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Copper and zinc equilibria in concentrated ammonia solutions: Relevance to stress corrosion cracking of alpha-brass
Corrosion Science, Vol. 39, No. 7, pp. 1307-1319, 1997 0 1997 Elsevier Science Ltd
Pergamon
Printed in Great Britain. All rights reserved 0010-938X/97 $17.00+0.00
PII: S001&938X(97)00030-9
COPPER AND ZINC EQUILIBRIA IN CONCENTRATED AMMONIA SOLUTIONS: RELEVANCE TO STRESS CORROSION CRACKING OF ALPHA-BRASS D. TROMANS University
of British Columbia,
Department of Metals and Materials Engineering, British Columbia, V6T 124, Canada
6350. Stores Road, Vancouver,
Abstract-Potential-pH diagrams for the Cu-NHx-Hz0 and Zn-NH,-Hz0 systems have been constructed for 15 M ammonia solutions commonly used for studying SCC of a-brass. Equilibria and pH calculations were based on activities of NH3 and water that were 37.4 and 0.68, respectively. The equilibrium diagrams are in agreement with the reported behaviour of brass in copper-containing 15 M ammonia, and with the active behaviour of zinc, They are the most useful diagrams to date for studies relating to SCC of a-brass in tarnishing and non-tarnishing 15 M ammonia solutions. Anodic polarisation tests indicated that formation of the zinc-tetrammine complex in 15 M ammonia may involve transient zincate ions. 0 1997 Elsevier Science Ltd Keywords:
A. brass, A. copper,
A. zinc, B. polarisation.
INTRODUCTION The stress corrosion cracking (KC) of a-phase Cu-Zn alloy (brass) in ammoniacal environments has received considerable attention, partly because the system is useful for developing and testing mechanistic models of environment-assisted crack propagation. In contrast, the relevant potential (E) and pH equilibria of copper and zinc have received less consideration, particularly in highly concentrated ammonia solutions. Some of the earliest E-pH equilibrium diagrams for the Cu-NH3-Hz0 system were constructed by Halpern,’ Mattsson,’ and Johnson and Leja3 for 1 M of total dissolved ammonia species (NH3 + NHd+), where the molarity was equated with activity. However, when considering the formation of Cur-ammine and Cu”-ammine complexes, and the stability of the copper oxides, these workers did not completely recognise the consequences of the equilibrium between the ionised and un-ionised forms of dissolved ammonia, NH;
= NH3 + H+, kl = 5.628 x lo-“,
pk, = 9.25
(1)
where k, is the equilibrium constant for eqn (1) at 25°C and pk, = -log k,. Consequently, below pH 9.25 aqueous equilibria between dissolved ammonia and copper species should be determined with respect to NH4+ whereas above pH 9.25 the equilibria should be determined with respect to NH3. Halpern,’ and Johnson and Leja3 determined copper oxide stability relative to NH 3, whereas Mattsson’ considered only NH4+. Thus, their E-pH diagramsIp were only partially correct. Letowski and Niemic,4
Manuscript
received 20 June 1996; in amended
form 14 February 1307
1997.
1308
D. Tromans
and Bartonicek and Lukasovska’ considered both forms of ammonia and showed that the solublities of the copper oxides reached a maximum when the pH was equal to pk’ ,4 leading to oxide stability domains in pH regions above and below 9.25.5,6 Hoar and Rothwell’ also considered both forms of ammonia and constructed E-pH diagrams for copper in 1 M total ammonia (assumed to be unit activity) that displayed two regions of oxide stability. Although Mattsson’s2 E-pH diagram was only partially correct, his work drew attention to the influence of cupric-ammine complexes and Cu20 formation on SCC of c1brass in near-neutral ammoniacal sulphate solutions (1 M ammonia species) containing dissolved copper (added as CuS04.5H20). His results indicated that intergranular cracking occurred in pH ranges where Cu20 was thermodynamically stable (tarnishing solutions) and that transgranular cracking was more likely in pH regions where Cu20 was unstable (non-tarnishing solutions). This work stimulated many studies on the mechanism of SCC under film formation (tarnishing) conditions in solutions of similar composition to those of Mattsson,8p’6 and generated interest in the specific role of cupric ammine complexes in the cracking process.““43’7-2’ Following the observations of Mattsson2 Pugh and co-workers’7*‘832’ developed Cucontaining tarnishing and non-tarnishing solutions for studying intergranular and transgranular SCC of cl-brass, respectively, that were based on highly concentrated solutions of ammonia (15 M). For non-tarnishing conditions, the copper was added either as CU(NO~)~ (e.g. 8 g/L of dissolved Cu), or added as small amounts of copper powder (e.g. 1 g/L of Cu) that were dissolved under oxygenated conditions. Tarnishing solutions were prepared by dissolving larger amounts of copper powder (e.g. 68 g/L Cu) under oxygenated conditions. In all cases, the final form of the soluble copper was a Cu”ammine. These highly concentrated ammonia solutions were not subject to the ageing and precipitation phenomena reported in Mattsson-type solutions2 and have been used widely in recent years for fundamental studies of cracking mechanisms in brass. 15,22 28 More recently, significant attention has been directed towards the transgranular SCC of a-brass in non-tarnishing 15 M ammonia solutions containing Cu’-ammines.29p33 Kaufman and Fink,29 and Shahrabi et ~1.~’ prepared test solutions by adding both CU(NO~)~ and Cu powder to 15 M ammonia in a sealed test cell so that all Cu” was reduced by the Cu to Cu’. Sieradzki, Newman and co-workers30m32 ensured the absence of Cu” by dissolving Cu20 in 15 M ammonia in the presence of copper powder under a nitrogen atmosphere. Dickson et a1.33 also followed the preparation procedures of Newman et a1.32 The published SCC studies on a-brass in 15 M ammonia, irrespective of whether Cu” or Cu’-ammines were present, were conducted with an imperfect knowledge of activities and equilibria in highly concentrated ammonia solutions. The work of Bjerrum34 showed that the dominant soluble forms of copper should be the Cu”-pentammine (Cu(NH,),‘+) and Cu’-diammine (CU(NH~)~+) complexes in concentrated (high activity) ammonia solutions, as assumed by Pugh et al.‘7,‘8,2’ and others. ‘9,20 The E-pH diagram reported by Leidheiser and Kissinger35 for 15 M ammonia did not contain the Cu”-pentammine complex and was based erroneously on an assumed activity of 15 for the ammonia species. Jenkins and Durham36 attempted to estimate the activity coefficient of NH3, using molar concentrations, equilibrium constants and visual observation of conditions under which Cu20 tarnish films formed on copper. To date, there are no equilibrium diagrams for the Cu-NHs-H20 system that are relevant to the concentrated 15 M ammonia situation and which provide a thermodynamic basis for a better understanding of dissolution and oxide formation phenomena in these solutions.
Copper and zinc equilibria
1309
There is a similar absence of published equilibria for the Zn-NHs-H20 system in 15 M ammonia, despite a long held view that SCC of a-brass is strongly related to preferential dissolution (de-alloying) of zinc.‘2,‘4,3@32,35,37A3 Early E-pH studies in 1 M ammonia by Mattsson2 and Johnson and Leja3 were in disagreement regarding the domain of stability of Zn(OH)z, but these differences were resolved by Letowski and Niemic4 in terms of the dominance of NH: at pH < pki and the dominance of NH3 at pH > pki. Leidheiser and Kissinger35 presented an equilibrium diagram for zinc in 15 M ammonia solution that incorrectly implied molarity was equal to activity. Also, Bertocci ‘I9 discussed the effect of the molar concentration of ammonia on the solubility of zinc in terms of the Zn”-tetrammine complex (Zn(NH3)42+) and the Zn”-oxyanion (ZnOi-), but he did not construct any E-pH diagrams. The present work was concerned with the construction of equilibrium diagrams that are applicable to the Cu-NH3-H20 and Zn-NH3-H20 systems in concentrated 15 M ammonia (NH3 + NH4+) solutions of the type commonly used for studying SCC of a-brass in the presence of copper additions. Particular attention was given to the estimation of activities of species in accordance with the conventions of aqueous thermodynamics.
ACTIVITIES
OF SPECIES
IN CONCENTRATED
AMMONIA
SOLUTION
Activities of solutes were obtained from the product of their molal concentrations (m) and molal activity coefficients (y), using estimated or calculated y-values. The activity of the solvent (H20) was obtained from the variation in vapour pressure with the mole fraction of constants (k) that were used in the calculations were the solvent.44 Any equilibrium determined from the standard chemical free energy data edited by Bard et a1.,45 as listed in Table 1. The activity of a species is indicated by enclosing it in square brackets, e.g.
N--Id, W201. Occasionally the un-ionised aqueous form of ammonia is written in the literature as ammonium hydroxide, NH40H (=NHs.H,O), with u“ = -263.8 kJ mol-‘.46 Inspection of Table 1 shows that this number is simply the sum of the u” values of the NH3 and H20 species, indicating that NH40H is not uniquely different from NH+ Hence, the unionised form may be treated entirely as NHs, consistent with the practice followed by others,3-5’7 and is the form used in the subsequent activity calculations. Ammonia solution A 15 M ammonia solution contains 28.5% by weight of total ammonia species (NH3 + NH4+) measured as NH3,47 which converts to 23.4 m total ammonia and a mole fraction of 0.7 for water. Interpolation of partial pressure data shows that the partial pressure of NH3 gas over aqueous ammonia solutions at 25°C is 6.74 kPa and 85.7 kPa at 2.92 m and 23.4 m total ammonia, respectively.a In these solutions, the dissolved ammonia exists almost wholly as NH3. Hence, we may write [NH3123.4/[NH3i2.93 = b23.4 X 23.4m2.92
X 2.92)
= 8w6.74
(2)
where the subscripts refer to the molal concentrations of the solutions. Furthermore, if it is assumed that ~2.92 is close to unity, because this solution is less concentrated, then eqn (2) shows that y23.4 is - 1.6. Therefore,
N-M,,., -
1.6 x 23.4 - 37.4
(3)
1310
D. Tromans Table
1.
Standard
Species
chemical
State*
Hf
free energies/m01 u”, (kJ.mol-‘) 0.00 -237.18
aq
Hz0
1
cu
s
g:+
cuo2*
0.00
i65.7 + 50.3 - 258.90 - 183.90 - 148.10 - 129.70 -359.50 - 10.30 -63. + 15.6 - 30.50 -73.20 -111.50
aq aq
HCu02-
aq -
aq
cu20
S
cue
s
CWW2
S
Cu(NHs)+ WNH3)2
aq +
aq
Cu(NHs)*+ CU(NH~)*~ + Cu(NHs),‘+ CU(NH&~+
aq aq aq aq
Cu(NHs)s2 + NH3 NH‘,+ Zn Zn2 + ZnO Zn(OH)42Zn(NH)s* * Zn(NH)42t
aq aq aq s
- 134.70 - 26.60 -79.37 0.00 - 147.16 -318.32 - 877.40 -268.60 -307.
aq S
aq aq aq
*aq = aqueous;
(u”) at 25°C Reference 45 45 48
48 48 48 48
48 49 48 48 48 48
48 48 48
Calc. from 5’ 46 46 52 52 52 52 3 52
I = liquid.
s = solid;
indicating that the [NH,] activity in a 15 M ammonia solution is much higher than previously assumed.‘9*35,36 Interpolation of published data shows that the partial pressures of water above ammonia solutions at 25°C are 3.21 kPa and 2.17 kPa when the mole fractions of water are 1 and 0.7,@ respectively. Hence, recognising that [Hz01 is unity for unit mole fraction, the activity at 0.7 mole fraction [H~0]a.,, is given by [HzO],,,
= 2.17/3.21
- 0.68
indicating that the activity of water in 15 M ammonia is slightly less than its mole fraction. The activities of [NHd+] and [OH-] may be obtained from the aqueous equilibrium reaction NH3 + Hz0
= NH;
+ OH-,
k5 = 1.78 x 1O-5
(5)
If there is a total of N moles of ammonia species, then N( 1 -a) moles of NH3 will be in equilibrium with Ncl moles of [NH4+] and No! moles of OH-. Consequently, in a solution of N molality we may write ks = [NH4+][OH-]
WWW201
=
]N%+t$[NWon-]
[NC1-
@YNH~I[HZOI
(6)
Copper
and zinc equilibria
1311
In a 23.4 m ammonia solution, Nis 23.4, and from eqns (3) and (4), yNHl is 1.6 and [Hz01 is 0.68. The value of iVclis sufficiently small (i.e. small concentration of species), as will be evident below, that the values of You- and YNH+may be approximated to unity. Hence, after substituting 1.78 x 1O-5 for kS and solving gqn (6), ct is 9.0932 x 10F4, and [OH-] and [NH:] are both 2.128 x 1OK’. The [H+] activity in 23.4 m (15 M) ammonia may be calculated from the water equilibrium Hz0 = H+ + OH-, k7 = 9.945 x IO-l5 Substituting for [Hz01 =0.68 and [OH-] =2.128 x lo-* 3.178 x lo-l3 that is equivalent to a pH of 12.5 (pH= -log[H+]).
(7) yields a [H+] value of
Addition of copper powder
The dissolution of copper powder under fully oxygenated conditions in concentrated ammonia solutions leads to the formation of Cu”-ammines and OH- ions (pH changes) in proportions of 1:2, as shown in eqn (8), Cu + nNH3 + H20 + 1/202 + Cu(NH&+
+ 20H-
(8)
where n is the number of NH3 ligands on the ammine complex. When 1 g/L of copper is dissolved in 15 M ammonia it is equivalent to a concentration of 2.4606 x lo-* m of soluble copper. This concentration is sufficiently small that its corresponding y may be approximated to unity so that Cu(NHs)z+ is 2.4606 x lo-*. Equation (8) shows that 2.4606 x lo-* moles of soluble copper generate 4.9212 x lo-* moles of OH-. A portion of the OH-, x moles, will interact with the pre-existing 2.128 x lo-* m of NH,f via eqn (5) and the remainder will add to the pre-existing 2.128 x lo-* m of OH-. Hence, the new equilibrium with respect to ammonia becomes k = [2.128 x lo-* - x][2.128 x lo-* + 4.9212 x lo-* - X] = 1 78 x 1o_5 5 [37.4 - 2.128 x lo-* + x][HzO]
(9)
where [NH31 is [37.4 - 2.128x10-*+x], and [Hz01 is 0.68 according to eqn (4). Solving eqn (9) gives x = 1.336 x lo-* and [OH-] = 5.713 x lo-*. Therefore, from eqn (7) [H+], = 1.184 x lo-l3 and the pH is 12.93. Similarly, the dissolution of 6 g/L of copper is equivalent to 1.4764 x 10-l m of soluble copper, which corresponds to 1.4764 x 10-l [Cu(NHs)F] after approximating y to unity. Dissolution is accompanied by the generation of 2.9528 x 10-l moles of OH-, via eqn (8) of which a portion, x moles, interacts with the pre-existing NH: and the remainder adds to the pre-existing OH- so that eqn (9) is now changed to k = [2.128 x lo-* - x][2.128 x 10-l + 2.9528 x lo-* -xl 5 [37.4 - 2.128 x lo-* + x][HzO]
= 1 78 x 1o_5
(10)
Solving eqn (10) gives x = 1.976 x lo-’ and [OH-] = 2.968 x 10-l. Hence, from eqn (7) [Hf] = 2.278 x 10-14, and the pH is 13.642. Similar calculations for the dissolution of 10 g/L of copper show that Cu(NHs)i+ is 2.4606x lo-‘, [OH-] is 4.931 x lo-‘, [H+] is 1.372x 10Vi4, and the pH is 13.86. The principal results of all the activity calculations are summarised in Table 2.
D. Tromans
1312 Table 2.
Activities
in 15 M ammonia
Solution
[NH,1
15 M l5M I5 M 15 M
37.40 37.40 37.40 37.40
ammonia +lg/LCu +6g/LCu + 10 g/L Cu
solutions
E-PH
containing
copper powder
dissolved
as Cu”
[Hz01
[Cl?‘]
PH
0.68 0.68 0.68 0.68
2.4606 x IO-* 1.4764x 10-l 2.4606 x 10-l
12.50 12.93 13.64 13.86
EQUILIBRIA
All E-pH equilibria in the Cu-NH3-Hz0 and Zn-NHs-Hz0 systems in 15 M ammonia were based on [NH,] and [H,O] activities in Table 2, together with the standard chemical free energies per mole of species (u”) listed in Table 1. Most of the u” data for the copper species were obtained from the compilation by Bertocci and Wagman, with the notable exceptions of CuO and Cu(NHs):+. The appropriate u” for CuO is uncertain.48 We have used the value of - 129.7 kJ mol-’ listed by Barner and Scheurman,49 which is lower than that reported by King et a1.” (- 127.9 kJ mol-‘) and higher than the - 134 kJ mol-’ suggested by Bertocci and Wagman. The u” for Cu(NH3):’ was calculated from the equilibrium constant (log kl, = - 0.6) reported by Sillen and Martel15’ for eqn (11) at 25°C using the u” values for NH3 and Cu(NH3):’ listed in Table 2, CUE+
4
+ NH3 = Cu(NH):+,
k,, = 0.251
(11)
Thermodynamic data for the zinc system were obtained principally from the compilation by Brodd and Werth.52 When consulting Ref. 52, the reader should note that there are some typographical errors relating to the zincate ion, Zn(OH$, e.g. p” is listed as + 877.4 instead of - 877.4 kJ mol- ’ and the text states that the chemical free energy is - 209.7 kJ mol-- ’ instead of kcal mol- ‘. The resulting E-pH diagrams were constructed for the pH range from 12 to 15, because of the high pH encountered in 15 M ammonia solutions (Table 2). Solid lines de-lineate regions of stability of solid phases. Coarse broken lines show equilibria between dissolved (aqueous) species. Fine broken lines, labelled (a) and (b) correspond to the hydrogen and oxygen electrodes at one atmosphere fugacity and [Hz01 = 0.68. Cu-NHrH20 system Equilibrium E-pH diagrams for the Cu-NH3-H20 system are shown in Figs l-3 for 1, 6, and 10 g/L of dissolved Cu species, respectively. The pH of the corresponding Cu”containing solution in Table 2 is indicated on each figure by an arrow. The dominant soluble copper species in equilibrium with the oxides are the Cu”-pentammine and the Cu’diammine complexes. The oxides, CuO and CuzO, are stable in the higher pH regions and their stability range extends to lower pH values as the dissolved copper activity increases. Some of the important reaction equilibria are shown below in eqns (12)-( 18), Cu(NHs):+ Cu(NHs)t
+ e- = Cu(NHs)l
+ 3NHj,Ey,
= +0.0839 V(SHE)
+ e- = Cu + 2NH3, Ey3 = -0.1015
Cu20 + 4NH3 + 2H+ = 2Cu(NH3)2f
V(SHE)
+ H20, k,4 = 1.1202 x lOI
(12) (13) (14)
Copper
and zinc equilibria
13
1313
14
15
PH Fig. 1, Cu-NHsPH20 equilibria of dissolved Cu species (activity
at 25°C in 15 M ammonia = 2.4606 x 10m2). Arrow Table 2.
CuO + 5NH3 + 2H+ = Cu(NH&+
solution ([NHs] = 37.4) containing 1 g/L shows pH of corresponding solution in
+ H20, kls = 1.3707 x 10”
2Cu(NH3):+ + Hz0 + 2e- = Cu20 + lONH3 + 2H+, Eye = -0.4791 V(SHE)
(15) (16)
CuzO + 2H+ + 2e- = 2Cu + H20, Ey, = +0.4614 V(SHE)
(17)
2CuO + 2H+ + 2e- = CuzO + HzO, E;, = +0.6521 V(SHE)
(18)
------.-
12
%__@_
13
___ 14
15
PH Fig. 2. Cu-NHs-Hz0 equilibria of dissolved Cu species (activity
at 25°C in 15 M ammonia solution ([NH31 = 37.4) containing 6 g/L = I .4764 x lo-‘). Arrow shows pH of corresponding solution in Table 2.
D. Tromans
1314
13
15
14
PH Fig. 3. Cu-NH3-H20 equilibria at 25°C in 15 M ammonia solution ([NH,] = 37.4) containing 10 g/ I_ of dissolved Cu species (activity = 2.4606 x lo-‘). Arrow shows pH of corresponding solution in
Table 2.
where the standard hydrogen electrode
electrode potentials, (SHE) at 25°C.
E”, are referenced
with respect
to the standard
Zn-NH,H20 system Aqueous equilibria in the Zn-NHs-Hz0 system are shown in Fig. 4 for soluble zinc activities of lop6 and lo- ‘. No oxide phase is stable in the pH range of interest and zinc dissolves at much lower (more active) potentials than copper. The dominant soluble zinc species are the Zn”-tetrammine, Zn(NHs), 2f , below pH 13.656, and the Zn’t-zincate, Zn(OH)i-, above pH 13.656. Thus, both species should be formed upon dissolution of zinc
0.5 f-------------a
..+.____._____
___ 1
O-
t!
:
-0.5
j
?-----.-----.*A
9 ui
Zn(NH$
Zn(OH)f
-
._.._.._._._._.._.._ ___L
-I-1 -
b
-1.5
-
10.6
Zn -2
12
i
I
I
13
14
15
PH Fig. 4, Z~_NH~-H~O equilibria at 25°C in 15 M ammonia solution t[NH~I=37.4). Diagram constructed for 1x 10e6and 1 x 10-l activities of dissolved Zinc species.
Copper and zinc equilibria
1315
in the Cu-containing solutions in Table 2, depending on the solution pH. The Cu-containing solutions are of interest because stress corrosion studies of brass in these solutions will lead to dissolution of zinc. Important reaction equilibria are shown below in eqns (19)-(21), Zn(NH3)p
+ 4H20
Zn(NH# Zn(OH$
= Zn(OH)i-
+ 4NHs + 4H+, kis = 2.1814 x 1O-48
+ 2e- = Zn + 4NH3, Ei, = -1.0414 + 4H+ + 2e- = Zn + 4H20,
V(SHE)
Ei, = +0.3694 V(SHE)
(19) (20) (21)
DISCUSSION Copper equilibria The E-pH diagrams in Figs 1-3 are in remarkably good agreement with the reported behaviour of brass in Cu-containing 15 M ammonia solutions,‘5X’7,‘8,2’-28 indicating that they closely describe the stability of the copper species. For example, Fig. 1 shows that when Cu is added as 1 g/L of copper powder to form Cu(NH3):’ at a pH of 12.93 (Table 2) the only possible reduced forms of the pentammine at this pH are Cu(NHs)p or Cu (i.e. a nontarnishing condition), consistent with observations. Complete reduction of the pentammine to Cu would be consistent with the reported formation of brown films on brass. l5 Figures 2 and 3 show that when Cu(NH# species are formed by the addition of 6 and 10 g/L of copper powder, respectively, the oxide stability regions enlarge and the resulting pH of the solutions, (13.64 and 13.86) places the solutions in situations that now allows the pentammine to be reduced to CUZO (tarnishing condition), consistent with observations.“,“,‘s,*‘-2s Also, Fig. 3 shows that 10 g/L of copper powder may be dissolved in 15 M ammonia under oxygenated conditions without forming CuO, which was the upper limit of powder dissolution investigated by Pugh et a1.‘7,‘8 Furthermore, if Cu” is added as Cu(NO&, in amounts up to the equivalent of 10 g/L of copper, the pH of the solution should remain near 12.5, corresponding to that of 15 M ammonia alone (Table 2) because no hydroxyl ions are generated by reactions shown in eqn (8). Figure 3 shows it is not possible to reduce Cu(NHs):+ to Cu20 at pH 12.5. Therefore, the ammoniacal Cu(NO3)* solutions should be non-tarnishing, as observed by Pugh er a1.17+*8*21 Similarly, 15 M ammonia with additions of Cu(NO& followed by additions of Cu powder to reduce the Cu”-pentammine to Cu(NH3); should be non-tarnishing, as observed.29,30 The E-pH equilibria were consistent with the preparation ofnon-tarnishing 0.05 M Cu’ammine solutions by the dissolution of 0.025 M CuzO powder (= 3.18 g/L Cu) in 15 M ammonia.3’ The dissolution reaction is shown in eqn (22), Cu20 + 4NHs + H20 = 2Cu(NH$
+ 20H-
(22)
where equal moles of OHand Cu(NH3): are formed. In 15 M ammonia, 0.05 M Cu(NH3): is equivalent to 7.8125 x lo-* m, leading to formation of 7.8125 x lo-* m OHof which a portion (x) interacts with the pre-existing NH: and the remainder contributes to a change in pH, as described in (9) and (IO), producing a final pH of 13.13. Inspection of Fig. 3 shows that even when [Cu(NH3):] is increased to 2.4606 x 10-l (E 10 g/L Cu) no Cu20 is formed until the pH reaches 13.363, confirming that additions of 0.025 M Cu20 are readily soluble in 15 M ammonia, in good agreement with published data.3’
D. Tromans
1316
It is evident from Figs l-3 and Table 2 that two important consequences of the oxygenated dissolution of copper powder are the movement of the stability lines on the EpH diagrams and the increase in pH, both of which dictate whether the resulting solution produces tarnishing or non-tarnishing conditions. It was difficult to confirm precisely the calculated pH values listed in Table 2, because the usual glass pH-electrode behaves unreliably above pH 12. However, the glass electrode confirmed the magnitude and direction of the predicted pH changes, lending confidence to the calculated values. For example, an increase of approximately 1.2-l .3 pH units was measured after the dissolution of 6 g/L of copper powder, as compared to the calculated increase of 1.14 pH units.
Zinc equilibria Figure 4 shows that the domain of stability of zinc occurs at much lower (more active) potentials than the stability region of copper in Figs 1-3, consistent with the required active behaviour of zinc in de-alloying-related mechanisms of SCC of brass in 15 M ammonia solutions. 3om32735Y4’113 It is evident that the dominant soluble forms of zinc in these solutions may be either the Zn”-tetrammine or the Zn”-zincate complexes, depending upon whether the pH is below or above 13.656, respectively. The active behaviour of zinc, and effect of pH, was confirmed by conventional anodic polarisation tests at 24°C on metallographically polished pure metal (99.999% by weight) in 15 M ammonia (pH 12.5) and in 15 M ammonia with sufficient NaOH added to increase the pH to 14. The tests were conducted in a single compartment glass cell containing a Pt counter electrode. The working electrode surface was 90 mm2 and the potential was measured with respect to an external HgO/Hg reference electrode in 2.5 m NaOH (pH 14.6),53 where EHgopg= +0.063 V(SHE). The reference electrode was interfaced to the working solution via a salt bridge and polyethylene Luggin probe assembly containing aqueous NaOH. The probe was end-fitted with a sintered zirconia plug. Concentrated ammonia solutions have poor electrical conductivity,35.47 and it was necessary to apply corrections for the ohmic potential drop between the tip of the Luggin probe and the working electrode surface. The ohmic correction was obtained from a linear plot of the potential vs current density (i). The anodic polarisation curves are shown in uncorrected and ohmic-corrected forms in Fig. 5, where log (i) is plotted against E. Uncorrected data were collected at potentials up to + 1.O V(HgO/Hg). All the corrected data are presented in Fig. 5, whereas uncorrected data above -0.5 V(HgO/Hg) are not shown. Figure 5 shows very low open circuit potentials (E,,& consistent with the predictions of (- 1.321 V(SHE)) and - 1.409 V(HgO/Hg) Fig. 4. The E,, were - 1.384 V(HgO/Hg) (- 1.346 V(SHE)) in the 15M NH3 and 15 M NH3 + NaOH solutions, respectively, where the slightly lower potential in the NaOH containing solution was consistent with the more active behaviour associated with the dominance of Zn(OH$ complexes at pH 14 (see Fig. 4). Furthermore, the ohmic-corrected curve in the NaOH-containing solution exhibited anodic Tafel behaviour with a slope of 41.3 mV at i< 32 A m-‘. This was close to the 40 mV reported by Bockris et aZ.54 for anodic behaviour of zinc in alkaline KOH solutions, where the rate determining anodic step was shown to be the transfer of the second electron according to eqn (23), Zn(OH>; The steeply rising section
+ OH-
+ Zn(OH),
of the ohmic-corrected
+ e-
polarisation
(23) curve at higher
i in the
Copper
H
-3
and zinc equilibria
Comcted
-2
-1 log(i),
Fig. 5.
Anodic
1317
0
1
2
3
(A.m -2)
polarisation behaviour of zinc in 15 M NH3 and 15 M NH3 + NaOH uncorrected and ohmic-corrected data. Scan rate 1 mVs- ’
showing
NaOH-containing solution indicated the possibility of mass transfer controlled anodic behaviour, possibly the transfer of OH- ions to the electrode surface. In contrast, the ohmic-corrected polarisation curve in 15 M ammonia did not exhibit a Tafel region and the anodic currents were considerably lower than when NaOH was present. The shape of the polarisation curve suggested that anodic behaviour was falling under mass transfer control much earlier (i.e. at lower current densities) than in the NaOH-containing solution, consistent with the lower concentration of OH- ions at pH 12.5. This implied that anodic dissolution was still determined by the rate of formation of zincate species, even in a pH regime where Zn(NH# complexes were the dominant soluble species and there was a high concentration of NHs available for their formation. If this interpretation is correct, transient zincate species may be formed initially at the electrode surface, which then interact with NH3 in solution to form the stable tetrammine complex via eqn (24) Zn(OH)i-
+ 4NHs + Zn(NHs)F
+ 40H-
(24)
CONCLUSIONS (1) The calculated [NHs] and [Hz01 activities in 15 M ammonia solutions are 37.4 and 0.68, respectively. (2) Based on the above activities, new E-pH diagrams were constructed for the CuNHs-Hz0 and Zn-NHs-Hz0 systems in 15 M ammonia solutions in the pH range 12-15. These diagrams are consistent with the reported behaviour of cl-brass in these solutions in relation to tarnish (CuzO) formation, non-tarnishing conditions, and the active behaviour of the zinc component. (3) The new equilibrium diagrams are particularly useful for understanding tarnishing and non-tarnishing phenomena in copper-containing 15 M ammonia solutions of the type that are used frequently for studying SCC of a-brass. (4) Anodic dissolution of zinc in 15 M ammonia to form the Zn”-tetrammine complex may involve the formation of transient zincate species.
1318
D. Tromans
Acknowledgements-The author wishes to thank the Natural Canada for financial support. Also, the experimental assistance
Sciences and Engineering of Mr C. Fong is gratefully
Research Council acknowledged.
of
REFERENCES 1. J. 2. 3. 4. 5. 6. 7. 8. 9. IO. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33.
34. 35. 36. 37. 38. 39. 40. 41.
Halpern, Trans. AIME 9,280-289 (1957). E. Mattsson, Electrochim. Acta 3, 279-291 (1961). H.E. Johnson and J. Leja, J. Electrochem. Sot. 112, 6388641 (1965). F. Letowski and J. Niemic, J. Electrochem. Sot. 113, 629-630 (1966). R. Bartonicek and M. Lukasovski, Corros. Sci. 9, 3542 (1969). R. Bartonicek and M. Lukasovski, Corros. Sci. 11, 111-l 14 (1971). T.P. Hoar and G.P. Rothwell, Electrochim. Acto 15, 1037-1045 (1970). A.J. Forty and P. Humble, Phil. Mag. 8, 247-264 (1963). T.P. Hoar and C.J.L. Booker, Corros. Sci. 5, 821-840 (1965). A.J. McEvily and A.P. Bond, J. Electrochem. Sot. 112, 131-138 (1965). H.E. Johnson and J. Leja, Corrosion 22, 178-189 (1966). A.J. Forty and P. Humble, Environment-Sensitve Mechanical Behavior (Eds. A.R.C. Westwood and N.S. Stoloff), pp. 403419, Gordon and Breach, New York (1966). S.S. Birley and D. Tromans, Corrosion 27, 297-311 (1971). L.P. Lee and D.Tromans, Environmenr-Sensitive Fracture of Engineering Materials (Ed. Z.A. Foroulis), pp. 232-249, TMS-AIME (1978). M. Kermani and J.C. Scully, Corros. Sci. 19, 89-l 10 (1979). M. Kermani and J.C. Scully, Corros. Sci. 19, Ill-122 (1979). E.N. Pugh and A.R.C. Westwood, Phil. Mag. 13, 167-183 (1966). E.N. Pugh, J.V. Craig and A.J. Sedriks, Fundamental Aspects of Stress Corrosion Cracking (Eds. R.W. Staehle, A.J. Forty and D. van Rooyen), pp. 118-151, NACE, Houston TX (1969). U. Bertocci, Elecrrochim. Metoll. 3, 275-280 (1968). A.W. Blackwood and N.S. Stoloff, Trans. ASM 62, 677-689 (1969). E.N. Pugh and J.A.S. Green, Metall. Trans. 2, 3129-3134 (1971). T.R. Pinchback, S.P. Clough and L.A. Heldt, Mefall. Trans. A 6A, 1479-1483 (1975). M. Kermani and J.C. Scully, Corros. Sci. 19, 489-506 (1979). J.A. Beavers and E.N. Pugh, Metall. Trans. A llA, 809-820 (1980). D.V. Beggs, M.T. Hahn and E.N. Pugh, Hydrogen Embrittlement and Stress Corrosion Cracking, pp. 181, ASM, Metals Park OH (1984). S.P. Lynch, Scriplo Metall. 18, 321-326 (1984). E.I. Meletis and R.F. Hochman, Corros. Sci. 24, 843-862 (1984). A. Parthasarathi and N.W. Polan, Corrosion 43, 747-755 (1987). M.J. Kaufman and J.L. Fink, Metali. Trans. A 18A, 153991542 (1987). T. Shahrabi, R.C. Newman and K. Sieradzki, /. Electrochem. Sot. 140, 348-352 (1993). K. Sieradzki, J.S. Kim, A.T. Cole and R.C. Newman, J. Electrochem. Sot. 134, 1635-1639 (1987). R.C. Newman, T. Shahrabi and K. Sieradzki, Scripta Metall. 23, 71-74 (1989). J.I. Dickson, S.Q. Li, J.-P. Bailon and D. Tromans, Parkins Symposium on Fundamental Aspects of Stress Corrosion Cracking (Eds. SM. Bruemmer, E.I. Meletis, R.H. Jones, W.W. Gerberich, F.P. Ford, and R.W. Staehle), pp. 3033322, TMS, The Minerals, Metals and Materials Society (1992). J. Bjerrum, Metal Ammine Formation in Aqueous Solution - Theory of the Reversible Step Reaclions, P. Haase and Son, Copenhagen (1957). H. Leidheiser and R. Kissinger, Corrosion 28, 218-221 (1972). L.H. Jenkins and R.B. Durham, J. Electrochem. Sot. 117, 768-774 (1970). C. Edeleanu, Physical Metallurgy of Stress Corrosion Fracture (Ed. T.N. Rhodin), pp. 79-96, Interscience, New York (1959). A.J. Forty, Physical Metallurgy of Stress Corrosion Fracture (Ed. T.N. Rhodin), pp. 99, Interscience, New York (1959). P. Hari, Corros. Sci. 15. 123-144 (1975). N.W. Polan, J.M. Popplewell and M.J. Pryor, J. Electrochem. Sot. 126, 129991300 (1979). A. Parthasarathi and N.W. Polan, Metall. Trans. A 13A, 2027-2033 (1982).
Copper
and zinc equilibria
1319
42. U. Bertocci, F.I. Thomas and E.N. Pugh, Corrosion 40, 439440 (1984). 43. K. Sieradzki and R.C. Newman, Phil. Msg. 51, 95-132 (1985). 44. (a) Chemical Engineers’ Handbook (Ed. J.H. Perry), pp. 404407, McGraw-Hill Book Co. Inc., New York (1941), (b) Perry’s Chemical Engineers Handbook, 6th Edition (Eds., R.H. Perry, D.W. Green and J.O. Maloney), pp. 3-71-3-73, McGraw-Hill, Inc., New York (1984). 45. Standard Potentials in Aqueous Solutions (Eds. A.J. Bard, R. Parsons and J. Jordan), IUPAC, Marcel Dekker Inc., New York (1985). 46. J.T. Maloy, Standard Potentials in Aqueous Solutions (Eds. A.J. Bard, R. Parsons and J. Jordan), pp. 127139, IUPAC, Marcel Dekker Inc., New York (1985). 47. CRC Handbook of Chemistry and Physics (Ed. R.C. Weast), pp. D-221, CRC Press, Inc., West Palm Beach, FL (1978). 48. U. Bertocci and D.D. Wagman, Standard Potentials in Aqueous Solutions (Eds. A.J. Bard, R. Parsons and J. Jordan), pp. 287-293, IUPAC, Marcel Dekker Inc., New York (1985). 49. H.E. Barner and R.V. Scheurman, Handbook of Thermochemical Data for Compounds and Aqueous Species, pp. 45, J. Wiley and Sons, Inc., New York (1978). 50. E.G. King, A.D. Mah and L.B. Pankratz, Thermodynamic Properties of Copper and Its Inorganic Compounds, pp. 13, INCRA Monograph II, The Metallurgy of Copper, International Copper Research Association, Inc. (1973). 51. L.C. Sillen and A.E. Martell, Stability Constants of Metal-Zon Complexes, pp. 152, Special Publication No. 17, The Chemical Society, London (1964). 52. R.J. Brodd and J. Werth, Standard Potentials in Aqueous Solutions (Eds. A.J. Bard, R. Parsons and J. Jordan), pp. 249-257, IUPAC, Marcel Dekker Inc., New York (1985). 53. D.D. Macdonald and M.C.H. McKubre, Temperature Limitations of Alkaline Battery Electrodes, Part ZZ, pp. 20-23, SRI Project No. 7775, SRI International, Menlo Park, CA (1979). 54. J.O’M. Bockris, 2. Nagy and A. Damjanovic, J. Electrochem. Sot. 119, 285-295 (1972).
Pergamon
Printed in Great Britain. All rights reserved 0010-938X/97 $17.00+0.00
PII: S001&938X(97)00030-9
COPPER AND ZINC EQUILIBRIA IN CONCENTRATED AMMONIA SOLUTIONS: RELEVANCE TO STRESS CORROSION CRACKING OF ALPHA-BRASS D. TROMANS University
of British Columbia,
Department of Metals and Materials Engineering, British Columbia, V6T 124, Canada
6350. Stores Road, Vancouver,
Abstract-Potential-pH diagrams for the Cu-NHx-Hz0 and Zn-NH,-Hz0 systems have been constructed for 15 M ammonia solutions commonly used for studying SCC of a-brass. Equilibria and pH calculations were based on activities of NH3 and water that were 37.4 and 0.68, respectively. The equilibrium diagrams are in agreement with the reported behaviour of brass in copper-containing 15 M ammonia, and with the active behaviour of zinc, They are the most useful diagrams to date for studies relating to SCC of a-brass in tarnishing and non-tarnishing 15 M ammonia solutions. Anodic polarisation tests indicated that formation of the zinc-tetrammine complex in 15 M ammonia may involve transient zincate ions. 0 1997 Elsevier Science Ltd Keywords:
A. brass, A. copper,
A. zinc, B. polarisation.
INTRODUCTION The stress corrosion cracking (KC) of a-phase Cu-Zn alloy (brass) in ammoniacal environments has received considerable attention, partly because the system is useful for developing and testing mechanistic models of environment-assisted crack propagation. In contrast, the relevant potential (E) and pH equilibria of copper and zinc have received less consideration, particularly in highly concentrated ammonia solutions. Some of the earliest E-pH equilibrium diagrams for the Cu-NH3-Hz0 system were constructed by Halpern,’ Mattsson,’ and Johnson and Leja3 for 1 M of total dissolved ammonia species (NH3 + NHd+), where the molarity was equated with activity. However, when considering the formation of Cur-ammine and Cu”-ammine complexes, and the stability of the copper oxides, these workers did not completely recognise the consequences of the equilibrium between the ionised and un-ionised forms of dissolved ammonia, NH;
= NH3 + H+, kl = 5.628 x lo-“,
pk, = 9.25
(1)
where k, is the equilibrium constant for eqn (1) at 25°C and pk, = -log k,. Consequently, below pH 9.25 aqueous equilibria between dissolved ammonia and copper species should be determined with respect to NH4+ whereas above pH 9.25 the equilibria should be determined with respect to NH3. Halpern,’ and Johnson and Leja3 determined copper oxide stability relative to NH 3, whereas Mattsson’ considered only NH4+. Thus, their E-pH diagramsIp were only partially correct. Letowski and Niemic,4
Manuscript
received 20 June 1996; in amended
form 14 February 1307
1997.
1308
D. Tromans
and Bartonicek and Lukasovska’ considered both forms of ammonia and showed that the solublities of the copper oxides reached a maximum when the pH was equal to pk’ ,4 leading to oxide stability domains in pH regions above and below 9.25.5,6 Hoar and Rothwell’ also considered both forms of ammonia and constructed E-pH diagrams for copper in 1 M total ammonia (assumed to be unit activity) that displayed two regions of oxide stability. Although Mattsson’s2 E-pH diagram was only partially correct, his work drew attention to the influence of cupric-ammine complexes and Cu20 formation on SCC of c1brass in near-neutral ammoniacal sulphate solutions (1 M ammonia species) containing dissolved copper (added as CuS04.5H20). His results indicated that intergranular cracking occurred in pH ranges where Cu20 was thermodynamically stable (tarnishing solutions) and that transgranular cracking was more likely in pH regions where Cu20 was unstable (non-tarnishing solutions). This work stimulated many studies on the mechanism of SCC under film formation (tarnishing) conditions in solutions of similar composition to those of Mattsson,8p’6 and generated interest in the specific role of cupric ammine complexes in the cracking process.““43’7-2’ Following the observations of Mattsson2 Pugh and co-workers’7*‘832’ developed Cucontaining tarnishing and non-tarnishing solutions for studying intergranular and transgranular SCC of cl-brass, respectively, that were based on highly concentrated solutions of ammonia (15 M). For non-tarnishing conditions, the copper was added either as CU(NO~)~ (e.g. 8 g/L of dissolved Cu), or added as small amounts of copper powder (e.g. 1 g/L of Cu) that were dissolved under oxygenated conditions. Tarnishing solutions were prepared by dissolving larger amounts of copper powder (e.g. 68 g/L Cu) under oxygenated conditions. In all cases, the final form of the soluble copper was a Cu”ammine. These highly concentrated ammonia solutions were not subject to the ageing and precipitation phenomena reported in Mattsson-type solutions2 and have been used widely in recent years for fundamental studies of cracking mechanisms in brass. 15,22 28 More recently, significant attention has been directed towards the transgranular SCC of a-brass in non-tarnishing 15 M ammonia solutions containing Cu’-ammines.29p33 Kaufman and Fink,29 and Shahrabi et ~1.~’ prepared test solutions by adding both CU(NO~)~ and Cu powder to 15 M ammonia in a sealed test cell so that all Cu” was reduced by the Cu to Cu’. Sieradzki, Newman and co-workers30m32 ensured the absence of Cu” by dissolving Cu20 in 15 M ammonia in the presence of copper powder under a nitrogen atmosphere. Dickson et a1.33 also followed the preparation procedures of Newman et a1.32 The published SCC studies on a-brass in 15 M ammonia, irrespective of whether Cu” or Cu’-ammines were present, were conducted with an imperfect knowledge of activities and equilibria in highly concentrated ammonia solutions. The work of Bjerrum34 showed that the dominant soluble forms of copper should be the Cu”-pentammine (Cu(NH,),‘+) and Cu’-diammine (CU(NH~)~+) complexes in concentrated (high activity) ammonia solutions, as assumed by Pugh et al.‘7,‘8,2’ and others. ‘9,20 The E-pH diagram reported by Leidheiser and Kissinger35 for 15 M ammonia did not contain the Cu”-pentammine complex and was based erroneously on an assumed activity of 15 for the ammonia species. Jenkins and Durham36 attempted to estimate the activity coefficient of NH3, using molar concentrations, equilibrium constants and visual observation of conditions under which Cu20 tarnish films formed on copper. To date, there are no equilibrium diagrams for the Cu-NHs-H20 system that are relevant to the concentrated 15 M ammonia situation and which provide a thermodynamic basis for a better understanding of dissolution and oxide formation phenomena in these solutions.
Copper and zinc equilibria
1309
There is a similar absence of published equilibria for the Zn-NHs-H20 system in 15 M ammonia, despite a long held view that SCC of a-brass is strongly related to preferential dissolution (de-alloying) of zinc.‘2,‘4,3@32,35,37A3 Early E-pH studies in 1 M ammonia by Mattsson2 and Johnson and Leja3 were in disagreement regarding the domain of stability of Zn(OH)z, but these differences were resolved by Letowski and Niemic4 in terms of the dominance of NH: at pH < pki and the dominance of NH3 at pH > pki. Leidheiser and Kissinger35 presented an equilibrium diagram for zinc in 15 M ammonia solution that incorrectly implied molarity was equal to activity. Also, Bertocci ‘I9 discussed the effect of the molar concentration of ammonia on the solubility of zinc in terms of the Zn”-tetrammine complex (Zn(NH3)42+) and the Zn”-oxyanion (ZnOi-), but he did not construct any E-pH diagrams. The present work was concerned with the construction of equilibrium diagrams that are applicable to the Cu-NH3-H20 and Zn-NH3-H20 systems in concentrated 15 M ammonia (NH3 + NH4+) solutions of the type commonly used for studying SCC of a-brass in the presence of copper additions. Particular attention was given to the estimation of activities of species in accordance with the conventions of aqueous thermodynamics.
ACTIVITIES
OF SPECIES
IN CONCENTRATED
AMMONIA
SOLUTION
Activities of solutes were obtained from the product of their molal concentrations (m) and molal activity coefficients (y), using estimated or calculated y-values. The activity of the solvent (H20) was obtained from the variation in vapour pressure with the mole fraction of constants (k) that were used in the calculations were the solvent.44 Any equilibrium determined from the standard chemical free energy data edited by Bard et a1.,45 as listed in Table 1. The activity of a species is indicated by enclosing it in square brackets, e.g.
N--Id, W201. Occasionally the un-ionised aqueous form of ammonia is written in the literature as ammonium hydroxide, NH40H (=NHs.H,O), with u“ = -263.8 kJ mol-‘.46 Inspection of Table 1 shows that this number is simply the sum of the u” values of the NH3 and H20 species, indicating that NH40H is not uniquely different from NH+ Hence, the unionised form may be treated entirely as NHs, consistent with the practice followed by others,3-5’7 and is the form used in the subsequent activity calculations. Ammonia solution A 15 M ammonia solution contains 28.5% by weight of total ammonia species (NH3 + NH4+) measured as NH3,47 which converts to 23.4 m total ammonia and a mole fraction of 0.7 for water. Interpolation of partial pressure data shows that the partial pressure of NH3 gas over aqueous ammonia solutions at 25°C is 6.74 kPa and 85.7 kPa at 2.92 m and 23.4 m total ammonia, respectively.a In these solutions, the dissolved ammonia exists almost wholly as NH3. Hence, we may write [NH3123.4/[NH3i2.93 = b23.4 X 23.4m2.92
X 2.92)
= 8w6.74
(2)
where the subscripts refer to the molal concentrations of the solutions. Furthermore, if it is assumed that ~2.92 is close to unity, because this solution is less concentrated, then eqn (2) shows that y23.4 is - 1.6. Therefore,
N-M,,., -
1.6 x 23.4 - 37.4
(3)
1310
D. Tromans Table
1.
Standard
Species
chemical
State*
Hf
free energies/m01 u”, (kJ.mol-‘) 0.00 -237.18
aq
Hz0
1
cu
s
g:+
cuo2*
0.00
i65.7 + 50.3 - 258.90 - 183.90 - 148.10 - 129.70 -359.50 - 10.30 -63. + 15.6 - 30.50 -73.20 -111.50
aq aq
HCu02-
aq -
aq
cu20
S
cue
s
CWW2
S
Cu(NHs)+ WNH3)2
aq +
aq
Cu(NHs)*+ CU(NH~)*~ + Cu(NHs),‘+ CU(NH&~+
aq aq aq aq
Cu(NHs)s2 + NH3 NH‘,+ Zn Zn2 + ZnO Zn(OH)42Zn(NH)s* * Zn(NH)42t
aq aq aq s
- 134.70 - 26.60 -79.37 0.00 - 147.16 -318.32 - 877.40 -268.60 -307.
aq S
aq aq aq
*aq = aqueous;
(u”) at 25°C Reference 45 45 48
48 48 48 48
48 49 48 48 48 48
48 48 48
Calc. from 5’ 46 46 52 52 52 52 3 52
I = liquid.
s = solid;
indicating that the [NH,] activity in a 15 M ammonia solution is much higher than previously assumed.‘9*35,36 Interpolation of published data shows that the partial pressures of water above ammonia solutions at 25°C are 3.21 kPa and 2.17 kPa when the mole fractions of water are 1 and 0.7,@ respectively. Hence, recognising that [Hz01 is unity for unit mole fraction, the activity at 0.7 mole fraction [H~0]a.,, is given by [HzO],,,
= 2.17/3.21
- 0.68
indicating that the activity of water in 15 M ammonia is slightly less than its mole fraction. The activities of [NHd+] and [OH-] may be obtained from the aqueous equilibrium reaction NH3 + Hz0
= NH;
+ OH-,
k5 = 1.78 x 1O-5
(5)
If there is a total of N moles of ammonia species, then N( 1 -a) moles of NH3 will be in equilibrium with Ncl moles of [NH4+] and No! moles of OH-. Consequently, in a solution of N molality we may write ks = [NH4+][OH-]
WWW201
=
]N%+t$[NWon-]
[NC1-
@YNH~I[HZOI
(6)
Copper
and zinc equilibria
1311
In a 23.4 m ammonia solution, Nis 23.4, and from eqns (3) and (4), yNHl is 1.6 and [Hz01 is 0.68. The value of iVclis sufficiently small (i.e. small concentration of species), as will be evident below, that the values of You- and YNH+may be approximated to unity. Hence, after substituting 1.78 x 1O-5 for kS and solving gqn (6), ct is 9.0932 x 10F4, and [OH-] and [NH:] are both 2.128 x 1OK’. The [H+] activity in 23.4 m (15 M) ammonia may be calculated from the water equilibrium Hz0 = H+ + OH-, k7 = 9.945 x IO-l5 Substituting for [Hz01 =0.68 and [OH-] =2.128 x lo-* 3.178 x lo-l3 that is equivalent to a pH of 12.5 (pH= -log[H+]).
(7) yields a [H+] value of
Addition of copper powder
The dissolution of copper powder under fully oxygenated conditions in concentrated ammonia solutions leads to the formation of Cu”-ammines and OH- ions (pH changes) in proportions of 1:2, as shown in eqn (8), Cu + nNH3 + H20 + 1/202 + Cu(NH&+
+ 20H-
(8)
where n is the number of NH3 ligands on the ammine complex. When 1 g/L of copper is dissolved in 15 M ammonia it is equivalent to a concentration of 2.4606 x lo-* m of soluble copper. This concentration is sufficiently small that its corresponding y may be approximated to unity so that Cu(NHs)z+ is 2.4606 x lo-*. Equation (8) shows that 2.4606 x lo-* moles of soluble copper generate 4.9212 x lo-* moles of OH-. A portion of the OH-, x moles, will interact with the pre-existing 2.128 x lo-* m of NH,f via eqn (5) and the remainder will add to the pre-existing 2.128 x lo-* m of OH-. Hence, the new equilibrium with respect to ammonia becomes k = [2.128 x lo-* - x][2.128 x lo-* + 4.9212 x lo-* - X] = 1 78 x 1o_5 5 [37.4 - 2.128 x lo-* + x][HzO]
(9)
where [NH31 is [37.4 - 2.128x10-*+x], and [Hz01 is 0.68 according to eqn (4). Solving eqn (9) gives x = 1.336 x lo-* and [OH-] = 5.713 x lo-*. Therefore, from eqn (7) [H+], = 1.184 x lo-l3 and the pH is 12.93. Similarly, the dissolution of 6 g/L of copper is equivalent to 1.4764 x 10-l m of soluble copper, which corresponds to 1.4764 x 10-l [Cu(NHs)F] after approximating y to unity. Dissolution is accompanied by the generation of 2.9528 x 10-l moles of OH-, via eqn (8) of which a portion, x moles, interacts with the pre-existing NH: and the remainder adds to the pre-existing OH- so that eqn (9) is now changed to k = [2.128 x lo-* - x][2.128 x 10-l + 2.9528 x lo-* -xl 5 [37.4 - 2.128 x lo-* + x][HzO]
= 1 78 x 1o_5
(10)
Solving eqn (10) gives x = 1.976 x lo-’ and [OH-] = 2.968 x 10-l. Hence, from eqn (7) [Hf] = 2.278 x 10-14, and the pH is 13.642. Similar calculations for the dissolution of 10 g/L of copper show that Cu(NHs)i+ is 2.4606x lo-‘, [OH-] is 4.931 x lo-‘, [H+] is 1.372x 10Vi4, and the pH is 13.86. The principal results of all the activity calculations are summarised in Table 2.
D. Tromans
1312 Table 2.
Activities
in 15 M ammonia
Solution
[NH,1
15 M l5M I5 M 15 M
37.40 37.40 37.40 37.40
ammonia +lg/LCu +6g/LCu + 10 g/L Cu
solutions
E-PH
containing
copper powder
dissolved
as Cu”
[Hz01
[Cl?‘]
PH
0.68 0.68 0.68 0.68
2.4606 x IO-* 1.4764x 10-l 2.4606 x 10-l
12.50 12.93 13.64 13.86
EQUILIBRIA
All E-pH equilibria in the Cu-NH3-Hz0 and Zn-NHs-Hz0 systems in 15 M ammonia were based on [NH,] and [H,O] activities in Table 2, together with the standard chemical free energies per mole of species (u”) listed in Table 1. Most of the u” data for the copper species were obtained from the compilation by Bertocci and Wagman, with the notable exceptions of CuO and Cu(NHs):+. The appropriate u” for CuO is uncertain.48 We have used the value of - 129.7 kJ mol-’ listed by Barner and Scheurman,49 which is lower than that reported by King et a1.” (- 127.9 kJ mol-‘) and higher than the - 134 kJ mol-’ suggested by Bertocci and Wagman. The u” for Cu(NH3):’ was calculated from the equilibrium constant (log kl, = - 0.6) reported by Sillen and Martel15’ for eqn (11) at 25°C using the u” values for NH3 and Cu(NH3):’ listed in Table 2, CUE+
4
+ NH3 = Cu(NH):+,
k,, = 0.251
(11)
Thermodynamic data for the zinc system were obtained principally from the compilation by Brodd and Werth.52 When consulting Ref. 52, the reader should note that there are some typographical errors relating to the zincate ion, Zn(OH$, e.g. p” is listed as + 877.4 instead of - 877.4 kJ mol- ’ and the text states that the chemical free energy is - 209.7 kJ mol-- ’ instead of kcal mol- ‘. The resulting E-pH diagrams were constructed for the pH range from 12 to 15, because of the high pH encountered in 15 M ammonia solutions (Table 2). Solid lines de-lineate regions of stability of solid phases. Coarse broken lines show equilibria between dissolved (aqueous) species. Fine broken lines, labelled (a) and (b) correspond to the hydrogen and oxygen electrodes at one atmosphere fugacity and [Hz01 = 0.68. Cu-NHrH20 system Equilibrium E-pH diagrams for the Cu-NH3-H20 system are shown in Figs l-3 for 1, 6, and 10 g/L of dissolved Cu species, respectively. The pH of the corresponding Cu”containing solution in Table 2 is indicated on each figure by an arrow. The dominant soluble copper species in equilibrium with the oxides are the Cu”-pentammine and the Cu’diammine complexes. The oxides, CuO and CuzO, are stable in the higher pH regions and their stability range extends to lower pH values as the dissolved copper activity increases. Some of the important reaction equilibria are shown below in eqns (12)-( 18), Cu(NHs):+ Cu(NHs)t
+ e- = Cu(NHs)l
+ 3NHj,Ey,
= +0.0839 V(SHE)
+ e- = Cu + 2NH3, Ey3 = -0.1015
Cu20 + 4NH3 + 2H+ = 2Cu(NH3)2f
V(SHE)
+ H20, k,4 = 1.1202 x lOI
(12) (13) (14)
Copper
and zinc equilibria
13
1313
14
15
PH Fig. 1, Cu-NHsPH20 equilibria of dissolved Cu species (activity
at 25°C in 15 M ammonia = 2.4606 x 10m2). Arrow Table 2.
CuO + 5NH3 + 2H+ = Cu(NH&+
solution ([NHs] = 37.4) containing 1 g/L shows pH of corresponding solution in
+ H20, kls = 1.3707 x 10”
2Cu(NH3):+ + Hz0 + 2e- = Cu20 + lONH3 + 2H+, Eye = -0.4791 V(SHE)
(15) (16)
CuzO + 2H+ + 2e- = 2Cu + H20, Ey, = +0.4614 V(SHE)
(17)
2CuO + 2H+ + 2e- = CuzO + HzO, E;, = +0.6521 V(SHE)
(18)
------.-
12
%__@_
13
___ 14
15
PH Fig. 2. Cu-NHs-Hz0 equilibria of dissolved Cu species (activity
at 25°C in 15 M ammonia solution ([NH31 = 37.4) containing 6 g/L = I .4764 x lo-‘). Arrow shows pH of corresponding solution in Table 2.
D. Tromans
1314
13
15
14
PH Fig. 3. Cu-NH3-H20 equilibria at 25°C in 15 M ammonia solution ([NH,] = 37.4) containing 10 g/ I_ of dissolved Cu species (activity = 2.4606 x lo-‘). Arrow shows pH of corresponding solution in
Table 2.
where the standard hydrogen electrode
electrode potentials, (SHE) at 25°C.
E”, are referenced
with respect
to the standard
Zn-NH,H20 system Aqueous equilibria in the Zn-NHs-Hz0 system are shown in Fig. 4 for soluble zinc activities of lop6 and lo- ‘. No oxide phase is stable in the pH range of interest and zinc dissolves at much lower (more active) potentials than copper. The dominant soluble zinc species are the Zn”-tetrammine, Zn(NHs), 2f , below pH 13.656, and the Zn’t-zincate, Zn(OH)i-, above pH 13.656. Thus, both species should be formed upon dissolution of zinc
0.5 f-------------a
..+.____._____
___ 1
O-
t!
:
-0.5
j
?-----.-----.*A
9 ui
Zn(NH$
Zn(OH)f
-
._.._.._._._._.._.._ ___L
-I-1 -
b
-1.5
-
10.6
Zn -2
12
i
I
I
13
14
15
PH Fig. 4, Z~_NH~-H~O equilibria at 25°C in 15 M ammonia solution t[NH~I=37.4). Diagram constructed for 1x 10e6and 1 x 10-l activities of dissolved Zinc species.
Copper and zinc equilibria
1315
in the Cu-containing solutions in Table 2, depending on the solution pH. The Cu-containing solutions are of interest because stress corrosion studies of brass in these solutions will lead to dissolution of zinc. Important reaction equilibria are shown below in eqns (19)-(21), Zn(NH3)p
+ 4H20
Zn(NH# Zn(OH$
= Zn(OH)i-
+ 4NHs + 4H+, kis = 2.1814 x 1O-48
+ 2e- = Zn + 4NH3, Ei, = -1.0414 + 4H+ + 2e- = Zn + 4H20,
V(SHE)
Ei, = +0.3694 V(SHE)
(19) (20) (21)
DISCUSSION Copper equilibria The E-pH diagrams in Figs 1-3 are in remarkably good agreement with the reported behaviour of brass in Cu-containing 15 M ammonia solutions,‘5X’7,‘8,2’-28 indicating that they closely describe the stability of the copper species. For example, Fig. 1 shows that when Cu is added as 1 g/L of copper powder to form Cu(NH3):’ at a pH of 12.93 (Table 2) the only possible reduced forms of the pentammine at this pH are Cu(NHs)p or Cu (i.e. a nontarnishing condition), consistent with observations. Complete reduction of the pentammine to Cu would be consistent with the reported formation of brown films on brass. l5 Figures 2 and 3 show that when Cu(NH# species are formed by the addition of 6 and 10 g/L of copper powder, respectively, the oxide stability regions enlarge and the resulting pH of the solutions, (13.64 and 13.86) places the solutions in situations that now allows the pentammine to be reduced to CUZO (tarnishing condition), consistent with observations.“,“,‘s,*‘-2s Also, Fig. 3 shows that 10 g/L of copper powder may be dissolved in 15 M ammonia under oxygenated conditions without forming CuO, which was the upper limit of powder dissolution investigated by Pugh et a1.‘7,‘8 Furthermore, if Cu” is added as Cu(NO&, in amounts up to the equivalent of 10 g/L of copper, the pH of the solution should remain near 12.5, corresponding to that of 15 M ammonia alone (Table 2) because no hydroxyl ions are generated by reactions shown in eqn (8). Figure 3 shows it is not possible to reduce Cu(NHs):+ to Cu20 at pH 12.5. Therefore, the ammoniacal Cu(NO3)* solutions should be non-tarnishing, as observed by Pugh er a1.17+*8*21 Similarly, 15 M ammonia with additions of Cu(NO& followed by additions of Cu powder to reduce the Cu”-pentammine to Cu(NH3); should be non-tarnishing, as observed.29,30 The E-pH equilibria were consistent with the preparation ofnon-tarnishing 0.05 M Cu’ammine solutions by the dissolution of 0.025 M CuzO powder (= 3.18 g/L Cu) in 15 M ammonia.3’ The dissolution reaction is shown in eqn (22), Cu20 + 4NHs + H20 = 2Cu(NH$
+ 20H-
(22)
where equal moles of OHand Cu(NH3): are formed. In 15 M ammonia, 0.05 M Cu(NH3): is equivalent to 7.8125 x lo-* m, leading to formation of 7.8125 x lo-* m OHof which a portion (x) interacts with the pre-existing NH: and the remainder contributes to a change in pH, as described in (9) and (IO), producing a final pH of 13.13. Inspection of Fig. 3 shows that even when [Cu(NH3):] is increased to 2.4606 x 10-l (E 10 g/L Cu) no Cu20 is formed until the pH reaches 13.363, confirming that additions of 0.025 M Cu20 are readily soluble in 15 M ammonia, in good agreement with published data.3’
D. Tromans
1316
It is evident from Figs l-3 and Table 2 that two important consequences of the oxygenated dissolution of copper powder are the movement of the stability lines on the EpH diagrams and the increase in pH, both of which dictate whether the resulting solution produces tarnishing or non-tarnishing conditions. It was difficult to confirm precisely the calculated pH values listed in Table 2, because the usual glass pH-electrode behaves unreliably above pH 12. However, the glass electrode confirmed the magnitude and direction of the predicted pH changes, lending confidence to the calculated values. For example, an increase of approximately 1.2-l .3 pH units was measured after the dissolution of 6 g/L of copper powder, as compared to the calculated increase of 1.14 pH units.
Zinc equilibria Figure 4 shows that the domain of stability of zinc occurs at much lower (more active) potentials than the stability region of copper in Figs 1-3, consistent with the required active behaviour of zinc in de-alloying-related mechanisms of SCC of brass in 15 M ammonia solutions. 3om32735Y4’113 It is evident that the dominant soluble forms of zinc in these solutions may be either the Zn”-tetrammine or the Zn”-zincate complexes, depending upon whether the pH is below or above 13.656, respectively. The active behaviour of zinc, and effect of pH, was confirmed by conventional anodic polarisation tests at 24°C on metallographically polished pure metal (99.999% by weight) in 15 M ammonia (pH 12.5) and in 15 M ammonia with sufficient NaOH added to increase the pH to 14. The tests were conducted in a single compartment glass cell containing a Pt counter electrode. The working electrode surface was 90 mm2 and the potential was measured with respect to an external HgO/Hg reference electrode in 2.5 m NaOH (pH 14.6),53 where EHgopg= +0.063 V(SHE). The reference electrode was interfaced to the working solution via a salt bridge and polyethylene Luggin probe assembly containing aqueous NaOH. The probe was end-fitted with a sintered zirconia plug. Concentrated ammonia solutions have poor electrical conductivity,35.47 and it was necessary to apply corrections for the ohmic potential drop between the tip of the Luggin probe and the working electrode surface. The ohmic correction was obtained from a linear plot of the potential vs current density (i). The anodic polarisation curves are shown in uncorrected and ohmic-corrected forms in Fig. 5, where log (i) is plotted against E. Uncorrected data were collected at potentials up to + 1.O V(HgO/Hg). All the corrected data are presented in Fig. 5, whereas uncorrected data above -0.5 V(HgO/Hg) are not shown. Figure 5 shows very low open circuit potentials (E,,& consistent with the predictions of (- 1.321 V(SHE)) and - 1.409 V(HgO/Hg) Fig. 4. The E,, were - 1.384 V(HgO/Hg) (- 1.346 V(SHE)) in the 15M NH3 and 15 M NH3 + NaOH solutions, respectively, where the slightly lower potential in the NaOH containing solution was consistent with the more active behaviour associated with the dominance of Zn(OH$ complexes at pH 14 (see Fig. 4). Furthermore, the ohmic-corrected curve in the NaOH-containing solution exhibited anodic Tafel behaviour with a slope of 41.3 mV at i< 32 A m-‘. This was close to the 40 mV reported by Bockris et aZ.54 for anodic behaviour of zinc in alkaline KOH solutions, where the rate determining anodic step was shown to be the transfer of the second electron according to eqn (23), Zn(OH>; The steeply rising section
+ OH-
+ Zn(OH),
of the ohmic-corrected
+ e-
polarisation
(23) curve at higher
i in the
Copper
H
-3
and zinc equilibria
Comcted
-2
-1 log(i),
Fig. 5.
Anodic
1317
0
1
2
3
(A.m -2)
polarisation behaviour of zinc in 15 M NH3 and 15 M NH3 + NaOH uncorrected and ohmic-corrected data. Scan rate 1 mVs- ’
showing
NaOH-containing solution indicated the possibility of mass transfer controlled anodic behaviour, possibly the transfer of OH- ions to the electrode surface. In contrast, the ohmic-corrected polarisation curve in 15 M ammonia did not exhibit a Tafel region and the anodic currents were considerably lower than when NaOH was present. The shape of the polarisation curve suggested that anodic behaviour was falling under mass transfer control much earlier (i.e. at lower current densities) than in the NaOH-containing solution, consistent with the lower concentration of OH- ions at pH 12.5. This implied that anodic dissolution was still determined by the rate of formation of zincate species, even in a pH regime where Zn(NH# complexes were the dominant soluble species and there was a high concentration of NHs available for their formation. If this interpretation is correct, transient zincate species may be formed initially at the electrode surface, which then interact with NH3 in solution to form the stable tetrammine complex via eqn (24) Zn(OH)i-
+ 4NHs + Zn(NHs)F
+ 40H-
(24)
CONCLUSIONS (1) The calculated [NHs] and [Hz01 activities in 15 M ammonia solutions are 37.4 and 0.68, respectively. (2) Based on the above activities, new E-pH diagrams were constructed for the CuNHs-Hz0 and Zn-NHs-Hz0 systems in 15 M ammonia solutions in the pH range 12-15. These diagrams are consistent with the reported behaviour of cl-brass in these solutions in relation to tarnish (CuzO) formation, non-tarnishing conditions, and the active behaviour of the zinc component. (3) The new equilibrium diagrams are particularly useful for understanding tarnishing and non-tarnishing phenomena in copper-containing 15 M ammonia solutions of the type that are used frequently for studying SCC of a-brass. (4) Anodic dissolution of zinc in 15 M ammonia to form the Zn”-tetrammine complex may involve the formation of transient zincate species.
1318
D. Tromans
Acknowledgements-The author wishes to thank the Natural Canada for financial support. Also, the experimental assistance
Sciences and Engineering of Mr C. Fong is gratefully
Research Council acknowledged.
of
REFERENCES 1. J. 2. 3. 4. 5. 6. 7. 8. 9. IO. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33.
34. 35. 36. 37. 38. 39. 40. 41.
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