CHAPTER 3.10

Thermodynamic Basis of Electrolysisand ElectrochemistryTetsuji HiratoDepartment of Energy Science and Technology, Kyoto University, Kyoto, Japan

3.10.1. ZINC ELECTROWINNING

The thermodynamic basis of zinc electrowinning can be understood from the

potential–pH diagram shown in Figure 1 of Chapter 3.9. In the process, the anode reac-

tion is oxygen evolution, which is shown as line b in the potential–pH diagram. The reac-

tion proceeds at potentials higher than line b. The cathode reaction is deposition of metal

zinc, which is shown as line 1 in the potential–pH diagram because the concentration of

Zn2þ(aq) is relatively high in solutions used in this process.

The reaction proceeds at potentials lower than line 1. Since this potential region is

lower than line a, it is easily understood that hydrogen evolution cannot be avoided

in this process.

The voltage required in this process is the difference between lines b and 1. The volt-

age increases with decreasing pH value, and at pH 0, a voltage of at least 1.99 V is required

for progression of the zinc electrowinning process. Of course, more voltage is required to

operate the real process, including overpotential of zinc deposition, and an IR drop in a

circuit with an electric current.

This diagram can also be used to consider the removal of Zn2þ fromwaste solutions as

Zn(OH)2. In the pH area between lines 4 and 5, the concentration of zinc in the solution

can be reduced to less than 10�5.

3.10.2. COPPER ELECTROWINNING

The thermodynamic basis of copper electrorefining can be understood from the

potential–pH diagram shown in Figure 3.10.1. In this process, the anode reaction is

oxygen evolution, which is shown as line b in the potential–pH diagram. The reaction

proceeds at potentials higher than line b. The cathode reaction is deposition of metal cop-

per shown as line 1 in Figure 3.10.1. The reaction proceeds at potentials lower than line

1. The potential indicated by line 1 is higher than line a, and thus hydrogen evolution

does not need to be considered in this process. The voltage required in this process is the

Treatise on Process Metallurgy, Volume 1 © 2014 Elsevier Ltd.http://dx.doi.org/10.1016/B978-0-08-096986-2.00037-0 All rights reserved. 653

Figure 3.10.1 Potential–pH diagram for the Cu–H2O system at 298 K.

654 Tetsuji Hirato

difference between lines b and 1. The voltage increases with decreasing pH value, and at

pH 0, a voltage of at least 0.899 V is required for progression of the copper electrowin-

ning process, which is much lower than that in zinc electrowinning process.

3.10.3. COPPER ELECTROREFINING

The thermodynamic basis of copper electrorefining can also be understood from

the potential–pH diagram shown in Figure 3.10.1. In this process, anode and cathode

reactions are expressed as follows:

at anode : Cu¼Cu2þþ2e

at cathode : Cu2þþ2e¼Cu

Both reactions are indicated by line 1 in Figure 3.10.1 and the required voltage is

0 from a thermodynamic understanding. More strictly, the activity of anode copper is

less than 1 because of impurities, so the voltage is a little bit more than 0.

3.10.4. ELECTROCHEMISTRY IN LEACHING

The electrochemistry in acid leaching or dissolution of copper oxides and hydrox-

ides can also be understood by the potential–pH diagram shown in Figure 3.10.1. The

dissolution reactions of copper oxides and hydroxides are as follows:

Cu OHð Þ2 sð Þþ2Hþ ¼Cu2þ aqð Þþ2H2O lð Þ ð3:10:1Þ

655Thermodynamic Basis of Electrolysis and Electrochemistry

Cu2O sð Þþ2Hþ aqð Þ¼ 2Cu2þ aqð ÞþH2O lð Þþ2e ð3:10:2ÞReactions (3.10.1) and (3.10.2) are indicated by lines 2 and 3 in the diagram, respec-

tively. Reaction (3.10.2) is a redox reaction as well as an acid–base reaction. For the dis-

solution of Cu2O to proceed, oxidant is required. When powders of Cu2O are put into

an aqueous sulfuric acid solution from which oxygen is purged, the following dispropor-

tionation reaction occurs:

Cu2O sð Þþ2Hþ aqð Þ¼Cu sð ÞþCu2þ aqð ÞþH2O lð Þ ð3:10:3ÞThe thermodynamic properties of sulfide minerals in an aqueous environment can

also be described by means of potential–pH diagrams. Such diagrams are as useful to

hydrometallurgists as are Ellingham diagrams to pyrometallurgists [1]. For example,

Figure 3.10.2 shows the potential–pH diagram for the Zn–S–H2O system. This diagram

is constructed by considering that zinc in zinc sulfide is decomposed to Zn2þ, Zn, and Zn(OH)2 and sulfur to H2S, S, HSO�

4 , or HS�. As can be seen from this figure, zinc sulfide

can be decomposed by oxidation, acidification, or reduction. The decomposition of zinc

sulfide is represented by the following reactions:

Oxidation

Figu

ZnS¼Zn2þþSþ2e�

ZnSþ4H2O¼Zn2þþHSO�4 þ7Hþþ8e�

ZnSþ4H2O¼Zn2þþSO2�4 þ8Hþþ8e�

ZnS

Zn(OH)2

Zn2+

SO42-HSO4

-

Zn2+

Zn

E/V

vs.

SH

E

S

0.8

0

-0.8

-1.60 4-4 8 12

pH

H2S

H2S

HS-

re 3.10.2 Potential–pH diagram for the Zn–S–H2O system at 298 K.

656 Tetsuji Hirato

Acidification

ZnSþ2Hþ ¼Zn2þþH2S

Reduction

ZnSþ2Hþþ2e� ¼ZnþH2S

ZnSþ2e� ¼ZnþS2�

In oxidative leaching, oxygen or Fe3þ is used as an oxidant. In real oxidative leaching of

sulfides, the stable area of elemental sulfur seems larger than shown in Figure 3.10.2. This

is caused by the slow reaction of oxidation from elemental sulfur to sulfate. This diagram

can also be used to consider the removal of Zn2þ fromwaste solutions as ZnS. In this case,

the activity of Zn2þ(aq) should be set at 10�6.

REFERENCES[1] E. Peters, Metall. Trans. B 7B (1976) 505–517.