InTERnIITIO01# JOURIllaLOF

mlnERHL PRO[ESSln6ELSEVIER Int. J. Miner. Process. 49 (1997) 171-183

Study of Merrill-Crowe processing. Part I: S,olubility of zinc in alkaline cyanide solutionGexla Chi *, Maurice C. Fuerstenau, John O. MarsdenUniversi~ of Nevada-Reno, Mackay School of Mines, Reno, NV 89557, USA

Abstract

A comprehensive study of zinc solubility was carded out in de-aerated alkaline cyanide solutions. Parameters investigated were: dissolved oxygen, cyanide and lead nitrate concentrations and pH Zinc solubility increases significantly with increased dissolved oxygen and cyanide concentrations. Eh-pH diagrams have been developed and used to predict regions of stability of zinc hydroxide, which may inhibit gold precipitation. The interaction between cyanide and hydroxyl ion activities to prevent the formation of zinc hydroxide is presented. With lead nitrate additions, maximum solubility of zinc occurs in the range of 10-15 mg/l lead nitrate. SEM photomicrographs of zinc particles in the presence of lead nitrate are presented.Keyword~: zinc solubility; Eh-pH diagrams; Merrill-Crowe processing

1. Introduction

Precious metal recovery from aqueous solution by zinc dust precipitation has been in use since the late 1800s [1]. This technology, originally patented by Salman and Pichard, was applied in 1897 to the Homestake operation in Lead, South Dakota by C.W. Merrill. In 1916, the process was refined with the introduction of T.B. Crowe's vacuum deaerator with which considerable improvement in efficiency was effected [1]. Since then, the Merrill-Crowe process has seen extensive application in the precipitation of precious metals from cyanide solution. Merrill-Crowe systems operate effectively most of the time, but at times and for unexplained reasons, precious metals are not precipitated efficiently. In other cases, well-de, signed systems do not achieve the expected low, barren-solution grades. When these phenomena occur, large quantities of precious metals can be retained in solution

* Corresponding author. 0301-75 L6/97/$17.00 Copyright 1997 Elsevier Science B.V. All rights reserved. Pll S0301-7516(96)00043-9

172

G. Chiet al. l i n t . J. Miner. Process. 49 (1997) 171-183

Au(CN) 2 C'-i " AulCN) 2 CN"

ON-//

~\

I "11~--~Depositgold ed

particle

"~"------~ ] ~ " ~

CN-

ZnlCN)4 2Fig. 1. Schematicof mechanismof gold precipitationon zinc [14].hampering gold production efficiency. It was decided to investigate the influence of some of the factors which affect the efficiency of precious metals precipitation from cyanide solutions by zinc dust under controlled conditions and with pure solutions. Although the Merrill-Crowe process is widely used in the precious metals industry, most of the published literature on the cementation of gold and silver prior to 1960 dealt largely with plant practice [2]. Detailed laboratory investigations have only been reported during the past two decades. Studies have been mostly confined to the kinetics of gold and silver cementation on rotating discs or cylinders [3-6]. In the late 1980s, comprehensive investigations were conducted by Parga et al. [2,7,8] to study the mechanisms of the cementation reactions. The cementation of gold and silver by zinc dust is an electrochemical process, proceeding by localized anodic and cathodic reactions. A schematic of this process is shown in Fig. 1. The main reactions for zinc dissolution are: Zn + 4 C N - + 1/202 + H20 = Zn(CN)~- + 20H 2Au(CN)2- + Zn = 2Au + Zn(CN)4 Zn + H20 + 2Au(CN)2- = 2Au + HZnO 2 + 3 H + + 4 C N (1) (2) (3)

To better understand the phenomena controlling zinc efficiency in these systems, the solubility of zinc in alkaline cyanide solution was established under various conditions in the absence of precious metals. The effects of cyanide, oxygen and lead nitrate concentrations and pH on zinc solubility were investigated.

2. Experimental methods and materialsIn order to simulate the Merrill-Crowe process, a laboratory procedure was established which allowed de-aeration of the solution to achieve a desired dissolved oxygen concentration. Zinc concentration was measured by Atomic Absorption Spectroscopy.

G. Chiet al. lint. J. Miner. Process. 49 (1997) 171-183

173

Solution

l

filter

sample bottle

[3

Fig. 2. Schematic of experimental apparatus. 1 = transfer tube for adding zinc; 2 = transfer tube for adding Pb(NO3):; 3 = oxygen meter probe; 4= tube for air escaping; 5= glass tube for solution withdrawal; 6 = magnetic stir bar; 7 = magnetic stirrer; 8 = gas dispersion tube. A11 of the experiments were conducted in a four-necked 500-ml glass reaction vessel. A dissolved oxygen probe, nitrogen dispersion tube, sampling device and zinc and lead nitrate addition tubes were placed into the reactor through the openings in the cover, as shown in Fig. 2. In all of the studies, purified nitrogen was passed through the solution via a dispersion tube prior to conducting the experiment. A YSI Model 58 dissolved oxygen meter was used to measure the concentration of dissolved oxygen in the solution. The test solution was prepared with reagent-grade NaCN. Reagent-grade NaOH was used to adjust the pH of the solution. Distilled water was used for all solutions. Before and after every batch experiment, the solution was titrated for cyanide concentration. In each test, 18 m g / l of Merrillite zinc (Pasco) was added with various amounts of Pb(NO3) 2. Solution was withdrawn by means of a peristaltic pump for analysis of zinc concentration when oxygen concentration did not decrease with time.

3. Thermodynamic considerations of the Zn--CN--H20 systemIn order to understand in more detail the dissolution of zinc in cyanide solution, it is useful to consider the thermodynamics of the system utilizing Eh-pH diagrams. These diagrams delineate the stability regions of various specious at equilibrium and, consequently, are useful in the study of the solubility of zinc in cyanide solution. It has been suggested that the following reactions occur when zinc dissolves in cyanide solution: Zn + 4 C N - = Z n ( C N ) ] - + 2e Zn + 3 O H - = HZnO2- + H 2 0 + 2e (4) (5)

These half reactions can be coupled with the water discharge half cell to produce hydrogen: Zn + 4 C N - + 2 H 2 0 = Z n ( C N ) ] - + 2 O H - + H 2 Zn + 2 H 2 0 = HZnO 2 + H + + H 2 (6) (7)

174

G. Chi et al./lnt. J. Miner. Process. 49 (1997) 171-183

All of these reactions are electrochemical processes and may be resolved into their appropriate half cells for the construction of pertinent Eh-pH diagrams. The essential half-cell reaction for zinc dissolution in aqueous solution is: Zn 2+ + 2e = Zn E = - 0 . 7 6 2 + 0.0295 log [Zn 2+ ]

(8)

Zinc is so strongly reducing that even at relatively high Zn 2+ concentrations, the potential lies well below that required to reduce water. In alkaline cyanide solutions, zinc forms numerous complexes by reaction with hydroxyl and cyanide ions. Of these, one, Zn(OH) 2, is sparingly soluble. The rest are soluble. The relevant equilibria involving metallic zinc are: Zn(CN)4z- + 2e = Zn + 4 C N E = - 1 . 2 5 + 0.0295 log [Zn(CN)4-] + 0.118pCN Zn(OH)2 + 2e = Zn + 2 O H E = - 0 . 4 1 9 - 0.059pH HZnO] + 3H + + 2e = Zn + 2H20 E = 0.054 - 0.0886 pH + 0.0295 log [HZnO 2 ] ZnO 2- + 2H20 + 2e = 4 O H - + Zn E = - 1.216 + 0.0295 log [ZnO~- ] + 0.118 pOH (12) (11)r _ _~ 1

(9)

(10)

In addition, hydrocyanic acid acts as a weak acid, and its hydrolysis can be represented as: H + + C N - = HCN pH + pCN = 9.4 - log[HCN] (13)

Eq. (13) indicates a considerable reduction in the cyanide activity when the pH falls below the value of 9.4. The total cyanide content, A, as referred to in practice, is defined by the following equation: A = [HCN] + [ C N - ] and Eq. (13) becomes: pH + pCN = 9.4 + log(1 + 10 pH-94)_ _

(14)

loga

(15)

By combining Eq. (9) and Eq. (15), one obtains: Zn(CN)4- + 2e = Zn + 4 C N E = - 1 . 2 5 + 0.0295 log[Zn(CN) 2- ] + 0 . 1 1 8 1 9 . 4 - l o g A + log(1 + 10 pH-9"4) - p H ] The Eh-pH diagrams describing these reactions are shown in Fig. 3a-c. The calculations are based on the thermodynamic data taken from references [9-11]. (16)

G. Chi et al. / Int. J. Miner. Process. 49 (1997) 171-183

175

(a)

1.0 0.0-I .0 -2.0

.............Zn ~

[" .....i

2+u,oG~-_-i~r ......

--1

] 11.... -1__~

iI I

>

.............

,

=,~=

0

2

4

6 pH

8

10

12

14

...............

;I

i

z.-

,

......

o.o . . . . . .

.-2.1

2.2o+~e~o.-~_.

.......

. .........

,,

z,c~4, / ~~, , , ~ ,/~/-~I

|

-I---t---~ ~ |,

---- ]=

J

~- . . . . . . . .

', ~/Co)

Zn 2 4 6 8 10 12

14%

pH

/(C) 1.0-

. _. .++ . .O 2 + H 2 0 + 4 e = 4 O H . ~"2

--I

'. . . . . . . . . . . . . . .'

I 1 ' 1l][--;--1 II o~' ', ~ I1~ ',~=/

0.0 . . . . . . . .-1.0 "-2.0 0

.~o+2o=~o,-,-+.~~2 4

..........

:.... . .......I

z"~cro,2n

z6 pH

8

10

12

14

Fig. 3. Eh-pH diagrams of the Zn-CN-H20 system at three cyanide concentrations,[Zn(II)]= 1.5 >