Question 3 of 6: Eh-pH diagrams for the Au-, Ag-, Cu- and Fe-CN-H2O systems

Problem 3 — Eh-pH diagrams for the Au-, Ag-, Cu- and Fe-CN-H2O systems (20 marks)

Question text not reproduced: the examination questions are © Engineers and Geoscientists BC. Open the official past paper (linked at the top of this page) to read the question, then follow the worked solution below.

Two water-stability reference lines recur on every one of the four diagrams and are not species-specific: $\text{O}_2/\text{H}_2\text{O}$ at $E_h = 1.23-0.0591\,\text{pH}$ (upper bound, aerated conditions), and $\text{H}^+/\text{H}_2$ at $E_h=-0.0591\,\text{pH}$ (lower bound). Aerated cyanide-leach pulps typically sit well above the $\text{H}^+/\text{H}_2$ line (measured Eh ≈ +0.2 to +0.4 V vs. SHE, well below the thermodynamic $\text{O}_2/\text{H}_2\text{O}$ line because O₂ reduction kinetics are sluggish) — this aerated-Eh window is the assumption used throughout parts (a)-(b) below.

[Figure not reproduced: Fig. 2 — schematic M(CN) x - /M(s) boundary (dashed red, pH-independent) for each of the four systems, plotted against the water-stability lines and marked at its own electrolysis/H2-reduction crossover pH c . The full multi-species diagrams on the exam page carry additional oxide/hydroxide fi. See the official exam paper.]

(a) Cu in Au(CN)₂⁻, pH 9.3, WITH air. (i) With air present the mixed corrosion potential is pinned near the $\text{O}_2/\text{H}_2\text{O}$ line (≈+0.2 to +0.4 V), which sits well above BOTH the $\text{Cu(CN)}_2^-/\text{Cu}$ line (-0.44 V) and the $\text{Au(CN)}_2^-/\text{Au}$ line (-0.60 V); the copper is therefore not thermodynamically stable and corrodes, but because O₂ reduction is the kinetically dominant cathodic partner (not the slower $\text{Au(CN)}_2^-$ reduction), the copper oxidizes into its OWN cyanide complex rather than cementing the dissolved gold — no significant gold deposition occurs, and the copper instead contaminates the pregnant solution while consuming cyanide and oxygen (the well-documented "soluble copper" problem in gold cyanidation). (ii) $$4\text{Cu(s)} + 8\text{CN}^- + \text{O}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{Cu(CN)}_2^- + 4\text{OH}^-$$

(b) Ag in Au(CN)₂⁻, pH 8, WITH air. (i) The same reasoning applies: aerated Eh (≈+0.2 to +0.4 V) sits above the $\text{Ag(CN)}_2^-/\text{Ag}$ line (-0.31 V), so silver also corrodes to its cyanide complex rather than staying metallic; since -0.31 V is LESS negative than copper's -0.44 V, silver needs a smaller overpotential to corrode and dissolves comparatively more readily. As in (a), O₂ (not $\text{Au(CN)}_2^-$) is the dominant cathodic reaction, so silver simply co-dissolves as $\text{Ag(CN)}_2^-$ alongside the gold rather than cementing it — matching the well-known observation that silver reports to the pregnant solution in most gold-silver cyanidation circuits. (ii) $$4\text{Ag(s)} + 8\text{CN}^- + \text{O}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{Ag(CN)}_2^- + 4\text{OH}^-$$

(c) Au in Fe(CN)₆³⁻, pH 9.3, NO air. (i) Even without air, the ferri/ferrocyanide couple ($E^{\circ}=+0.36$ V) sits well above the $\text{Au(CN)}_2^-/\text{Au}$ line (-0.60 V) — ferricyanide alone is thermodynamically capable of oxidizing metallic gold to $\text{Au(CN)}_2^-$, given free cyanide is available, so gold dissolves via this auxiliary oxidant with no oxygen required (ferricyanide has been used industrially as an auxiliary/alternative oxidant for oxidant-starved gold leaches). (ii) $$\text{Au(s)} + 2\text{CN}^- + \text{Fe(CN)}_6^{3-} \rightarrow \text{Au(CN)}_2^- + \text{Fe(CN)}_6^{4-}$$

(d) Fe in Au(CN)₂⁻, pH 9.3, NO air. (i) Metallic iron's own line ($E^{\circ}\approx -1.3$ V, complexed by cyanide) sits far BELOW the $\text{Au(CN)}_2^-/\text{Au}$ line (-0.60 V), so iron is a strong enough reductant to cement gold out of solution even with no air present — iron dissolves (as the stable hexacyanoferrate(II) complex, the large stability field shown on the Fe-CN-H₂O diagram at this pH) while metallic gold is deposited. This is the classical cementation mechanism (historically iron/steel scrap, and later zinc dust in the Merrill-Crowe process, both exploit exactly this potential gap; cementation is deliberately run oxygen-free for the same reason aeration suppressed cementation in parts (a)-(b)). (ii) $$\text{Fe(s)} + 2\text{Au(CN)}_2^- + 2\text{CN}^- \rightarrow \text{Fe(CN)}_6^{4-} + 2\text{Au(s)}$$

Approach for (iii)/(iv). Because each $\text{M(CN)}_x^-/\text{M}$ boundary is a horizontal (pH-independent) line while $E_h(\text{H}^+/\text{H}_2)=-0.0591\,\text{pH}$ falls steadily with increasing pH, the two lines cross at exactly one pH: $\text{pH}_c=-E^{\circ}(\text{M})/0.0591$. For $\text{pH}>\text{pH}_c$ the metal's own line sits above the $\text{H}_2$ line, so a cathode can be set to deposit the metal before H₂ evolves (electrolysis works with good current efficiency). For $\text{pH} \lt \text{pH}_c$ the $\text{H}_2$ line sits above the metal's line, so H₂ gas is thermodynamically capable of reducing the complex to metal (chemical/hydrogen reduction is favoured) — the two regimes are exact mirror images of each other about $\text{pH}_c$.

1. Compute each crossover pH. $$\text{pH}_c=\frac{-E^{\circ}(\text{M})}{0.0591}$$ For Au ($E^{\circ}=-0.60$ V): $\text{pH}_c=\boxed{10.15}$. For Ag ($E^{\circ}=-0.31$ V): $\text{pH}_c=\boxed{5.25}$. For Cu ($E^{\circ}=-0.44$ V): $\text{pH}_c=\boxed{7.45}$. For Fe ($E^{\circ}\approx -1.3$ V): $\text{pH}_c=22.0$, outside the 0-14 window entirely.
2. (iii) Electrolysis pH ranges (metal deposits above the H₂ line). Au: efficient only for $\text{pH}>10.15$ (explains why direct electrowinning of dilute, moderate-pH cyanide pregnant solution is inefficient in practice — zinc cementation is preferred instead, except from concentrated, high-pH strip liquors). Ag: efficient for $\text{pH}>5.25$, i.e. essentially the ENTIRE practical cyanide-leach range (pH 8-11) — silver electrowinning is comparatively easy. Cu: efficient for $\text{pH}>7.45$, workable at typical alkaline leach pH but marginal below it. Fe: since $\text{pH}_c=22$ lies outside 0-14, the H₂ line is ABOVE the Fe line at every pH from 0 to 14 — iron can NEVER be electrowon efficiently from a cyanide bath; hydrogen co-evolution always dominates.
3. (iv) Hydrogen-reduction pH ranges (H₂ line above the metal line). This is the mirror condition, $\text{pH} \lt \text{pH}_c$. Au: thermodynamically favoured for $\text{pH} \lt 10.15$, i.e. across essentially the whole practical leach range — though H₂ gas reduction of dissolved gold cyanide is kinetically far too slow to be used industrially at ambient conditions (this is exactly why solid-metal cementation, part (d), is used instead of bubbling H₂ gas). Ag: only favoured below $\text{pH}=5.25$ — NOT achievable at typical alkaline leach pH without first destroying the cyanide (HCN volatilization hazard). Cu: only below $\text{pH}=7.45$, similarly impractical at leach pH. Fe: favoured at EVERY pH from 0 to 14 (since $\text{pH}_c=22$ is never reached) — thermodynamically H₂ could always reduce the iron complex, but again kinetics rule this out industrially, which is why metallic iron itself (not H₂ gas) is the practical cementation reagent in part (d).