Journal of Materials Science and Chemical Engineering
Vol.06 No.01(2018), Article ID:82075,10 pages

Electrochemical Study of Redox Reaction of Various Gold III Chloride Concentrations in Acidic Solution

Afolabi Ayeni, Shafiq Alam*, Georges Kipouros

Department of Chemical and Biological Engineering, University of Saskatchewan, Saskatoon, Canada

Copyright © 2018 by authors and Scientific Research Publishing Inc.

This work is licensed under the Creative Commons Attribution International License (CC BY 4.0).

Received: December 13, 2017; Accepted: January 26, 2018; Published: January 29, 2018


The redox reaction of gold III chloride in acid solutions has been electrochemically investigated using a cyclic voltammetry technique. This paper em- phasizes the current and potential sites at which gold III chloride is reduced in hydrochloric acid that is vital to electrochemical evaluation of gold recovery. The solutions were prepared by reacting HCl with AuCl3 in various concentrations thus 30 and 60 mg/L AuCl3 in 0.1 and 0.5 M HCl, respectively. Solutions of 0.1 and 0.5 M HCl containing 0, 30 and 60 mg/L AuCl3, respectively were tested for possible reduction and oxidation reactions by cyclic voltammogram experiment using a glassy carbon, a saturated calomel and a platinum wire mesh as working, reference and counter electrodes, respectively. The results showed no peak in the case of the absence of AuCl3 in the solutions, but appreciable cathodic and anodic peaks for the reduction and oxidation of various concentrations of AuCl3 in acid solutions. The reaction between AuCl3 and HCl was found to be reversible because the ratio of oxidation peak current and reduction peak current was 1. The concentration of AuCl 4 on the surface of the working electrode at the reduction site for each AuCl3 concentration using Nernst equation was 1.22 × 109 ppm and 2.44 × 109 ppm. The reduction potentials were independent of concentration, while the current was highly dependent of concentration.


Redox Reaction, Cyclic Voltammetry, Nernst Equation, Gold, Chloride

1. Introduction

The demand for gold in the global market has jolted researchers into gold recovery methods from either a lean ore or waste products of consumer electronics (urban mining). Sometimes, synthetic solutions are used to study the extraction of gold in a preliminary laboratory experiment. Gold III chloride is the auric salts commonly used to achieve this purpose [1] . Hydrochloric acid has been the popular leachant for precious metals from secondary sources [2] . Hence, in this study, solutions of gold III chloride in HCl were prepared to investigate reversible and redox reactions using a cyclic voltammetry technique.

Few studies have been carried out on the electrochemical reduction to metal, such as the electrochemical reduction of silver from iodide solutions [3] . Fourcade and Tzedakis [3] used a potentiostat as the electrochemical apparatus with a silver disk working electrode, a saturated calomel reference electrode and a platinum counter electrode to measure all the electrode potentials during the adsorption experiments. Tao et al. [4] reported scanning tunneling microscopy (STM) and electrochemical study of the interplay between redox properties, adsorption, and self-assembly processes of porphins on Au surfaces.

The cyclic voltammetry technique is generally used to study the electrochemical properties of an analyte in solution [5] [6] [7] . The theory of voltammetric methods is based on the solution of the Nernst Equation (1). Voltammetry is a method in which information about an analyte is obtained by measuring the current generated as the applied potential to the working electrode is varied. Potential is measured between the working electrode and the reference electrode, while current is measured between the working and the counter electrode [8] [9] . The Nernst Equation is expressed in terms of potential at the working electrode.

E = E 0 R T n F ln Q (1)

where, E = measured potential, E0 = standard electrode potential, R = gas constant, T = temperature (˚K), Q = reaction quotient, n = number of electrons exchanged, and F = Faraday’s constant.

The redox reaction and half-cell reaction of gold III chloride in HCl are represented in reactions R1 and R2 respectively thus,

AuCl 3 + HCl [ AuCl 4 ] + H + (R1)

For half-cell reaction,

AuCl 3 + e AuC 4 (R2)

The result from the experiment showed no reasonable cathodic or anodic peak for hydrochloric acid solution without the presence of gold III chloride, while peaks were observed during the measurement of various concentrations of the gold III chloride in hydrochloric acid solution. It could be deduced from the experiment the reaction of gold III chloride with HCl is a redox and reversible reaction.

The objective of this work was to investigate the reaction processes in gold III chloride acid solution using cyclic voltammetry.

2. Experimental

2.1. Materials and Instrumentations

Glassware, analytical grade hydrochloric acid, gold (III) chloride, and de-ionized water were used for the preparation of the solutions. The electrochemical equipment consisted of a PAR 283 Potentiostat/Galvanostat (PS/GS) and a Solartron 1260 Frequency response analyzer, Glassy carbon electrode was used as working electrode (WE), a platinum mesh served as counter electrode (CE), while a saturated calomel electrode (SCE) was the reference electrode (RE).

2.2. Experimental Procedure

2.2.1. Preparation of Solution

Solutions of 30 ppm and 60 ppm of AuCl3 in 200 mL of 0.1 and 0.5 M HCl were prepared. The reaction between AuCl3 and HCl is a redox as shown in R1. AuCl3 was reduced to AuCl 4 and HCl was oxidized to H+. This redox reaction was electrochemically measured by subjecting the gold chloride solution to cyclic voltammogram experiment, in which the cathodic and anodic current peaks were determined relative to applied potentials. This was achieved with a core driven software PAR 283 Potentiostat/Galvanostat (PG/GS) and a Z-plot driven solartron instrument [3] . Cleaning of the electrodes was done prior to the experiment, for the purpose of revealing the surface of the electrodes which might have been covered by impurities, and conditioned them for the experiment as reported by Feng et al., [10] . The 0.5 M HCl was prepared, and the electrodes were immersed in the diluted solution. The electrodes were then connected to the Potentiostat which was set at cyclic voltammogram experiment mode for cleaning. The vertex potential 1 and 2 were set at −0.25 and 1.25 V respectively at a scan rate of 5.0 mV/s. The cleaning was done for 1 h.

Subsequently, about 100 mL of 30 ppm AuCl3 (in 0.1 M HCl) was poured into a 250 mL beaker. The electrodes were immersed into the solution and connected with connecting cables to the Instrument accordingly.

2.2.2. Cyclic Voltammetry Measurement

The cyclic voltammogram experiment was performed with a PAR potentiostat/galvanostat. The schematic of the setup is as shown in Figure 1. The applied potential was set between −0.25 V and 1.25 V. The scan rate was set at 5 mV/s, and the No. of cycles was maintained at 1 mV/point. During the measurement, the scanning of the potential was from −0.25 to 1.25, and then back to −0.25 at a rate of 5 mV/s for 1 cycle. This procedure was performed on all the prepared solutions and the measurement of peaks was recorded via the plot of current density as a function of the potential. The experiment was carried out at a room temperature of 25˚C.

3. Results and Discussion

The experimental results are depicted in Figures 1-6. The system moved through the various dynamic regime (start-finish), and the oxidation and reduction

Figure 1. Cyclic voltammetry measurement of 30 ppm AuCl3 in 0.1 M HCl at room temperature (25˚C).

Figure 2. Cyclic voltammetry measurement of 60 ppm AuCl3 in 0.1 M HCl solution at room temperature (25˚C).

Figure 3. Cyclic voltammetry measurement of 0 ppm AuCl3 in 0.1 M HCl solution at room temperature (25˚C).

Figure 4. Cyclic voltammetry measurement of 30 ppm AuCl3 in 0.5 M HCl solution at room temperature (25˚C).

Figure 5. Cyclic voltammetry measurement of 60 ppm AuCl3 in 0.5 M HCl solution at room temperature (25˚C).

peaks could be observed through the voltammogram. The figures represent the current generated by applied cyclic potentials on 30 ppm, 60 ppm, 0 ppm AuCl3 in 0.1 and 0.5 M HCl, respectively. Where, Epa and Epc were oxidation and reduction potential peaks, respectively, and Ipa and Ipc were the corresponding current at the peak of oxidation and reduction in that order.There are reports on the formation of these asymmetry peaks during voltammetric measurements [3] [4] [11] .

Figure 6. Cyclic voltammetry measurement of 0 ppm AuCl3 in 0.5 M HCl solution at room temperature (25˚C).

In Figure 1, oxidation peak was reached at about 1 V (Epa), and 0.0028 A/cm2 current density (Ipa) was generated. At this point the HCl was oxidized completely to H+. On the other hand, the complete reduction of AuCl3 to AuCl 4 was achieved at a potential Epc and current Ipc of 0.55 V and 0.0028 A/cm2, respectively. Khunathai et al. reported the standard reduction potential E0 of AuCl 4 to be 1.0 V [11] . The Nernst equation (1) was employed to determine the concentration of reduced auric chloride ion ( AuCl 4 ) at the surface of the working electrode in 30 ppm solution, thus:

E = E 0 R T n F ln Q

From the half-reaction depicted in reaction R2,

ln Q = ln concentrationofAuCl 4 concentrationofAuCl 3 = ( E 0 E ) n F R T = ( 1 0.55 ) × 1 × 96500 8.314 × ( 273 + 25 ) = 17.52

ConcentrationofAuCl 4 = 30 × e 17.52 = 1.22 × 10 9 ppm

For a reversible reaction [6] ,

I p c : I p a = 1 (2)

In Figure 1, I p c = 0.0035 0.00075 = 0.00275

I p a = 0.004 0.00125 = 0.00275 .

Hence, I p c : I p a = 0.00275 : 0.00275 = 1 .

From this experiment, 30 ppm of AuCl3 would be reduced to AuCl 4 at the reduction peak on the surface of the working electrode at a concentration of 1.22 × 109 ppm. The reaction was also a reversible reaction, considering the ratio of Ipc to Ipa which was found to be 1 [6] .

Figure 2 shows voltammetry measurement of 60 ppm AuCl3 in 0.1 M HCl solution, with the Epc and Epa obtained at 0.55 V and 1.23 V, respectively. The corresponding cathodic and anodic peak current density (Ipc and Ipa) were 0.0022 and 0.0024 A/cm2. From Equations (1) and (2) respectively, the concentration of the AuCl 4 at reduction peak on the surface of the electrode was 2.44 × 109 ppm, and the ratio of Ipc to Ipa was 1. Hence, the reaction was also a reversible one.

Figure 3 depicts the voltammetry measurement of 0.1 M HCl without AuCl3 (0 ppm AuCl3), no peak was feasible either at the oxidation zone or reduction zone. This confirmed electrochemically, that the reaction between HCl and water was not a redox reaction but a dissociation

HCl + H 2 O H 3 O + + Cl (R3)

Figure 4 and Figure 5 represented the higher concentration of AuCl3 and HCl, however the values of Epc was the same as that of Figure 1 and Figure 2 (0.55 V), indicating evidence that the standard potential and reduction potential peaks did not change with concentration. The reduction peak current varied with concentration as shown in Figure 1 and Figure 4 was −0.0035 and −0.0065 A/cm2, respectively. Figure 6, just like Figure 3, had no peak because it was not a redox reaction.

Summary of Cyclic Voltammetry (CV) Curves for Various Concentrations

The form of CV curves as depicted in Figure 7 are like those published in various literature under similar experimental conditions [11] [12] [13] [14] [15] .

4. Conclusion

The redox reaction of AuCl3 in HCl solution of various concentrations has been electrochemically studied using a cyclic voltammetry technique. The concentration of the reduced AuCl 4 ion was determined using the Nernst equation. The results showed that anodic and cathodic peaks were present in the solution containing AuCl3 of various concentrations, while no peaks were generated in HCl solution in the absence of AuCl3. The measured Epc in the tested solutions was the same showing evidence that the reduction potentials were independent of the concentration. However, the reduction current varied with the concentration of solution, affirming the dependency of current on concentration. This study has interesting implication to determine the electrochemical parameters of gold reduction during leaching and adsorption processes with more accurate results obtained for specific process.

(a) (b) (c) (d)
(e) (f)

Figure 7. Cyclic voltammetry curves of (a) 30 ppm AuCl3 in 0.1 M HCl; (b) 60 ppm AuCl3 in 0.1 M HCl; (c) 30 ppm AuCl3 in 0.5 M HCl; (d) 60 ppm AuCl3 in 0.5 M HCl; (e) 0 ppm AuCl3 in 0.1 M HCl and (f) 0 ppm AuCl3 in 0.5 M HCl at room temperature.


The equipment was calibrated with the support of Prince Yuan Ding (Dalhousie University, Halifax) and Will Judge (University of Toronto).

Cite this paper

Ayeni, A., Alam, S. and Kipouros, G. (2018) Electrochemical Study of Redox Reaction of Various Gold III Chloride Concentrations in Acidic Solution. Journal of Materials Science and Che- mical Engineering, 6, 80-89.


  1. 1. Barakat, M.A. and Mahmoud, M.H.H. (2004) Recovery of Platinum from Spent Catalyst. Hydrometallurgy, 72, 179-184.

  2. 2. Adhikari, B.B., Gurung, M., Alam, S., Tolnai, B. and Inoue, K. (2013) Kraft Mill Lignin—A Potential Source of Bio-Adsorbents for Gold Recovery from Acidic Chloride Solution. Chemical Engineering Journal, 231, 190-197.

  3. 3. Fourcade, F. and Tzedakis, T. (2000) Study of the Mechanism of the Electrochemical Deposition of Silver from an Aqueous Silver Iodide Suspension. Journal of Electroanalytical Chemistry, 493, 20-27.

  4. 4. Ye, T., He, Y. and Bourguet, E. (2006) Adsorption and Electrochemical Activity: An In-Situ Electrochemical Scanning Tunneling Microscopy Study of Electrode Reactions and Potential-Induced Adsorption of Porphyrins. Journal of Physical Chemistry, 110, 6141-6147.

  5. 5. Bard, A.J. and Faulkner, L.R. (2000) Electrochemical Methods, Fundamentals and Applications. 2nd Edition, Wiley, Hoboken.

  6. 6. Nicholson, R.S. and Irving, S. (1964) Theory of Stationary Electrode Polarography: Single Scan and Cyclic Methods Applied to Reversible, Irreversible, and Kinetic Systems. Analytical Chemistry, 36, 706-723.

  7. 7. Heinze, J. (1984) Cyclic Voltammetry—“Electrochemical Spectroscopy”: New Analytical Methods. AngewandteChemie International Edition, 23, 831-847.

  8. 8. Kissinger, P. and William, R.H. (1996) Laboratory Techniques in Electro-analytical Chemistry. 2nd Edition, Revised and Expanded, CRC, Boca Raton.

  9. 9. Zoski, C.G. (2007) Handbook of Electro-chemistry. Elsevier Science, Amsterdam.

  10. 10. Feng, J., Gao, Q., Lv, X. and Epstein, I.R. (2008) Dynamic Complexity in the Electrochemical Oxidation of Thiourea. Journal of Physical Chemistry, 112, 6578-6585.

  11. 11. Khunathai, K., Matsueda, M., Biswas, B.K., Kawakita, H., Ohto, K., Harada, H., Inoue, K., Funaoka, M. and Alam, S. (2011) Adsorption Behavior of Lignophenol Compounds and Their Dimethylamine Derivatives Prepared from Rice and Wheat Straws for Precious Metal Ion. Journal of Chemical Engineering of Japan, 44, 781-787.

  12. 12. Ahmed, M.I., Aziz, A.M., Helal, A. and Sheikh, M.N. (2016) Direct Electrodeposition of Nanogold on Gallium-Doped Zinc Oxide by Cyclic Voltammetry and Constant-Potential Techniques: Application to Electro-Oxidation of Sulfite. Journal of the Electrochemical Society, 163, D277-D281.

  13. 13. Naumowicz, M. (2016) Cyclic Voltammetry and Chloroamperometry Techniques in Description of the Surface-Active Phospholipid Bilayer Relative to Acid-Base Equilibria. Journal of the Electrochemical Society, 163, H750-H756.

  14. 14. Lee, J.B. and Kim, S.W. (2007) Semiconducting Properties of Passive Films Formed on Fe-Cr Alloys Using Capacitance Measurements and Cyclic Voltammetry Techniques. Materials Chemistry and Physics, 104, 98-104.

  15. 15. Sullivan, A.M. and Kohl, P.A. (1997) Electrochemical Study of the Gold Thiosulfate Reduction. Journal of the Electrochemical Society, 144, 1686-1690.