Anelectrochemicalinvestigationofthe …ammonia with oxygen pressures of up to 10atmospheres, found that both the amount and rate ofnickel dissolution were considerably lower than t~ose

An electrochemical investigation of thedissolution of copper, nickel, and copper-nickelalloys in ammonium carbonate solutions

SYNOPSIS

by M. J. NICOL*, Ph.D., B.Sc. (Hons.) (Wits.) (Visitor)

The dissolution of copper and nickel in ammonium carbonate solution was studied by techniques involving electrc-chemical anodic dissolution and ring-disc electrodes. The rate of oxidation of both copper and nickel by copper (11)was shown to be controlled mainly by diffusion to the surface. Variations in rate with changes in copper (11) concen-tration, agitation, pH, ammonia concentration, temperature, and copper (I) concentration were investigated andrationalized in terms of their effects on the electrochemical half-reactions involved. Evidence was found for thepassivation of nickel and nickel-copper alloys of more than 50 per cent nickel.

SAMEVATTING

Die oplossing van koper en nikkel in 'n ammoniumkarbonaatoplossing is bestudeer deur tegnieke waarby elektro-chemiese anodiese oplossing en ringskyfelektrodes betrokke was. Daar is getoon dat die oksidasietempo van sowelkoper en nikkel deur koper (11)hoofsaaklik deur diffusie na die oppervlak beheer word. Tempovariasies in ooreen-stemming met veranderinge in die koper (11)konsentrasie, agitasie, pH, ammoniakkonsentrasie, temperatuur enkoper (I) konsentrasie is ondersoek en gerasionaliseer in terme van hul uitwerking op die elektrochemiese half-reaksies daarby betrokke. Daar is tekens van die passivering van nikkel en nikkelkoperlegerings met meer as 50persent nikkel gevind.

INTRODUCTION

In the presence of oxygen, copperdissolves readily in ammoniacal solu-tion in accordance with the reaction

Cu+4N!f3+i02+H2O:?Cu(NH3)4Lj--20H-. .., (1)The Cu(NH3):+ formed can also

oxidize copper according to thereaction

Cu +Cu(NH3)~+:?2Cu(NH3h+, ...,... (2)

the copper (1) produced being rapidlyoxidized by oxygen to copper (H).Several studies of the overall processhave been reported in the literature,the most important of which arecritically reviewed below.

Lane and McDonaldl studied thekinetics of the dissolution of rotatingcopper cylinders in oxygenated am-monia solutions at normal tempera-tures and pressures. The rate ofproduction of copper (H) in solutionwas found to be zero-order at highammonia concentrations, althoughan approximately twofold decreasein rate was noted at a well-definedpoint during the reaction. Theyfound no evidence of autocatalysisby copper (H), the maximum con-centration of which never exceeded2 X 1O-2M during the experiments.Both apparent zero-order regionswere found to be affected by agita-tion, temperature, and ammonia

*National InstituteJohannesburg.

for Metallurgy,

concentration in a similar way. Theformation of basic films at thesurface due to local increases inpH as a result of reaction (1) wereinvoked to account for the observeddecreases in rate and dependenceof the rate on ammonia concentra-tion. However, no detailed mechan-ism of the reaction was suggested.

Lu and Graydon2 studied the dis-solution in both oxygenated andaerated ammonia solutions at atmo-spheric pressure with and withoutthe addition of copper (H) before thestart of the reaction. The maximumconcentration of copper (H) duringany run was 1O-2M. They foundthe reaction to be zero-order inoxygenated solutions, but apparentlyhalf-order in aerated solutions, al-though limited data at low concen-trations of copper (H) showed azero-order dependence on copper(H). The reactions were found tobe controlled by mass transportunder all conditions of the experi-ments. There are two main criti-cisms of their interpretation of theresults.(1) Autocatalytic behaviour will be

observed only above a minimumratio of copper (H) to oxygen,which can be deduced as follows.For diffusion-controlled reduc-tions of oxygen and copper (H),the relative contribution ofeach to the reaction will beproportional to nD2J 30 (see theAppendix), where n is the

JOURNAL OF THE SOUTH AFRICAN INSTITUTE OF MINING AND METALLURGY

number of electrons in thereduction of the oxidant, Dis the diffusion coefficient, and0 the concentration. Thus, thecontributions will be equal ata concentration ratio

[CU(H)]/[02] = 4Do2/DcU(II)

= 17(for DO2=2,6 X 10-5 cm2s-1

and DclI(II)=6 X 1O-6cm2s-1).Thus, in oxygenated solutionsat 25 QC ([02] = 1 X 10- 3M),autocatalysis will become ap-preciable only at concentrationsof copper above 2 X 1O-2M.The corresponding figure foraerated solutions is 4 X 10- 3M.Thus, the absence of an auto-catalytic effect in the case ofLu and Graydon's experimentsin oxygenated solutions isreadily explained. Furthermore,the half-order dependence inaerated solutions must be ap-parent only in that a continu-ously variable order from zeroat low copper (H) to one athigh copper (H) concentrationswould be expected.

(2) Their attempt to rationalizethis half-order dependence interms of a mechanism is invalidfor the following reasons. Theyassume that the slow step isthe removal of copper (I) fromthe surface and that this copper(I) is produced via the assumedrapid equilibrium (2) at thesurface.

JUNE 1975 291

EOMedium log10K V vs NHE

0 0,518

2M NH,NO3 5,93

2M NH,NO3 4,93

Derived from above -0,134

0 0,338

2M NH,NO3 4,27

2M NH,NO3 3,59

2M NH,NO3 3,00

2M NH,NO3 2,19


Derived from above 0,026

Derived from above 1,33

0 -0,232

2M NH,NO3 8,74


2M NH,NO3 9,50

These assumptions result in arate equation of the form

Rate = k[Cu(II)]l s,where [Cu(II)] s is the concentrationof copper (II) at the surface. Thefinal assumption required is thatthis concentration is the same asthat in the bulk. This, however, istrue only if the above equilibriumresults in a small perturbation ofthe bulk concentration at the inter-face. Thermodynamic data (to bepresented later) show that theequilibrium constant for reaction (2)under the conditions of their experi-ments is large (50M), so that theinterfacial concentration must neces-sarily be much lower than that inthe bulk. Halpern 3,4 and Habashi5have erroneously subscribed to thesame mechanism.

Halpern 3, 4, in a study of the

reaction at various oxygen pressuresfrom 1 to 8 atmospheres, foundthat the reaction could be describedin terms of two distinct rate-

determining processes. At lowoxygen pressures, the rate is limitedby diffusion of oxygen to the surface,while at high pressures and, par-ticularly, at low ammonia concentra-tions, the rate was found to beindependent of oxygen pressure andagitation, but approximately firstorder in ammonia concentration.The slow step in this region waspostulated in terms of the reactionof NH3 and NH, + with an oxidefilm on the surface. Habashi5 treatedHalpern's 3,4 data in terms of anelectrochemical mechanism, which,however, is qualitatively invalid inthat it ignores the potential de-pendence of the half-reactions andrequires continuously variable sepa-rate areas to be available for theanodic and cathodic processes.

A recent study of reaction (2)by Smith, Lower, and Hockings6has, to some extent, confirmedHalpern's 3, 4 results, i.e., a diffusion-controlled process at low concentra-

TABLE ITHERMODYNAMIC DATA

tions of copper (II) with decreasedrates, and the formation of a visibleoxide film at high concentrations.Their treatment of the data in theregion of no film formation is similarto that of Lu and Graydon2 in thatit tacitly, but incorrectly, assumesa small equilibrium constant forreaction (2). They also postulatethe formation of an oxide film(CU20 in particular), and the passiva-tion is attributed, rightly, to thedecrease in the concentration offree ammonia at the higher concen-trations of copper (II).

A process for the ammoniacalleaching of nickel metal in reducediron ore has been described by thestaff of Inc07. They found thatboth nickel and iron tended topassivate in oxygenated ammoniumcarbonate solutions, and frequentreactivation by hydrogen reductionwas necessary. Morioka and Shima-kage8, in a study of the dissolutionof nickel and copper powders in

Reaction

Cu+ +e~Cu

Cu++NH3~Cu(NH3)+

Cu(NH3)++NH3~Cu(NH3).+

Cu(NH3).++e:?Cu+2NH3

I

-1Temp.

°C

25

18

18

25Cu'+ +2e:?Cu

Cu'++NH3 :?Cu(NH3)Z+

Cu(NH3)Z++ NH3:?Cu(NH3).Z+

Cu(NH3).Z+ +NH3~Cu(NH3)3Z+

Cu(NH3)3z+ + NH3~Cu(NH3),z+

Cu(NH3)"++2e~Cu+4NH3

Cu(NH3),z+ +e~Cu(NH3).+ + 2NH3

Cu(NH3),,++Cuz2Cu(NH3).+

NiZ++2e~Ni

Niz+ +6NH3~Ni(NH3).Z+

Ni(NH3).'++2e~Ni+6NH3

25

25

25

25

25

30

25NH3+H+:?NH,+

JOURNAL OF THE SOUTH AFRICAN INSTITUTE OF MINING AND METALLURGY292 JUNE 1975

ammonia with oxygen pressures ofup to 10 atmospheres, found thatboth the amount and rate of nickeldissolution were considerably lowerthan t~ose for copper. Althoughthey dId not rationalize this be-haviour in terms of passivation ofthe nickel, the results tend tosupport this view.

Since, as will be demonstratedthese processes are indeed electro~chemical in nature, an electro-chemical approach was consideredto be the most useful in arrivingat a complete description of themechanisms of these reactions. Thispaper summarizes the results ofsuch an investigation.

Table I lists published (Martenand Sillen9) and derived thermo-dynamic data relevant to the leach-ing of copper and nickel in ammonia-cal solutions. These data illustrateseveral relevant points.(I) The greater stability of the

copper (I) ammines relative tothe ?opper (H) ammines (ap-proxImately an order of magni-tude difference in the K values)results in a favourable shift inthe Cu (H) + Cu :? 2Cu (I)equilibrium in ammoniacal solu-tions. However, since the actualvalue of this formal equilibriumconstant will depend on theconcentration of free ammoniait was thought that actualmeasurement of K at variouspH values would be of value inan assessment of any thermo-dynamic limitations to the cop-per-leaching reaction.

(2) The relatively high stability ofthe copper ammines means thatthey can compete with H + forthe free NHa and will evenpersist at pH values as low as 8.

(3) The difference in the Eo valuesof about 0,5 V between thosefor oxidation of nickel andreduction of copper (H) inammonia solutions means thatthere are no thermodynamicrestrictions to the leaching ofnickel with copper (H) as theoxidant.

Thermodynamic Measurements

The formal reduction potentialsof the Cu (H)-Cu (I)-Cu (0)couples were obtained by directmeasurement as follows. A solutionof 4 X 10- aM CUS04 in IM~NH4)2COa was reduced cathodicallym a thermostatted, well-stirred de-oxygenated cell. A large platinum-gauze cathode and an isolated anodewere used, and a constant currentof 5,0 mA was passed for a periodthat was calculated as being equiva-lent to half-reduction of thecopper (H) present. The potentialof a platinum plate electrode wasthen measured against a saturatedcalomel electrode in the solutionand taken as the formal reductionpotential of the Cu (H) - Cu (I)system at the particular pH value.The pH value of the solution wasthen recorded. Similar experimentswere carried out at different pHvalues, either concentrated sulphuricac~d or sodium hydroxide pelletsbemg used for the adjustment of pH.

Complete reduction of thecopper (H) solution resulted in acopper (I) solution, and the potentialof a copper-plated platinum plateelectrode versus a saturated calomelelectrode was used in the calcula-tion of the formal reduction poten-tial of the Cu (I) - Cu (0) couple.Coml(lete reduction ofthe copper (H)solutIOn was essential, and thefollowing criteria were employed:(I) persistence of a visible copper

plate on the cathode for at leasthalf an hour,

(2) a platinum potential of at least200 m V more cathodic thanthat for half-reduction (thiscorresponds to a reduction inthe ratio of copper (H) tocopper (I) of about 104), and

(3) a difference of less than 2 m Vin stirred and unstirred solu-tions, which would confirm thatany 'mixed potentials' due tothe reaction Cu + Cu (H) :?2Cu (I) are minimal.

Electrochemical Measurements

Rotating discs of the metals (A.R.grade) set in PTFE holders wereused in the study of the anodic

Three types of measurements were dissolution of copper and nickel.involved: thermodynamic, electro- The areas of the exposed faces werechemical, and chemical-dissolution . 0 32 d, an 0,39 cm2 respectively.


EXPERIMENTAL METHOD

Discs of copper-nickel alloys suppliedby Johnson, Matthey having anexposed area of 1,77 cm2 weresimilarly employed. Steady-statepotentiostatic and cyclic voltam-metric measurements were made ina three-electrode system by the useof a potentiostat and triangularwave generator assembled from solid-state operational amplifiers. Unlessstated otherwise, all measurementswere made at 22 :t O,loC, and allpotentials are quoted with respectto a saturated calomel referenceelectrode (+0,245 V versus normalhydrogen electrode).

Chemical-dissolution Measure-ments

The rates of the dissolution re-actions

Cu (H) + Cu ~ 2Cu (I) . . . (3)and 2Cu (U) + Ni ~ 2Cu (I) +

Ni (H)""""

(4)were studied when they were allowed

t~ occur on the copper-plated (ornIckel-plated) platinum disc of aring-disc electrode (dimensionsrl =0,48 cm, r2=0,51 cm, andra=0,56 cm). The copper (I) pro-duced by the reactions at the discwas monitored at the platinum ringby the reaction

Cu (I) ~ Cu (H) + e.The theoretical background for thisuse of a ring-disc electrode is out-lined in the Appendix. The ring wasconnected in a conventional three-electrode potentiostatic circuit andthe disc to a variable-voltage ~owersupply for plating of the disc. Ingeneral, the disc was plated from thesame solution that was used forstripping, the amount of copper

rem?v:ed in a 5-minute plate beingneglIgIble. On completion of theplate, the disc was open-circuitedand the ring current at a suitabl~potential was monitored as a func-tion of rotation speed. A study of theanodic oxidation of copper (I)showed that the process becomesdiffusion-controlled at potentialsanodic to 200 m V versus SCE. Inall experiments, the ring potentialwas therefore set at 400 m V.

.In experiments with nickel-plated

dIscs, a solution of O,IM NiS04 inIM (NH4)2COa at a pH value of 8was used for plating. Owing to thelow current efficiency, longer deposi-

JUNE 1975 293

tion times (up to 20 minutes) werenecessary.

The collection efficiency of thering was determined from copperanodic-dissolution experiments, carebeing taken to exclude oxygen fromthe system. A mean value of 33,2::!: 0,3 per cent for the efficiencywas obtained in the range 50 to 1000rev/min.

Unless stated otherwise, all experi-ments were carried out at 22 ::!:0,1 DCin a thermostatted MetrohmEA-880T cell fitted with reference(saturated calomel) and counterelectrodes for both ring and disc,and with a gas-inlet tube for de-oxygenation with high-purity nitro-gen. Solutions were made up bythe dissolution of A.R.-grade coppersulphate and A.R.-grade ammoniumcarbonate in triply distilled water.

RESULTS AND DISCUSSION

The three types of measurementsare again treated separately.

Thermodynamic Measurements

Results of the equilibrium poten-tial measurements are shown inFig. 1 as plots of the formal poten-tials (versus SCE) for the Cu (1)-

0,3

E7vvs

SCE

0,2

8,5

Cu (0) and Cu (II)-Cu (I) couplesas a function of pH. From the figure,it is apparent that K for the reactionCu (11) + Cu?2Cu (I) increases withpH, and calculation shows that itreaches a limiting value of about60 at pH values above 9,2. Theslope of the linear portion (below apH value of 9) of the Cu (I)-Cu (0)line of -120 m V per pH unit isconsistent with the formation of astrong Cu(NHa)2 + complex. A valueof 10,3 for log KIK2 is obtainedfrom the extrapolated value of Eo f

at a pH value of 8, on the assumptionthat ECu(I)/Cu(O) =0,518 V (versusNHE). This is in good agreementwith the value of 10,86 quoted inTable I.

Similarly, the slope of the Cu (11)-Cu (I) line of - 68 m V per pH unitconfirms the stability-constant data,which indicate increased stabiliza-tion of the copper (11) state relativeto copper (1) at higher pH values interms of increasing amounts of thetriammincs and tetrammines. Thevalue of -68 mV per pH unitsuggests that the predominant cop-per (11) species in this range of pHis the triammine (predicted - 60

9 9,5pH

Fig. I-Variation of the formal reduction potentials of the Cu (lI)jCu (I) andCu (l)jCu (0) couples with pH in I M (N H4).C03

294 JUNE 1975

m V) with small amounts of thetetrammine (predicted -120 m V).

Studies of Anodic Dissolution

CopperLinear sweep voltammetric curves

for the anodic dissolution of arotating copper disc in IM (NH4)2COasolutions of various pH values at asweep rate of 0,25 V min-l areshown in Fig. 2. The curves arenot compensated for iR drops andcan therefore be treated quantita-tivoly only at low currents. Theinflection at about - 50 m V in thecurves at the lower pH values isprobably associated with the forma-tion of copper (11).

A plot of the potential at constantcurrent (1,0 mA) versus the pHvalue is linear, with a slope of104 m V versus pH value, suggestingsecond-order dependence on OH-(or free NHa) concentration.

Partial passivation of the copperis observed at low concentrationsof free ammonia as shown by thecyclic voltammograms in Fig. 3. Thepeak at about -0,15 V can beassigned to the process

2Cu + H2O--+Cu20 + 2H + + 2e(E(pH 9)= -0,31 V versus SCE),

. . . . . . . . (5)and that at +0,040 V to the process

CU20+ H2O--+2CuO + 2H ++2e(E(pH 9)= -0,11 V versus SCE).

. . . . . . . . (6)The voltammogram (c) obtained inthe absence of ammonia also ex-hibits a shoulder at potentials anodicto -0,2 V and a cathodic peak at

- 0,32 V, which agrees well withthat predicted for the reverse ofreaction (5). In addition, a cathodictransient is evident in the region of-0,10 V during the cathodic sweepin voltammogram (b). The excellentagreement of this value with thatpredicted for the reverse of reaction(6) lends further support to theabove assignments. The absence ofcathodic peaks for reaction (5) inthe presence of ammonia suggeststhat the CU20 formed dissolvesrelatively rapidly in the solution. Acomparison of the peak heights forvoltammograms (a) and (b) indicatesthat the rate of the oxide-dissolutionreaction is approximately propor-tional to the square of the ammoniaconcentration.


pH1. 7,92. 8,63. 9,14. 9,65. 9}9

3

«E

I-~

ZWQ:Q:::>(J

-400 -300 -200 -100

POTENTIAL} mV vs SCE

Fig. 2-Linear sweep voltammograms (0,25 V min-1) for the anodic dissolution ofa rotating copper disc (250 revjmin) at various pH values in I M (N H.).COa

600

Ibl

<{

'"

+200-+-

POTENTIAL, mY vs ~CE

-200

Fig. 3-Cyclic voltammograms for a rotating copper disc (1000 revjmin) in IMNaCIO. containing (a) O,OIM (NH.).COa, (b) O,02M (NH.).COa at 0,3 V min-"

and (c) 0 (NH.).COa at 3 V min-1

JOURNAL OF THE SOUTH AFRICAN INSTITUTE OF MINING AND METALLURGY JUNE 1975 295

4

<E

I--~ZUJococ:;)<J

2

NickelVoltammograms for the anodic

dissolution of a rotating nickelelectrode in lM (NH4}zCO3 solutionat a pH value of 9,65 are shown inFig. 4. The curves exhibit a typicalpassivation peak at about - 300 m V.A comparison between these curvesand that obtained in the absence ofammonium ions, i.e., in lM NaClO.solution having a pH value of 10,is shown in Fig. 5. It is apparent thatthe formation of oxide or chemi-sorbed oxygen at -0,50 V in NaC1O4

(El for Ni - NiO at a pH value of 10

is -0,70 V) is suppressed in thepresence of ammonium carbonate.This is presumably due to adsorptionof NH.+ (or NH3) at the negativepotentials involved or to dissolutionof the oxide film in the ammoniasolution. However, the actual passiv-ating species is possibly similarbecause of the presence of a cathodicpeak in the region of - 800 to - 900m V in both cases.

As shown in Fig. 4, the passivelayer can be slowly removed if the

GOmin

-480 -320 -160

POTENTIAL, mY vs SCE

Fig, 4---Linear sweep voltammograms (0,25 V min-1) for the anodic dissolution of arotating nickel disc (1000 revjmin) in I M (N H.).CO3,

(a) Held open-circuit for various times before sweep(b) Held at -1030 mV for 30 s before sweep(c) Held at -700 mV for 30 s before sweep

400 1000

«~

'"zUJet:et::)U

200 500-~

-600 -400 -200

POTENTIAL,.mV vs SCE

Fig, 5-Cyclic voltammograms (6V min-1) for a rotating nickel disc (1000 revjmin)in (a) IM NaCIO., pH 10 and (b) IM (NH.).CO3

296 JUNE 1975 JOURNAL OF THE SOUTH AFRICAN INSTITUTE OF MINING AND METALLURGY

disc is simply left to stand in de-oxygenated (NH4)2CO3 solutionf'!, in-creasing amounts being removed in10, 30, and 60 minutes. Reduction ofthe passive layer is also possibleas shown by the curve in which thecathodic limit was increased to-1030 mY, i.e., into the region of

hydrogen evolution. These observa-tions are in accord with the observa-tions made by Inco 7 in the leachingof nickel that activation in ammonia-cal solutions in the absence ofair was necessary prior to dissolu-tion.Copper-Nickel Alloys

Steady - state potentiostaticcurrent- potential curves for theanodic dissolution of various copper-nickel alloys in IM (NH4)2CO3 solu-tion are shown in Fig. 6. It isapparent from the curves thatpassivation is observed only at lowpotentials for the alloys having acopper content of 50 per cent orless, the 70 per cent alloy givingrise to a curve similar to that forpure copper. Analysis of the pro-ducts of dissolution shows that nopreferential solution of one of thecomponents occurs. Similar resultswere obtained by UhliglO for theanodic dissolution of copper-nickelalloys in sulphuric acid solutions.Passive properties were found tobegin at a critical copper contentof between 60 and 65 per cent,which coincided with an unfilledd-band of electronic-energy levelsas measured by magnetic saturationmoment or electronic specific heat.Uhlig proposed the theory that thepassive film on copper-nickel alloysis made up largely of chemisorbedoxygen, the formation of which isfavoured by surface d-electronicvacancies. When such vacancies(0,6 per atom for nickel) are filled byvalence electrons of the alloyed atom(1 unpaired electron per atom forcopper), passive properties are nolonger observed, presumably be-cause oxygen is no longer chemi-sorbed.

Studies of Chemical Dissolution

CopperPreliminary experiments sug-

gested that the rate of dissolution ofcopper from the disc of the ring-disc electrode as a result of reaction(5) is controlled by mass transport,

i.e., the rate (as measured in terms ofa ring current, i R) is proportionalto the copper (11) concentration anddependent on the rotation speed.Application to a rotating disc of theLevich equationll for convectivediffusion results in a relation be-tween the ring current and the

product [Cu(lI)]wi, where w is thedisc rotation speed as derived inthe Appendix. A log-log plot of thisrelation for data in the range3 X 10- 3M to 8 X 1O-2M copper (11)is shown in Fig. 7. The results overa more-extensive range (2 X 1O-2Mto 0,6M) are summarized in Fig. 8.

!z-UJ

'"'"::>(J

10

<{E 1

0,1"IoCu

1. 1002. 703. 504. 305. 106. 0

-400 -300 -200 -100

POTENTIAL. mV vs SCE

Fig. 6-Steady-state current-potential curves for the anodic dissolution of rotatingcopper-nickel alloy discs (1000 rev/min) in IM (NH.).CO.

10

<{E

....-zUJ

'"'"i31

'"z;;:

0,01~

0) ~

[CuIm] wi I MREV! MIN!)

Fig. 7-Variation of the ring current due to copper dissolution with disc rotationspeed and copper (11) concentration in IM (N H.).CO.


10

1....'z

'"'"'"aC>za:

4 15E 2,01-'ZUJ0::0:::)U

C>za::

101,5

The deviations from linearity athigh copper (11) concentrations aremore apparent in this case. Asoutlined in the Appendix, the slopeof the lines in Figs. 7 and 8 shouldbe given by 19,5D22j 3, where D2 isthe diffusion coefficient of copper (11)

.- calculation yields D2 values of5,7 X 1O-6cm2s-1 and 5,0 X 1O-6cm2s-1respectively. An analysis of thediffusion-limited current for the re-duction of copper (11) to copper (I)at t,he rotating disc measured fromthe curves in Fig. 13, yields avalue of 6,9 X 1O-6cm2s-1 for D2,which is in reasonable agreementwith the above values.

The effect of pH on the dissolutionrate in IM (NH4hC03 containing0,02M copper (11) is shown in Fig.9. The pH value was adjusted by theaddition of sulphuric acid or sodiumhydroxide. Below a pH value of9,3, pH has little effect on the rate,which decreases with increasing pHabove 9,3. A similar decrease wasnoted in a 0,2M copper (11) solution,and, in addition, the rate decreasedat pH values below about 9,3. Theaddition of 1M (NH4)2S04 to thesolution had the effect of increasingthe rate at low pH values (below8,5) and decreasing it above 8,5.The effects on the rate of variationsin the ammonium carbonate con-centration, as demonstrated by thecurves in Fig. 10, parallel those ob-served as a result of variations inpH, which are therefore simply areflection of variations in free NH3concentration. Plots of ring currentversus rotation speed for severalpH values in 0,20M copper (11) areshown in Fig. H. The differencebetween the measured rates ofdissolution and that for a processcontrolled by mass transport (brokenline) are seen to be greatest at thelow pH values.

Reasons for the above effects be-come apparent when the potentialof the copper disc is measuredduring dissolution. The effect of thecopper (11) concentration on thispotential is shown in Fig. 12 as aplot of log [Cu (11)] versus potential.The resulting curve parallels thatfor the anodic dissolution (alsoshown), thus confirming that thechemical dissolution is first orderin copper (11) over the range 1O-4Mto 10- 3M. Dissolution rates in 10- 3M

298 JUNE 1975

and 5 X 10- 3M solutions, as meas-ured by the ring current, are shownin Fig. 12. These points fall closeto the anodic current-potentialcurve, thus confirming the electro-chemical nature of the process. Alsoshown are the potentials observedin oxygenated-saturated and air-

100

0,1

saturated solutions. It is apparentthat the rate in oxygenated solution(1O-3M °2) is equivalent to that in5 X 10- 3M copper (11) solution. Since

the ratio of Do2 to DOu(I!) is about

2,6, this implies that there are onlytwo electrons involved in the reduc-

""

[eu IIIU wr IM R£yt Mu,it)

1,0 10

Fig. 8-Variation of the ring current due to copper dissolution with disc rotationspeed and copper (11) concentration over a wider range in IM (NH.).COa

ex 204>2,5

8

c. In}

0 O,ZM

. O.1M~IM"zSO,

4>O,OZM

9 10

pH value

Fig. 9-Variation of ring current due to copper dissolution with pH at 250 revjminin IM (N H.).COa


,/

",/

/"

",<>:E

!Zwox:!510u

'"zii .../m;".. '00'x ".

///O'..x

./x-5

0,5 1,0

RNH,lzCO,],M

1,5 2,0

Fig. IO---Variation of the ring current due to copper dissolution with {N H.).CO3concentration at a pH value of9

30

//

/ x

// /»~J

,/ ~

,/// ~"""

/~/.~," V'

x, .(-2J3. A '--

,,/ ...~)~.! .-----.

... ..... ...... 1,1.. ..'

<>:E,...' 20zW0'.ox:::>u

'"""0::

10

20

W\ REV \ MIN-\

3010

Fig. I I-Variation of the ring current due to copper dissolution with disc rotationspeed in IM (N H.).CO3 containing O,2M Cu(II). Values in parentheses are the

potentials of the copper disc measured during dissolution

tion of each O2 molecule (i.e.,peroxide, and not water, is theinitial product).

For the chemical-dissolution re-action to be mass-transport con-trolled, these potentials must fallin the diffusion-controlled region forthe reduction of copper (II) tocopper (1). Fig. 13 shows linearsweep voltammograms for this reduc-

tion on a platinum disc at variouspH values. The second wave at about

- 0,5 V corresponds to reduction tothe metal. Several selected valuesof the potentials measured duringdissolution are indicated on thecurves. These potentials were ob-tained during the experiments, theresults of which are given in Fig. 11,in which the corresponding points


have been indicated. It is apparentthat both decreasing pH and in-creasing agitation result in mixedpotentials that are below (less cath-odic than) those required for thediffusion-controlled reduction of cop-per (II), and hence a dissolutioncontrolled by mass transport. Itmust be pointed out that noquantitative estimates of theseeffects can be made on the basis ofthe available data, which were ob-tained on a platinum electrode atcopper (II) concentrations consider-ably lower than that used for themeasurements of mixed potentials.However, the general conclusion isprobably valid, i.e., that the differ-ence in the formal reduction poten-tials of the two couples involved(100 mV at pH values greater than9 and 50 mV at a pH value of 8)is not sufficient to ensure that themixed potential will always fall inthe potential region where thereduction of copper (II) is diffusioncontrolled.

The effect of temperature in therange 20 to 50°0 on the rate ofdissolution in 0,08M copper (II) wasinvestigated. Activation energies of24,7 kJ mol-l at 1000 rev/min,19,3 kJ mol-1 at 250 rev/min, and17,2 kJ mol-l at 50 rev/min wereobtained. The increase with increas-ing agitation is in accord with theincreasing departure of the ratefrom mass-transport control withincreasing rotation speed.

The effect of added copper (1) onthe dissolution rate was determinedfrom the addition of known amountsof cuprous chloride to the leachingsolution. The concentration of cop-per (1) was determined in situ atthe end of the run by amperometrictitration with standard dichromlttesolution, zero ring current beingused to indicate the end-point. Aplot of [Ou (II)]w i versus i R fora series of experiments at approxi-mately 0,04M copper (II) with theaddition of increasing amounts ofcopper (1) is shown in Fig. 14. Itcan be seen that, except for smalldecreases in i R at the highestconcentration of copper (1) - aratio of copper (1) to copper (II) of7 to 5 - there is a negligible effecton the dissolution rate under theseconditions.

As shown in the Appendix, the

JUNE 1975 299

10CD

--------------------------------------

......

')'Eu

«E

>-.....iiim~O,::,

:=:z-W :>0::80::",:::>0u-. CD

1 -------------------------.

)xAIR

/.011,

~

!:J DISSOLUTION RATES CALCULATED

FROM RING CURRENTS

X MIXED PD,EN,IALS IN

IM {NH~12C02 SA""ATEDWI,H AIR OR OXYGEN

I

-300-400POTENTIAL,mY vs SCE

Fig. 12-Lower: Steady-state current-potential curve for the anodic dissolution ofa rotating copper disc (500 revjmin) in IM (NH.),COa

Upper: Variation of the mixed potential of the disc with Cu (11) concen-tration in I M (N H.),COa

dissolution rate (measured by i R)is proportional .to the copper (II)concentration if K is large, but isproportional to the square root ofthe concentration if K is small.Data obtained in 1M (NH4)2S04having a pH value of 6,6 where Kis certainly less than 1, are shownin Fig. 15, in which the better fitto a root dependence is confirmationof the theory. These results rein-force the remarks made in theintroduction about the source ofthe apparent square-root dependenceof the rate on the copper (II)concentration as observed by someworkers, i.e., such a dependence

300 JUNE 1975

can be obtained only under condi-tions in which K is small.Nickel

Preliminary experiments showedthat the nickel passivated at rela-tively low copper (II) concentrations,as would be predicted from theanodic-dissolution characteristicsoutlined previously. Plots of[Cu (II)]wt versus i R are shown inFig. 16 for the pre-passive region.The excellent linearity and slope,which yield an estimate of5,5 X 1O-6cm2s-1 for the diffusioncoefficient of copper (II) in reason-able agreement with the valuesdetermined from the copper leaching

results, confirm the diffusion-controlled nature of the reaction.It was noted that passivationoccurred at a copper (II) concentra-tion of 6,0 X 10- 3M in a 1M(NH4)2CO3 solution adjusted to apH value of 10 with sodium hydrox-ide, and at a concentration of1,6 X 1O-3M in aIM (NH4)2CO3 solu-tion at a pH value of 9. In the pre-passive region, however, all threesolutions gave the same dissolutionrates. It was found that a passivateddisc could be re-activated if left tostand in a deoxygenated (NH4)2CO3solution for approximately an hour,after which it exhibited dissolutionbehaviour similar to its pre-passivecharacteristics. On passivation, thepotential of the nickel disc increasedfrom about -400 mV to approxi-mately -200 mY, i.e., into thepassive region as found in the anodicdissolution studies (Fig. 4).

CONCLUSIONS

This study has conclusively de-monstrated that the dissolution ofboth copper and nickel in ammonia-cal solution containing copper (II)occurs electrochemically as reactionsin which the anodic dissolution ofthe metal is coupled to cathodicreduction of copper (II). In theregions of active dissolution, therate of dissolution has been shownto be determined mainly by masstransport of copper (II) to the metalsurface, although deviation fromsuch behaviour under conditions ofhigh copper (II) concentrations, lowpH, and increased agitation in thecase of copper have been observed.Similar observations have been madeby other investigators and, as de-monstrated in this work, are to beexpected as a result of the smalldifference in the equilibrium poten-tials of the half-reactions underthese conditions. No such limitationis expected in the case of nickeldissolution, although a practicaldemonstration of this cannot bemade because passivation of thenickel surface occurs at relativelylow copper (II) concentrations. Thispassivation in the case of nickel,and also nickel-copper alloys ofmore than about 50 per cent nickel,is an obvious disadvantage in anammoniacal leaching procedure fornickel, unless, as was done by


POTENTIAL, mV vs SCE

-600 -400mV (SCE)

-....zUJ

'"'"=>400 u

-200

200

600

Fig. 13-Linear sweep voltammograms (0,25 V min-1) forthe reduction of Cu (11) on a rotating platinum disc (250 rev/min)inIM(NH.).COacontaining(l) 1,82xlo-a M Cu(lI) atpH 8,6 (2) 1,88 x 10-aM Cu (11) at pH 8,8 (3) 1,92 x 10-aM Cu(11) at pH 9,2. Points (X) correspond to those indicated in

Fig. II

150

100

iR)JA

50

102:;

205,0

30 "104 !i:u (IJ)]1

1,5 x1OI[eu IU)J~

Fig. IS-Variation of the ring current due to copperdissolution with Cu (11) concentration in I M (N H.).SO. pH

6,6 at 1000 rev/min

8

4:=>.

4:E

I.

A.)

I "i/

'"9./1

'"il

c"III

"0

... 0,""

'"0,""

'ZUJ~ 4

=>u

<!)zi:<

°°

'"

...

1,50,5 1,0

@u(IIU W\ (M REV\ MIN\)

Fig. 14--Variation of the ring current due to copperdissolution with disc rotation speed and Cu (11) concentrat-

ions in IM (N H.).COa

100

4:

"'ZUJ'"'"=>u<!)~ 10

'"

10,001 0,01

[cum)] W! (M REV! MIN1)

. )-:I

:I

:IIIIIIIIII::III,IIII

PASSIVATION !IIIII

iIIIIIIIIII

:i:IIII

V

"0,1

Fig. I6--Variation of the ring current due to nickeldissolution with disc rotation speed and Cu (11) concentra-

tion in I M (N H.).CO.


in this work,for frequent

Inco 7 and confirmedprovision is madeactivation.

Passivation of copper was notobserved during chemical dissolutionunder the conditions employed dur-ing this investigation. However, asSmith et al.6 and Halpern 3,4 haveshown, passivation can occur if theratio of copper (II) to ammonia ishigh. Anodic dissolution experi-ments under such conditions haverevealed the existence of two passiv-ating oxides that have been electro-chemically identified as CU20 andCuO. The latter, however, cannot beformed if copper (II) is the oxidant.The existence of an optimum pHvalue in the region of 9 to 9,5 forthe rate of copper dissolution andthe decreased rates with high am-monia concentrations have beeninterpreted in terms of their effectson the rates of the half-reactionsinvolved.

The absence of any appreciableeffect of copper (I) on the rate hasan important bearing on the extentto which the copper (II) level canbe lowered during leaching. Thisobservation agrees with that madeby Smith et al.6, who found a small(about 20 per cent) decrease in ratewhen copper (I) was added inamounts equivalent to the cop-per (II).

Finally, several of the dis-crepancies in the literature regard-ing the mechanism of the dissolutionreaction have been clarified. Inparticular, conditions under whichthe dependence of the rate on thecopper (II) concentration willbecome one-half order have beenpredicted and confirmed experi-mentally in ammonium sulphatesolutions at low pH values.

APPENDIX

RING - DISC THEORY FORCHEMICAL STRIPPING

REACTIONS

Consider the chemical stripping ofcopper at a rotating disc by a

302 JUNE 1975

solution of copper (II).

KCu + Cu(II) :?: 2Cu(I) where

K = [Cu(I)]2/[Cu(II)],

and assume that this equilibriumis maintained on the surface of thedisc, Le., that the process is transportlimited.

At the disc surface,flux of Cu (II) to surface

= - t flux of Cu (I) fromsurface. (8)

The Levichll equation for a rotat-ing disc can be written as

flux =0,62D2/ 3y-l/6 wl/2{concentration in bulk -concentration atsurface}, (9)

where D is the diffusion coefficient, y

the kinematic viscosity of the solu-tion, and w the rotation speed inrad S-I.

Substitution of (9) in (8) yieldsD22/3 ([Cu(II)] 0 - [Cu(II)] s}

=tDI2/3 [Cu(I)]s=O, . . . (10)where subscripts 8 and 0 denotesurface and bulk quantities respect-ively, and subscripts 1 and 2 referto Cu (I) and Cu (II) respectively,and [Cu (I)] 0=0.

Now if Cu (I) is detected at thering by the reaction

Cu(I)~Cu(II)+e,then one can write for the collectionefficiency, N, N =i RjF A(flux ofCu (I) produced at the disc), whereiRis the ring current and A is thedisc area;

i.e., N=iRj{(FA)0,62DI2/3y-l/6wlJ2[Cu(I)] s},

or Cu(I)s=oiRjwt,where 0= Ij(0,62F ADI2/3y-l/6N).Substitution in (10) yields, on

rearrangement,Cu(II)s =Cu(II) 0

-DI2J 30i Rj2wlJ2D22/ 3,and substitution in (7) yields

wt [Cu(II)]0=DI2/30+ oiR

~R 2D22/3 K wl/i.

This equation has two limitingconditions.(i) K very large, in which instance

(7)

. 2D22/3wl/2 [Cu(II)] 0~R-

D 2/30 '1

i.e., in this case, [Cu(II)]s=Oand i R will be proportional to theproduct wt[Cu(II)] o.(ii) K very small, in which instance

. Ktwt[Cu(II)]ot~R- ot

i.e., in this case, [Cu(II)] s= [Cu(II)] 0and iRis proportional to the productwt[Cu(II)]ot.

ACKNOWLEDGEMENT

This paper is published by per-mission of the Director General,National Institute for Metallurgy.

REFERENCES

1. LANE, R. W., and McDoNALD, H. J.Kinetics of the reaction betweencopper and aqueous ammonia. J.Am. Chem. Soc., vo!. 68. 1946.pp. 1699-1704.

2. Lu, B. C. Y., and GRAYDON,W. F.Rates of copper dissolution inaqueous ammonium hydroxide solu-tions. J. Am. Chem. Soc., vo!. 77.1955. pp. 6136-6139.

3. HALPERN, J. The kinetics of thedissolution of copper in aqueousammonia. J. Electrochem. Soc., vo!.lOO. 1953. pp. 421-428.

4. HALPERN, J. Some aspects of thephysical chemistry of hydrometal-lurgy. J. Metals, 1957. pp. 280-289.

5. HABASHl, F. Kinetics and mechan-ism of copper dissolution in aqueousammonia. Ber. Bunsengesellschajt,vo!. 67. 1963. pp. 402-406.

6. SMITH, K. A., LowER, G. W., andHOCKINGS, W. A. Kinetics of leach-ing metallic copper in aqueouscupric ammonium carbonate solu-tions. PROC. INT. SYMP. ONHYDROMETALLURGY, Chicago,1973. pp. 265-292.

7. INco. Development of the Inco ironore recovery process. Can. Min.Metall. Bull., vo!. 59. 1956. pp.337-343.

8. MORlOKA, S., and SHIMAKAGE, K.Studies on the ammonia pressureleaching of metallic nickel andcopper powders. Trans. Japan Inst.Metals, vo!. 12. 1971. pp. 197-205.

9. MARTELL, A. E., and SILLEN, L. G.Stability constants. London, Chemi-cal Society, Special Publication no.17. 1964.

10. DHUG, H. H. Effect of electronconfiguration on the passive pro-perties of Cu-Ni alloys. Elechtro-chim. Acta, vo!. 16. 1971. pp.1939-1947.

11. LEVICH, V. G. Physicochemicalhydrodynamics. Engelwood Cliffs(New Jersey), Prentice-Hall, 1962.


Documents

Anelectrochemicalinvestigationofthe …ammonia with oxygen pressures of up to 10atmospheres, found that both the amount and rate ofnickel dissolution were considerably lower than t~ose