (5.27)
© 2006 by Taylor & Francis Group, LLC
quantify the amount of cyanide associated with them (see Chapter 7).
given in Figure 5.6. (Note: Equilibrium constants from Sehmel [5] were not used in calculating
Figure 5.7. The adsorption of nickel observed in the system with CN
of nickel. Figure 5.8 shows data for similar sets of batch adsorption experiments performed with
The stability constants given in Table 5.3 provide an indication of the propensity of weak acid

Physical–Chemical Properties and Reactivity 69
345678910
60
40
20
80
3456789 10
60
40
20
80
pH
Ni
2+
Ni
2+
(a)
(b)
pH
Ni(CN)
4
2–
Ni(CN)
4
2–
CN
T
=10
– 4
M

CN
T
=10
– 3
M
0
100
Percent total Ni
0
100
Percent to tal Ni
FIGURE 5.6 Theoretical distribution of nickel in the presence of (a) 10
−4
M cyanide (Ni
T
= 10
−4.77
M, I =
0.01 M), and (b) 10
−3
M cyanide (Ni
T
= 10
−4.77
M,I= 0.01 M). (Source: Reprinted with permission from
Theis, T.L. and Richer, R.O., Particulates in water, 189, 73, 1980. Copyright 1980 American Chemical Society.)
where ≡Fe
2
(OH)
2+

2
is a surface hydroxyl site on the surface of goethite in aqueous suspension, and
≡Fe
2
–(CN)
2
–Ni–(CN)
2
is the surface species formed by adsorption of Ni(CN)
2−
4
on the goethite.
The uptake of H
+
shown in the reaction occurs commonly in adsorption of inorganic anions on
oxides, and is related to the commonly observed pH dependence for anion adsorption: maximum
adsorption at lower pH and decrease in adsorption with increasing pH [13]. Formation of a metal–
ligand–metal ternary surface complex as shown in the reaction of Equation (5.27) has been proposed
for other metal–ligand systems [35–38].
Theis and West [8] studied the adsorption of cadmium, copper, and zinc divalent cations and their
metal–cyanide complexes on goethite in aqueous suspensions. Some typical results for adsorption of
Cd
2+
,Cu
2+
, and Zn
2+
were conducted with total metal concentration of approximately 2 × 10
−5
M and with 0.6 g/L

© 2006 by Taylor & Francis Group, LLC
in the absence of cyanide are presented in Figure 5.9. All of the experiments
70 Cyanide in Water and Soil
345678910116
0
30
20
10
50
40
80
70
60
100
90
pH
Percent nickel removed
FIGURE 5.7 Nickel adsorption as a function of pH in the presence of silicon dioxide and Cyanide. Ni
T
=
10
−4.77
M, I = 0.01 M, SiO
2
= 29.41 g/L. ()CN
T
= 10
−5
M,()CN
T

= 10
−4
M,()CN
T
= 10
−3
M,
(
•)CN
T
= 0. (Source: Reprinted with permission from Theis, T.L. and Richer, R.O., Particulates in water,
189, 73, 1980. Copyright 1980 American Chemical Society.)
cation adsorption on metal oxides: an increase from 0 to 100% adsorbed with increasing pH. Batch
adsorption experiments conducted with free cyanide showed no adsorption of the free cyanide on
goethite for any pH from 3 to 11 (data not shown). Experiments with free cyanide added in excess of
At lower pH values, adsorption of the cadmium, copper, and zinc was unaffected by the free cyanide
as may be seen by comparisonwithFigure 5.9. Above pH 6.5 to 7.0, however, adsorption of the metals
was inhibited by the presence of the cyanide. At the higher pH values, metal–cyanide complexes
dominate the speciation of the metals (e.g., see the aqueous phase speciation diagram for zinc in
have no affinity for the goethite surface at neutral to alkaline pH values.
The examples presented in this section demonstrate that some weak metal–cyanide complexes
can adsorb on soils and soil components under some conditions, but the extent of adsorption depends
strongly on the particular metal–cyanide species, mineral adsorbent, and solution conditions. Solu-
tion pH is an especially important governing parameter, as is the case for adsorption of all ions
on oxidic minerals [13]. The data presented also demonstrate that the presence of free cyanide in
© 2006 by Taylor & Francis Group, LLC
FeOOH(s). The pH adsorption edge plots shown in Figure 5.9 exhibit the typical characteristics for
the metal concentrations were also performed. Results for the three metals are given in Figure 5.10.
Figure 5.5). The data in Figure 5.10 indicate that the cadmium–, copper–, and zinc–cyanide species
Physical–Chemical Properties and Reactivity 71

3456789
10
11
0
30
20
10
50
40
80
70
60
100
90
Percent nickel removed
pH
FIGURE 5.8 Nickel adsorption as a function of pH in the presence of iron oxide (goethite) and cyanide.
Ni
T
= 10
−4.77
M, I = 0.01 M, α-FeOOH = 0.59 g/L. ()CN
T
= 10
−5
M,()CN
T
= 10
−4
M,

()CN
T
= 10
−3
M,(•)CN
T
= 0. (Source: Reprinted with permission from Theis, T.L. and Richer, R.O.,
Particulates in water, 189, 73, 1980. Copyright 1980 American Chemical Society.)
a systems with metals, leading to the formation of metal–cyanide complexes, can result in enhanced
or reduced adsorption of the metals. The metal–cyanide complexes may interact with the surface to
a greater or lesser extent than the metals alone. An interrelated, complex group of factors governs
metal–cyanide species adsorption, and it is diffcult to form generalizations.
5.2.1.4 Oxidation
Weak metal–cyanide complexes generally arereadilyoxidizedby oxidizing agents such as chlorine or
ozone. Themore strongly bonded complexes in theWAD 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.
most common approach used to treat waters bearing free cyanide. A number of weak metal–cyanide
complexes are also readily oxidized in this process. In order to identify the fraction of measured
total cyanide, which includes metal–cyanide complexes and free cyanide, that is treatable by alkaline
© 2006 by Taylor & Francis Group, LLC
As 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 •FeOOH = 0.6 g/L
I = 0.1
M
10
20
30
40
50
60
70
80
90
Percent removed
0
100
34567891011
pH
FIGURE 5.9 Adsorption of Cd, Cu, and Zn on goethite. I = 0.01 M, α-FeOOH = 0.6 g/L. ()Cu
T
=
10
−4.72

M,()Zn
T
= 10
−4.62
M,()Cd
T
= 10
−4.68
M.(Source: Theis, T.L. and West, M.L., Environ.
Technol. Lett., 7, 309, 1986.)
0
10
20
30
40
50
60
70
80
90
100
34567891011
pH
Percent removed
Zn
T
= 1.73 ϫ10
Ð 5
M/ CN
T

=2.00ϫ10
Ð 3
M
Cu
T
= 1.95 ϫ10
Ð 5
M/ CN
T
= 1.00 ϫ10
Ð 4
M
Cd
T
= 1.88 ϫ10
Ð 5
M/ CN
T
= 2.00 ϫ10
Ð 3
M
a• FeOOH = 0.6 g/L
I = 0.1M
FIGURE 5.10 Effect of cyanide on adsorption of Cd, Cu, and Zn on goethite. I = 0.01 M, α-FeOOH =
0.6 g/L. ()Cd
T
= 10
−4.73
M,CN
T

= 10
−2.70
M;()Zn
T
= 10
−4.76
M,CN
T
= 10
−2.70
M;()Cu
T
=
10
−4.71
M,CN
T
= 10
−4.00
M.(Source: Theis, T.L. and West, M.L., Environ. Technol. Lett., 7, 309, 1986.)
chlorination, an analytical measurement known as “cyanide amenable to chlorination” [15] has long
been employed. The method involves measurement of total cyanide on samples with and without
treatment by chlorination, with the difference giving the amount of cyanide in the sample amenable
existence of the method speaks to the facile oxidation of a number of weak metal–cyanide complexes.
In somecases, the presence of weak metal–cyanidecomplexescanenhancethe rate of free cyanide
decomposition through catalysis by the metal. This has been demonstrated for copper cyanide com-
plexes [20]. Gurol and Holden [20] studied the effect of copper(I) on the removal of free cyanide by
© 2006 by Taylor & Francis Group, LLC
to chlorination (Chapter 7). While the CATC method has limitations, as discussed in Chapter 7, the
Physical–Chemical Properties and Reactivity 73

ozone in alkaline solution. They performed experiments in solutions at pH 11.5 in systems with an
excess of free cyanide over copper(I), giving Cu(CN)
3−
4
as the dominant copper species. They found
that the presence of copper increased the rate of free cyanide oxidation significantly. Comparison of
initial rates of cyanide disappearance for systems with and without copper indicated a fivefold higher
rate for the system with copper cyanide species. Further investigations revealed that the observed
enhancement was likely due a very fast, independent oxidation–reduction reaction between Cu(I)
and free cyanide. The following reaction sequence was proposed:
2Cu(CN)
3−
4
+2O
3
→ 2Cu(CN)
2−
3
+2CNO
−
+2O
2
(5.28)
2Cu(CN)
2−
3
+O
3
+2H
+

→ 2Cu(CN)
−
3
+O
2
+H
2
O (5.29)
2Cu(CN)
−
3
→ 2Cu(CN)
−
2
+(CN)
2
(5.30)
2Cu(CN)
−
2
+4CN
−
= 2Cu(CN)
3−
4
(5.31)
(CN)
2
+2OH
−

→ CN
−
+CNO
−
+H
2
O (5.32)
The reaction in Equation (5.28) represents the direct oxidation of cyanide to cyanate by ozone. In
Equation (5.29), Cu(I) is oxidized to Cu(II) in a fast reaction. The Cu(II) species subsequently oxid-
izes cyanide to cyanogen (C
2
N
2
), being reduced back to Cu(I) in the process (Equation [5.30]).
An equilibrium between the copper(I) cyanogen species and Cu(CN)
3−
4
is rapidly established
(Equation [5.31]). In the last step, Equation (5.32), cyanogen goes through a disproportionation
reaction to yield free cyanide and cyanate. The net reaction from the above sequence is thus as
follows:
2Cu(CN)
3−
4
+3CN
−
+3O
3
→ 2Cu(CN)
3−

4
+3CNO
−
+3O
2
(5.33)
Thus, the oxidation of 3 mol free cyanide requires 2 mol of ozone and produces 3 mol of cyanate, as
would be expected, but the rate of the reaction is much accelerated due to the presence of the Cu(I).
5.2.2 STRONG METAL–CYANIDE COMPLEXES
5.2.2.1 Formation
The cyanide anion can form strong complexes with a number of transition heavy metals, the most
notable among them are cobalt, platinum, gold, palladium, and iron. Most of these metals fall in
Groups IB, IIB, and VIII of the periodic table. As iron is by far the most abundant of these elements in
the environment and in process waters, iron–cyanide complexes are the strong metal–cyanide com-
plexes of greatest interest. Gold–cyanide complexes are of great commercial interest, as the strength
of the gold–cyanide bond is exploited in hydrometallurgical gold mining for aqueous extraction of
arranged in tetrahedral or octahedral forms with strong bonding energy existing between the heavy
metal atom and the cyanide ligand [1]. Because they can only dissociate under strongly acidic
pH conditions (pH < 2), they are referred to as strong acid dissociable complexes, or simply as
strongly-complexed cyanide [15].
As some of these species are formed very slowly [1,33], it is difficult to determine the equilibrium
formation constants. Formation data determined by direct thermodynamic methods are available only
for complexes of gold(I), and palladium(II) [1]. For other metals, like iron, electron transfer between
complex ions of the element in the +2 and +3 oxidation states is rapid enough such that the ratio of
the formation constants can be determined from measurement of redox potentials. This ratio can then
be combined with standard enthalpy and entropy change measurements for the formation reaction
of interest.
© 2006 by Taylor & Francis Group, LLC
gold from ores (see Chapters 4 and 26). The metal–cyanide bonds in these complexes are mostly
74 Cyanide in Water and Soil

TABLE 5.4
Equilibrium Constants for Formation of Selected Strong
Metal–Cyanide Complexes
Reaction log K (at 25
◦
C, I = 0)
Ba
2+
+Fe
2+
+6CN
−
= BaFe(CN)
2−
6
49.40
Ba
2+
+Fe
3+
+6CN
−
= BaFe(CN)
−
6
55.44
Ca
2+
+Fe
2+

+6CN
−
= CaFe(CN)
2−
6
49.69
Ca
2+
+Fe
3+
+6CN
−
= CaFe(CN)
−
6
55.47
2Ca
2+
+Fe
2+
+6CN
−
= Ca
2
Fe(CN)
0
6
51.00
Ca
2+

+H
+
+Fe
2+
+6CN
−
+e
−
= CaHFe(CN)
2−
6
52.71
Fe
2+
+6CN
−
= Fe(CN)
4−
6
45.61
Fe
2+
+H
+
+6CN
−
= HFe(CN)
3−
6
50.00

Fe
2+
+2H
+
+6CN
−
= H
2
Fe(CN)
2−
6
52.45
Fe
3+
+6CN
−
= Fe(CN)
3−
6
52.63
2Fe
2+
+6CN
−
= Fe
2
(CN)
0
6
56.98

2K
+
+Fe
2+
+2H
+
+6CN
−
= K
2
H
2
Fe(CN)
0
6
52.31
3K
+
+Fe
2+
+H
+
+6CN
−
= K
3
HFe(CN)
0
6
50.22

K
+
+Fe
2+
+6CN
−
= KFe(CN)
3−
6
48.12
2K
+
+Fe
2+
+6CN
−
= K
2
Fe(CN)
2−
6
48.98
K
+
+Fe
2+
+H
+
+6CN
−

= KHFe(CN)
2−
6
51.47
Li
+
+Fe
2+
+6CN
−
= LiFe(CN)
3−
6
47.69
2Li
+
+Fe
2+
+6CN
−
= Li
2
Fe(CN)
2−
6
48.53
Li
+
+Fe
2+

+H
+
+6CN
−
= LiHFe(CN)
2−
6
51.22
Mg
2+
+Fe
3+
+6CN
−
= MgFe(CN)
−
6
55.39
Mg
2+
+Fe
2+
+6CN
−
= MgFe(CN)
2−
6
49.43
NH
+

4
+Fe
2+
+6CN
−
= NH
4
Fe(CN)
3−
6
48.07
2NH
+
4
+Fe
2+
+6CN
−
= (NH
4
)
2
Fe(CN)
2−
6
48.87
NH
+
4
+H

+
+Fe
2+
+6CN
−
= NH
5
Fe(CN)
2−
6
51.40
Na
+
+Fe
2+
+6CN
−
= NaFe(CN)
3−
6
47.99
2Na
+
+Fe
2+
+6CN
−
= Na
2
Fe(CN)

2−
6
48.74
Na
+
+Fe
2+
+H
+
+6CN
−
= NaHFe(CN)
2−
6
51.43
Sr
2+
+Fe
3+
+6CN
−
= SrFe(CN)
−
6
55.62
Tl
+
+Fe
2+
+6CN

−
= TlFe(CN)
3−
6
48.75
Source: Data from Sehmel, G.A., Cyanide and antimony thermodynamic data-
base for the aqueous species and solids for the EPA-MINTEQ geochemical code,
PNL-6835, Pacific Northwest Laboratory, Richland, WA, 1989, (Table 5).
Table 5.4 lists the equilibrium constants for the reversible formation of iron–cyanide complexes,
which are of primary interest with respect to cyanide in the environment. The constants reported are
from the compilation by Sehmel [5], which was selected for the reasons discussed in Section 5.2.1.1.
Among all the iron–cyanide complexes, the most commonly occurring are ferrocyanide, Fe(CN)
4−
6
,
where iron is the +2 oxidation state, and ferricyanide, Fe(CN)
3−
6
, where iron is in the +3 oxidation
state. Another iron–cyanide complex only recently identified and not presented in Table 5.4, with
a chemical formula, Fe(CN)
5
NHCH
4−
3
, has been found to dominate groundwater at many former
manufactured gas plant sites [39].
© 2006 by Taylor & Francis Group, LLC
Physical–Chemical Properties and Reactivity 75
FIGURE 5.11 Predominance diagram for dissolved cyanide species in equilibrium with hydrous ferric oxide

at T = 25
◦
C, as calculated with MINEQL+ (Schecher et al., 1998) using the reactions in Equation (5.1)
0.06 mM, and I = 0.06 M NaCl. (Source: Ghosh, R.S. et al., Water Environ. Res., 71, 1205, 1999.)
In soils and aquifer systems where iron is ubiquitous, the aqueous speciation of cyanide is
influenced significantly by reactions with iron dissolved from iron oxides [40,41]. Equilibrium with
hydrous ferric oxide, the common amorphous iron oxide, typically is important because Fe(OH)
3
(s)
serves as the source of iron that becomes dissolved, which 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 the area denoted
“Fe(CN)
6
(tot),” cyanide is predicted to exist at equilibrium primarily as the iron cyanide species
Fe(CN)
3−
6
(oxic conditions) or Fe(CN)
4−
6
(anoxic conditions). In the remaining area HCN is the
predominant form of dissolved cyanide, except for a small region at pH > 9.2, the pK
a
for HCN,
above which CN
−
dominates free cyanide speciation.

5.2.2.2 Dissociation
As indicated in the equilibrium species predominance diagram of Figure 5.11, iron–cyanide com-
plexes require acidic conditions to dissociate and form free cyanide. It is important to remember,
however, that the species distribution shown in Figure 5.11 reflects equilibrium conditions. Actual
species distributions for systems with iron present are strongly governed by kinetics.
Dissociation of iron–cyanide complexes in the dark is very slow [42]. Like weak metal–cyanide
complexes, the dissociation properties of iron–cyanide complexes in aqueous solutions are functions
of their stability constants, pH, temperature, and redox potential of the solution. Meeussen et al. [42]
studied the dissociation of ferrocyanide, Fe(CN)
4−
6
,in1mM solutions in the dark at 15
◦
C. Based on
the results, Meeussen et al. [42] projected half-lives ranging from 1 year under reducing conditions
(pE ≈ 5) at pH 4 to 1000 years at the same pH under oxidizing conditions (pE ≈ 10). In these
laboratory experiments, maximum decompostion rates were observed at low pH and pE. Actual
decomposition rates in the environment could be quite different, for example, through enhancement
by catalysts. Nevertheless, the results of Meeussen et al. demonstrate the high degree of stability of
© 2006 by Taylor & Francis Group, LLC
and Table 5.4. TOTCN = 0.6 mM, TOTFe = 0.5 mM,TOTK= 0.4 mM, TOTNa = 0.06 mM, TOTCl =
25
20
15
10
HCN
Water oxidized
Water reduced
Fe(CN)
6

(tot)
CN
–
5
pE
0
–5
–10
–15
0246
pH
810
76 Cyanide in Water and Soil
FIGURE 5.12 Ferro- (a) and ferricyanide (b) photodissociation reaction pathways. (Sources: Information
from: Gaspar, V. and Beck, M.T., Polyhedron, 2, 387, 1983; Fuller, M.W., Aust. J. Chem., 39, 1411, 1986.)
iron–cyanide complexes and the importance of considering kinetics when evaluating dissociation of
these complexes. Since strong acid conditions are required to dissociate these complexes, the total
cyanide analytical method, which is designed to recover strong metal–cyanide complexes in addition
to free cyanide and weak metal–cyanide complexes, employs strong acid pH conditions (pH = 1.5)
and heat (125
◦
C, for 2 h) to achieve dissociation of all metal–cyanide complexes present ([15,43];
While ferro- and ferricyanide complexes are quite stable in the dark, they can dissociate rapidly
when exposed to light [42,44–46]. For example, in experiments with 1.28 µM ferrocyanide solutions
at pH 12 exposed to diffuse daylight, Meeussen et al. [42] observed an initial decomposition rate
of approximately 8% per hour. Light in the ultraviolet (UV) range (wavelengths less than 420 nm)
is responsible for the photolysis of ferro- and ferricyanide [45]. Some proposed photodissociation
pathways for ferro- and ferricyanide are shown in Figure 5.12.
The photoactivated dissociation rate depends on light intensity, light wavelength, temperature,
presence of catalysts, and other parameters [45,47–49], but the results of Meeussen et al. [42]

demonstrate clearly that photodissociation of iron–cyanide complexes upon exposure to natural light
can be very rapid. Based on photolysis experiments with 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 degradationofferro-andferricyanide followsapproximatelyfirst-orderkinetics [45,49], at
least initially, but the rate slows as free cyanide accumulates in solution [49]. While some differences
in the ratesof photolysis of hexacyanoferrate have been observed at different pH values [45,49], Kuhn
and Young [49] found no consistent pattern of initial rate coefficient dependence on pH in studies
on solutions at pH 4 to 12. The presence of natural organic matter, or other photoreactive substances
in water, can significantly decrease the rate of hexacyanoferrate photolysis [49].
The rate of photochemical 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 produced [45]. Free cyanide could possibly be undetectable or short-
lived [51]. There are scenarios, however, where sunlight intensity may be uniform across the entire
water column, such as in shallow ponds and in surface runoff. Hexacyanoferrate contamination of
the latter can occur, for example, through the spreading of road salt containing iron cyanide as an
anticaking agent [52], that is, the commonly used “blue salt.” More details on the photolysis of iron
cyanide species and the role of photolysis in fate and transport of hexacyanoferrate in surface waters
5.2.2.3 Adsorption on Soil and Sediment
Like the weak metal–cyanide complexes (Section 5.2.1.3), strong complexes such as ferrocyanide,
Fe(CN)
4−
6
, and ferricyanide, Fe(CN)
3−
6
, can adsorb on common soil and sediment components
© 2006 by Taylor & Francis Group, LLC

Chapter 7).
are provided in Chapter 9.
[Fe(CN)
5
H
2
O]
3–
+ CN
–
[Fe(CN)
6
]
4
*
k
6
k
-
6
I
c
k
r
k
f
K
d
HCN + OH
–

k
6
k
-
6
I
c
k
r
k
f
[Fe(CN)
*
h
OH
–
[Fe(CN)
5
(OH)]
3–
h
…
?
approx. 0.98
approx.
0.02
[Fe(CN)
5
H
2

O]
2–
+ CN
–
[Fe(CN)
6
]
3–
hn
hn
…
(a) (b)
[Fe(CN)
6
]
4–
[Fe(CN)
6
]
3–
[Fe
2
(CN)
10
]
6–
Fe(OH)
3
[Fe
2

(CN)
10
]
5–
[Fe2(CN)
10
]
5–
Physical–Chemical Properties and Reactivity 77
6789
0
25
50
75
100
3.3 mg g
– 1
1.7 mg g
– 1
0.83 mg g
– 1
0.50 mg g
– 1
Percent Fe(CN)
6
– 4
adsorbed
5 10
pH
FIGURE 5.13 Ferrocyanide adsorption on alumina at different adsorbate/adsorbent ratios (mg Fe(CN)

4−
6
as
CN per g γ-Al
2
O
3
(s)).10
−5.19
M Fe(CN)
4−
6
and various γ-Al
2
O
3(s)
solid doses in 100 mL 0.01 M NaCl
solution. (Source: Bushey, J.T. and Dzombak, D.A., J. Coll. Int. Sci., 272, 46, 2004, John Wiley & Sons.
Reproduced with permission.)
such as iron, aluminum, and manganese oxides, and clays [8,12,53]. Adsorption of metal–cyanide
complexes occurs through a combination of electrostatic attraction and surface complexation [8]. As
most strong metal–cyanide complexes are anionic, they can be substantially adsorbed onto soils with
high anion exchange capacity. Solution conditions, especially pH, also affect the extent of adsorption
of metal–cyanide complexes in aqueous systems.
Alessi and Fuller [12] conducted laboratory column mobility tests in which ferricyanide solution
was passed through five different soils of varying physical and chemical characteristics. Based on
these tests, itwas concluded that soil properties, suchaslow pH(pH < 5), free iron oxide content, and
kaolin, chlorite, and gibbsite type clay (high anion exchange capacity) material increased adsorption
of iron–cyanide complexes to soil material.
Conversely, soils, sediments, and aquifer materials dominated by sand or other components

with high cation exchange capacities tend to be weaker adsorbents for iron–cyanide complexes
[40]. Mobility tests performed in fixed-bed columns packed with sand-dominated aquifer material
and ferrocyanide-contaminated site groundwater as the mobile phase revealed minimal interaction
between the dissolved ferrocyanide complexes and site sand [40]. In a mobility test performed by
Ghosh et al. [40], ferrocyanide was observed to break through the column in one pore volume, similar
to transport of a conservative tracer.
It has been demonstratedin a number of studies[8,53,54] that aluminum and ironoxides, two very
common and surface-reactive components of soils and sediments [6,55], can adsorb iron–cyanide
species significantly, especially at lower pH values (<7). Bushey and Dzombak [53] studied the
equilibrium adsorption of ferrocyanide on the aluminum oxide, γ-Al
2
O
3
, and found that adsorp-
tion increased as pH decreased (Figure 5.13). At pH < 7 and for all lower adsorbate/adsorbent
ratios (given as mg Fe(CN)
4−
6
as CN per g γ-Al
2
O
3
) examined, removal of ferrocyanide from
solution was essentially complete. As adsorbate/adsorbent ratio was increased, reflecting a greater
loading of Fe(CN)
4−
6
relative to available solid mass and surface area in the system, the percent
of the total Fe(CN)
4−

6
adsorbed decreased. The adsorptive characteristics of Fe(CN)
4−
6
shown
© 2006 by Taylor & Francis Group, LLC
78 Cyanide in Water and Soil
34 5 678 91011
pH
a– FeOOH = 0.6 g/L
I =0.1 M KNO
3
Percent Fe(CN)
6
3–
adsorbed
0
50
30
20
100
70
10
40
60
80
90
2 × 10
– 5
M Fe(CN)

6
– 3
FIGURE 5.14 Ferricyanide adsorption on goethite. I = 0.01 M, α-FeOOH = 0.6 g/L, 10
−4.70
M Fe(CN)
3−
6
.
(Source: Theis, T.L. and West, M.L., Environ. Technol. Lett., 7, 309, 1986.)
oxides [13].
Theis and West [8] studied ferricyanide, Fe(CN)
3−
6
, adsorption on goethite (α-FeOOH). Based
on their work, it is evident that ferricyanide can adsorb substantially onto the goethite surface
under acidic conditions (see Figure 5.14). Rennert and Mansfeldt [54] also studied the adsorption
of ferricyanide, as well as ferrocyanide, on goethite. Ferrocyanide adsorbed to a somewhat greater
extent than ferricyanide, suggesting that ferrocyanide could be less mobile than ferricyanide in a soil
system. Further, Rennert and Mansfeldt [56] investigated the effect of sulfate, an adsorbing anion that
competes for adsorption sites on oxides [13], on the adsorption of the iron–cyanide complexes. Over
the studied range of pH 3.5 to 8, iron–cyanide complex adsorption was strongly dependent on sulfate
concentration and vice versa. iron–cyanide complex adsorption was decreased by the presence of
the sulfate, especially at lower pH values.
Another study by Rennert and Mansfeldt [57] evaluated the influence of different soil properties
on iron–cyanide complex adsorption in soils. This study concluded that adsorption of both iron–
cyanide complexes is dependent upon soil organic carbon content, clay, oxalate-extractable Fe (free
iron) and oxalate-extractable Al. The latter two parameters are indirect measures of the iron and
aluminum oxide content of the soil. For soils with low organic carbon content (<10 g/kg), the
adsorption of ferro- and ferricyanide was found to be governed by pH and clay content. For soils
with high organic carbon content, clay and oxalate-extractable Al, and pH, were found to regulate

adsorption behavior for ferricyanide, whereas ferrocyanide adsorption was regulated only by oxalate-
extractable Al and pH under the same soil conditions. The findings demonstrated that soil organic
matter can have an important role in enhancing the adsorption of both iron–cyanide complexes,
possibly by reaction between the iron–cyanide nitrogen and reactive functional groups of surface
organic matter (e.g., quinone).
5.2.2.4 Oxidation–Reduction
Ferricyanide can be readily reduced to ferrocyanide by a variety of reducing agents. The
Fe(CN)
3−
6
/Fe(CN)
4−
6
redox couple has been used to oxidize phenolic compounds, such as resorcinol,
© 2006 by Taylor & Francis Group, LLC
in Figure 5.13 are similar to those observed generally for adsorption of anionic species on metal
Physical–Chemical Properties and Reactivity 79
for example [58]:
2C
6
H
6
O
2
+2OH
−
+2Fe(CN)
3−
6
→ C

12
H
10
O
4
+2Fe(CN)
4−
6
+2H
2
O (5.34)
Ferricyanide has also been used as a redox titrant to investigate oxidation–reduction properties of
organic compounds and mixtures of organic compounds, including humic acids [59]. For example,
ferricyanide has been used to study the one-electron transfer of hydroquinone at low pH:
QH
2
+2Fe(CN)
3−
6
→ Q + 2H
+
+2Fe(CN)
4−
6
(5.35)
where QH
2
and Q represent hydroquinone and quinone species, respectively [59].
Ferrocyanide, Fe(CN)
4−

6
, can be oxidized to ferricyanide, Fe(CN)
3−
6
, by molecular oxygen in
the dark, but the kinetics of this reaction are slow. Asperger et al. [60] determined the first order rate
constant for this reaction to be 10
−4
sec
−1
in the dark at 40
◦
C and pH 4.5.
In the absence of UV light, onlyvery strong oxidants like ozone, persulfate, andpermanganatecan
oxidize ferrocyanide ion, in acidic solutions, to ferricyanide [61]. It is difficult to oxidize ferrocyanide
under neutral to alkaline pH conditions without UV light. Under oxidant concentrations in excess
of stoichiometric requirements, however, ferrocyanide would decompose to CO
2
and other benign
end products at a very slow rate. Gurol and Holden [20] showed, for example, that oxidation of
one mole of iron-complexed cyanide to CO
2
required in excess of 30 mol of ozone. Similarly, the
presence of excess permanganate can decompose any ferricyanide ion formed from ferrocyanide to
CO
2
, ferric oxide, and nitrogen oxides [61]. Further discussion of ambient temperature oxidation of
To oxidize ferrocyanide and ferricyanide most efficiently, ultraviolet light can be employed to
photodissociate the complexes, so that lower doses of oxidizing agents will be effective in oxidizing
the free cyanide released. This has been demonstrated by Schaefer [18], and the process is examined

in depth in Chapter 20. Consider, for example, the photocatalytic dissociation of ferrocyanide and
subsequent oxidation of free cyanide and cyanate [18]:
Fe(CN)
4−
6
+3H
2
O + hυ → 6CN
−
+Fe(OH)
3
(s) + 3H
+
+e
−
(5.36)
CN
−
+oxidant → CNO
−
(5.37)
CNO
−
+oxidant → NO
−
3
+CO
2
(5.38)
All of these oxidation reactions follow first-order kinetics [18].

5.3 METAL–CYANIDES: SOLID PHASE COMPOUNDS
metal–metal cyanide complex salts. Solids form upon reaction of free cyanide or metal–cyanide spe-
cies with metal ions in solution, which results in solid precipitation when the reactants are sufficiently
abundant. These solids form and dissolve in natural aquatic systems. Salts of many metal–cyanide
species are also synthesized by various routes and sold commercially [62,63]. This section provides
detailed information on the physicochemical properties and reactivities of these metal–cyanide
species.
Solid forms of cyanide can be divided into three forms according to their properties and react-
ivity. They are (i) simple cyanide solids, like NaCN(s), KCN(s), Ca(CN)
2
(s), etc., (ii) alkali or
alkaline earth metal–metal cyanide complex solids, like K
3
[Fe(CN)
6
](s),Na
3
[Fe(CN)
6
]·H
2
O(s),
etc., and (iii) other metal–metal cyanide complex solids, especially the iron–iron cyanide solids like
Fe
4
[Fe(CN)
6
]
3
(s) and Fe

3
[Fe(CN)
6
]
2
(s), which are of special interest in water and soil systems.
© 2006 by Taylor & Francis Group, LLC
cyanide complexes is provided in Chapter 20.
As outlined in Chapter 2, many solid forms of metal–cyanide species occur, including double
80 Cyanide in Water and Soil
5.3.1 SIMPLE CYANIDE SOLIDS
These forms of cyanide solids include a metal or ammonium ion bonded ionically to the cyanide
anion. These simple cyanide salts generally dissociate readly in aqueous solution to form free cyanide
anion, according to the following general reaction:
M(CN)
x
= M
x+
+xCN
−
(5.39)
where M represents a metal cation with charge x.
Simple cyanides are named because of their structural simplicity and their ability to dissociate
substantially under most solution conditions. However, some simple cyanides, where the cyanide ion
is bonded to a heavy metal, like CuCN(s) and Zn(CN)
2
(s), are somewhat less soluble than others.
Most simple cyanide solids release the free cyanide ion readily in aqueous solution.
characteristics. Solubility products available in Sehmel [5] forsomeof the less soluble simple cyanide
constants have been calculated in a consistent manner and with the most current thermodynamic data

5.3.2 ALKALI OR ALKALINE EARTH METAL–METAL CYANIDE
COMPLEX SOLIDS
These forms of cyanidesolidsincludeanalkali(Na
+
,K
+
) or alkaline earth (Mg
2+
,Ca
2+
,Sr
2+
) metal
cation, or ammonium ion (NH
+
4
), bonded to a metal–cyanide complex. Some of these solids also
carry one or more water of hydration within the crystal lattice structure, for example, K
4
(FeCN)
6
)·
3H
2
O(s). These types of complexed cyanide solids are moderately soluble in aqueous solutions
under a wide range of pH conditions, releasing metal–cyanide complexes according to the following
general reaction:
A
x
[M(CN)

y
]·nH
2
O = xA
+
+[M(CN)
y
]
m−
(5.40)
where A representsan alkali or alkaline earth cation or ammonium ion, M is a metalcation complexed
with cyanide, x is the number of alkali metal atoms (Equation [5.40] is written for A with charge +1),
and n is the number of water molecules incorporated in the solid structure. Complex cyanide solids
are so named because of their structural complexity and their incomplete dissociation under most
solution conditions. The complexed metal–cyanide salts do not dissociate completely, and cyanide
is liberated in the form of dissolved metal–cyanide species.
solubility characteristics. Solubility products available in Sehmel [5] for some alkali and alkaline
Sehmel [5] are given because these constants have been calculated in a consistent manner and with
5.3.3 OTHER METAL–METAL CYANIDE COMPLEX SOLIDS
These forms of complex cyanide solids contain one or more transition metals or soft-sphere B-type
metal cations [6] bonded to a transition metal–cyanide complex. Some compounds in this class
are also referred to as “double metal cyanide complex salts.” The best known such compound is the
ferric ferrocyanide solid known as Prussian Blue, Fe
4
(Fe(CN)
6
)
3
(s). Most strong metal–cyanide
complexes form stable precipitates under acidic to neutral pH and under excess metal conditions.

The iron–iron cyanide solid compounds, are of special importance
in water and soil systems. Where cyanide has been introduced in water and soil systems, these
© 2006 by Taylor & Francis Group, LLC
Table 5.5 provides alistofsomesimplecyanide solids, andsomegeneralinformationonsolubility
(see discussion in Section 5.2.1.1).
salts are given in Table 5.6. Only solubility products given in Sehmel [5] are given because these
Table 5.7 provides a list of some complex cyanide solids, and some general information on
earth metal–metal cyanide complex solids are given in Table 5.8. Only solubility products given in
presented in Table 5.9,
y is the number of cyanide groups, m is the ionic charge of the liberated metal–cyanide complex,
the most current thermodynamic data (see discussion in Section 5.2.1.1).
Physical–Chemical Properties and Reactivity 81
TABLE 5.5
Some Simple Cyanide Solids and Solubility Information
Aqueous solubility
d
Solid Formula Physical form
a
(g/100 g H
2
O) Temp (
◦
C)
Qualitative
solubility
a,c
Barium cyanide Ba(CN)
2
wh cry pow vs H
2

O; s EtOH
Cadmium cyanide Cd(CN)
2
wh cub cry 1.7
a
15
Calcium cyanide Ca(CN)
2
wh rhom cry; hyg s H
2
O, EtOH
Cyanogen bromide CNBr wh hyg needles
Cesium cyanide CsCN wh cub cry; hyg vs H
2
O
Cobalt(II) cyanide Co(CN)
2
blue hyg cry i H
2
O
Cobalt(II) cyanide
dihydrate
Co(CN)
2
·2H
2
O pink-brn needles i H
2
O, acid
Copper(I) cyanide CuCN wh pow; or i H

2
O, EtOH;
grn orth cry s KCN soln
Copper(II) cyanide Cu(CN)
2
grn pow i H
2
O; s acid, alk
Gold(I) cyanide AuCN yel hex cry i H
2
O, EtOH, eth,
dil acid
Gold(III) cyanide
trihydrate
Au(CN)
3
·3H
2
O wh hyg cry vs H
2
O; sl EtOH
Mercury(II) cyanide Hg(CN)
2
col tetr cry 11.4
a
25 s EtOH; sl eth
Nickel(II) cyanide
tetrahydrate
Ni(CN)
2

·4H
2
O grn plates i H
2
O; s dil acid;
sNH
4
OH
Potassium cyanide KCN wh cub cry; hyg 69.9
a
20 sl EtOH
Rubidium cyanide RbCN wh cub cry s H
2
O; i EtOH, eth
Silver cyanide AgCN wh-gray hex cry 0.0000011 i EtOH, dil acid
Sodium cyanide NaCN wh cub cry; hyg 58.2
a
, 58.7
b
20 sl EtOH
Sodium
cyanoborohydride
NaBH
3
CN wh hyg pow vs H
2
O; s, thf;
sl EtOH; i bz, eth
Strontium cyanide
tetrahydrate

Sr(CN)
2
·4H
2
O wh hyg cry vs H
2
O
Thallium(I) cyanide TlCN wh hex plates s H
2
O, acid, EtOH
Zinc cyanide Zn(CN)
2
wh pow 0.00047
a
20 reac acid
a
Source: Data from Lide, D.R., CRC Handbook of Chemistry and Physics, 85th ed. (online edition), CRC Press, Boca Raton,
FL, 2004.
b
Source: Data from Dean, J.A., Lange’s Handbook of Chemistry, 15th ed. (online edition), McGraw-Hill, New York, 1999.
c
Abbreviations (after Lide [82]): ace: acetone; acid: acid solutions; alk: alkaline solutions; aq: aqueous; bz: benzene; col:
colorless; cry: crystals; cub: cubic; dil: dilute; eth: ethyl ether; EtOH: ethanol; grn: green; hex: hexagonal; hyg: hygroscopic;
i: insoluble in; orth: orthorhombic; pow: powder; reac: reacts with; rhom: rhombohedral; s: soluble in; sl: slightly soluble
in; soln: solution; temp: temperature; tetr: tetragonal; thf: tetrahydrofuran; tol: toluene; vs: very soluble in; wh: white; yel:
yellow.
d
Blanks indicate no values reported in Lide [81] or Dean [82].
compounds often form because of their low solubility and the abundance of iron in the environment
[10,40,41,64–67]. Moreover, iron–iron cyanide solids, most notably Prussian Blue, have been pro-

duced for use in commercial products such as dyes, inks, pharmaceuticals, and cosmetics for over
150 years [62,63].
with the primary distinguishing feature being the ratio of Fe
2+
and Fe
3+
in the crystalline structure.
The best-known iron–iron cyanide solids are Prussian Blue, Fe
4
(Fe(CN)
6
)
3
(s), and Turnbull’s Blue,
Fe
3
(Fe(CN)
6
)
2
(s), which have different proportions of Fe
2+
and Fe
3+
; Berlin White or Williamson’s
Salt, Fe
2
(Fe(CN)
6
)(s), which contains all Fe

2+
; and iron–iron cyanide solids containing predom-
inantly Fe
3+
, which exhibit colors ranging from brown (Prussian Brown) to green (Berlin Green),
© 2006 by Taylor & Francis Group, LLC
As indicated in Table 5.9, various iron–iron cyanide solids have been reported in the literature,