A calculation method for determining equilibria in metal-ammonia-water systems

hydrometallurgy ELSEVIER

Hydrometallurgy 38 (1995) 15-37

A calculation method for determining equilibria in metal-ammonia-water systems Shaolin Zhong ‘, Malcolm T. Hepworth b aColeraine Minerals Research Laboratory, Natural Resources Research Institute, Universiry of Minnesota, Duluth, Coleraine, MN 55722, USA b Department of Civil Engineering, University of Minnesota, Minneapolis, MN 55455, USA

Received 14 October 1992; accepted 26 October 1994

Abstract Based on mass and charge balances, a simple and convenient calculation method is proposed, which gives equilibrium relations between a solid phase and the associated aqueous phase in terms of total concentrations rather than the activities of the dissolved species. This technique is illustrated with examples in the metal systems: Ni-, Cu- and Mg-ammonia-water. For a given aqueous system, equilibrium diagrams, potential-pH diagrams, solubility curves and voltage-concentration ratio curves can be calculated as a function of ammonia addition. This analysis has practical value in interpreting the process conditions required to conduct hydrometallurgical processes.

1. Introduction The thermodynamic basis of complex hydrometallurgical extraction operations has been analyzed in terms of predominance diagrams, such as Eh-pH (or Pourbaix) and log [Me]pH [ 1,2]. However, the practical application of predominance diagrams is limited because of the differences between actual operating conditions and theoretical assumptions. For example, the stability relations in the metal-ammonia-water system are described by maintaining the activities of total dissolved metal and total free ammonia constant [ 3-61. This assumption does not take into account the presence of ammonia complexed in amines, a serious drawback to quantitative analysis of practical hydrometallurgical systems. An alternative for generating Eh-pH diagrams for metal-ammonia-water system has been proposed [ 71. By maintaining total dissolved metal and total ammonia constant, all equilibria are taken into consideration simultaneously. The diagrams then show stable domains for the solution and solids. However, one shortcoming still exists: in actual hydrometallurgical processes ammonia rather than a non-ammoniacal acid or base is generally added to adjust 0304-386X/95/$09.50

0 1995 Elsevier Science B.V. All rights reserved

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16

S. Zhong, M. T. Hepworth / Hydrometallurgy 38 (1995) 15-37

solution pH. As a result, calculations must be made for variable ammonia additions rather than for a fixed total ammonia content. In this paper, an attempt is made to introduce a simple and convenient calculation method, Based on both mass balance and charge balance equations, calculations are made to express equilibrium relations between solid phases and associated aqueous phases in terms of total concentrations rather than the activities of individual dissolved species. Since these equations differ from one system to another, the calculation procedure is also a little different. The detailed technique is described for examples for three metal (Cu, Mg and Ni)-ammonia-water systems separately. When an initial aqueous solution system is specified, along with the progressive ammonia additions, the corresponding equilibrium diagrams: potentialpH diagrams for the copper-ammonia-water system, solubility curves for the magnesiumammonia-water system, and voltage-concentration ratio curves for the nickel-ammoniawater system, can be calculated. All of the calculations can be finished on popular spreadsheet software (e.g., Lotus 123, Quattro Pro or Excel) and mathematic software (e.g., Mathcad). Since the pH in a metal-ammonia-water system is normally greater than 3, the concentration of HSO; is very low. At the same time, in the presence of free ammonia, the concentration of free metal ion is also low. Therefore, the concentration of ion pairs is expected to be very low. In our calculations, the species HSO; and ion pairs are ignored. This assumption is required by the simple calculation method used in this paper and is also verified by the calculation using the concentration constants of HSO, and ion pairs and the calculated concentrations.

2. Equilibrium data For each system, the mathematical consequences are expressed in terms of equilibrium constants, a charge balance equation and a set of mass balance equations. Each thermodynamic equilibrium constant can be rigorously written in terms of activity, and the mass and charge balance equations can be rigorously written in terms of concentration. In dilute solutions, activity coefficients depend only on the ionic strength of the solution and the charge on the ion. The theoretical prediction of activity coefficient is fairly accurate at ion concentrations below 0.1 mol/l. In more concentrated solutions the activity coefficients become dependent upon the nature of the ions in solution and a simple formulation is not feasible. The validity of the original and of the different extended forms of the DubyeHuckel equation is fairly limited. Butler [ 81 indicated that, at concentrations above 1 M, the errors involved in estimating activity coefficients are usually of the order of magnitude of the activity coefficients themselves. Therefore, it is just as accurate, and also much simpler, in this situation to assume that the concentration constant (usually measured in 14 M NaC104, if available) is independent of ionic strength. Results obtained with this approximation are usually not in error by more than a factor of 2, if all equilibria have been taken into account. In hydrometallurgical systems, when a concentrated solution is considered, it is preferable to use concentration constants rather than activity constants. In our calculations the concentration constants at ionic strengths of 2.0 mol/l are assumed. The concentration constants

S. Zhong, M.T. Hepworth / Hydrotnetallurgy 38 (I 995) 15-37

11

of metal ions and their complexes measured at constant ionic strength are extracted from the literature [ 1,9-151. The standard electrode potential is used for the calculation of solution potentials, and the activity coefficients for single ions are calculated from the mean activity coefficient using the mean salt method [ 2,161. High temperature equilibrium constants can be extrapolated from low temperature values using the Correspondence Principle [ 17,181. These thermodynamic data are presented for copper, magnesium and nickelammonia-water systems. 2.1. Analysis of the nickel-ammonia-water system In the following analysis of the nickel-ammonia-water system, the primary equilibria are represented in Table 1. In the Shenitt Gordon process a Fe-C+Ni-Cu sulfide ore is leached with ammonia at about 105°C and under an air pressure of about 8 atm. This brings cobalt, nickel and copper into solution and iron is precipitated as a hydroxide. Sulfur in the ore is mostly oxidized to sulfate but some of it goes to thiosulfate and polythionate ions. After separation of the undissolved solids, the pregnant solution is boiled at atmospheric pressure to expel the excess ammonia. This causes the metastable thiosulfate and polythionate ions to decompose and react with copper to precipitate copper sulfide, which is filtered off and shipped to a copper smelter. The remaining solution contains nickel and cobalt ammonia complexes in a ratio of approximately 40: 1. The solution is first treated with air at a pressure of 50 atm at about 200°C to oxidize all the thiosulfate and is then reduced with hydrogen to give metallic nickel. The hydrogen reduction is carried to the point where most of the nickel is precipitated but without affecting the cobalt in solution. In a typical reduction procedure the reduction feed solution, analyzing 46 g/l Ni and 350 g/l (NH&SO, with enough free ammonia to give an NH,:Ni molar ratio of 2, is delivered to an autoclave to react with hydrogen at a total pressure of 30 atm and a temperature of 204°C [ 19,201. Table 1 Thermodynamic data for Ni*+-NHS-H,O [ 1,10,12,15] Reaction

2H+ +2e=H2 H,O=H++OHNH3+H+=NH: Ni*+ +2e=Ni Ni*+ +NH,=Ni(NH, )*+ NiZf+2NHs=Ni(NH3)$+ Ni2++3NH3=Ni(NH3)$+ NiZ++4NHs=N~(NH3)4 2+ NiZ++5NH3=Ni(NH,)gf NiZ++6NHs=Ni(NH3)g+

Property

Eh” Gv KN

Eh’ PI P2 P3 P4 Ps P6

Temperature 25°C

100°C

200°C

0

0

0

lo-

13.96

1()9._w

- 0.236 lo* l9

1p3 ,06.X l(y.95 108.70 108.73

lo-

12.23

lo’.@ - 0.226 ,01.92 103.50 104.51 ,05.08 105.26 104.*

lo-“20 1

(p.0

- 0.216 100.52 101.27

1p3 1oL.23 1(pm 10-m

Eh” = the standard reduction potential for the reaction indicated with respect to the standard hydrogen electrode; K, = the concentration constant for the dissociation of water; KN= the concentration constant for the formation of ammonium ion; and the p values are the concentration constants for the association reactions indicated for the progressive amines of nickel.

18

S. Zhong, MT. Hepworth/Hydrometallurgy

38 (1995) 15-37

The reduction of nickel by hydrogen in an autoclave can be described by one of equations; Nizf+H2=Ni+2H+ either and H++NH,=NHz or Ni(NH,)z+ +H, = Ni + 2NHz + (X - 2)NH3. When all equilibria in solution are considered there is no difference between the calculated driving force of hydrogen reduction, no matter which equation is used. This driving force can be expressed as a combination of two half-cell reactions: Ni*+ + 2e = Ni, and 2H + + 2e = H,, for which the reduction potentials can be expressed as: EhNiz+/Ni = Ehsiz+ /Ni +-

Eh,+ ,nz = Eho,+ ,uz -

2.303RT 2F

10&WNi2 +

H _ 2.303RT ypgP,,,

2.303RT F

Combining these half cell reactions, reaction to be expressed as follows:

= EhONiZ+ /Ni +-

2.303RT 2F

enables the driving force of the overall reduction

lOga,iz+ - Ehi+ /HZ+

2.303RT 2.303RT F pH + ~lo@n,

(1)

Increases in nickel ion concentration, pH and hydrogen pressure favor nickel reduction. However, in order to avoid hydrolysis of nickel ions at high pH, ammonia is added, to serve the double purpose of complexing the nickel and raising the solution pH. High pH values promote reduction but the stable complexes of nickel act to keep nickel in solution. The most favorable concentration ratio of [NH,] / [Ni2+] must be found to maximize nickel reduction by hydrogen. For the purpose of illustration of the calculation technique, an assumption is made that the initial concentration of NiS04 is 1 .O M, and that of (NH,) 2S04 is ‘m’ M, where m is a variable. The mass balance for Ni is given by: [Ni2’],=

[Ni2+] + 2 [Ni(NH,)?+]

= 1

j=l

[Ni2’]

{ 1+ kPj[NH,]‘j

(2)

= 1

j=l

The mass balance for NH3 is given by: [NH31T= [NH,] + [NH,+] + &[Ni(NH,)_j?+]

=2m+

[NH3jadded

j=1

[NH,] + [H+] [NH,]&+

[Ni2+] tj&[NH,lj=2m+ j=l

The charge balance is given by:

[NH,].,,

(3)

S. Zhong, MT. Hepworth/Hydrometallurgy

38 (1995) IS-37

19

0.10

2 W 4

0.08

-

--

m=lY m=2M

Fig. 1. Ni-NH,-Ha0 system: the effect of ammonia addition on chemical driving force as a function of initial (NH,),SO, concentration. Conditions: 25°C; Pm= 1 atm; 1 M NiSO,.

2[Ni*+] +2 &Ni(NH,)j+]

+ [H+] + [NH:] = [OH-] +2[SOz-]

j=l

2,Ni2’]11+~P,[NH,I’)+[H+]+[HC][NH,]K~= j=l

Kw -+2(1+m)

W+l

combining with Eq. (2) : (l+K,[NH,]}[H+]*-2m[H+]

-K,=O

(4)

The results’, calculated for two hydrogen over-pressures ( 1 and 30 atm, respectively) are shown in Figs. 1 and 2. These diagrams show the most favorable ratios of [NH,] / [ Ni*+ ] for nickel reduction. With ammonia addition, the raised pH of the solution increases the chemical driving force. After AE reaches a maximum, the driving force drops because the added ammonia forms stable complexes with nickel to keep nickel in solution, although the solution pH is still increased. The driving force of AE for the reduction reaction reaches a maximum value when the NHJNi ratio is maintained between 2 and 2.5. These results ’The calculation procedure is as follows: ( 1) For a fixed initial concentration of ( NH,+)$04 and pressure of hydrogen, a series of values from 1Om6 to 1 M is assigned to the concentration of free ammonia [NH,] as a variable. (2) The concentration of free nickel ion is calculated directly from Eq. (2) and the concentration of hydrogen ion [H+ ] is obtained by solving the quadratic equation (4). Then the total added ammonia [NH31added is calculated by Eq. (3). (3) The pH and AE are acquired from [Hf ] and E!q. (1) by introducing the activity coefficients of hydrogen and nickel ions.

S. Zhong, M. T. Hepworth / Hydrometallurgy 38 (1995) 15-37

20

0.25

,

/0.20

/ : -

/

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

/

/ I I

? -0.15 w 4

I j//p-------:::-,.

0.10 -I ’ --, I

_ .--.--

25°C iooOc 200DC

0.05

Fig. 2. Ni-NH3-Hz0 system: the effect of ammonia addition on chemical driving force as a function of temperature. Conditions: 2 M (NH.,),SO,; 1 M MSO,; PH2= 30 atm.

correspond with the values employed in commercial practice for hydrogen reduction of ammoniacal nickel solutions. In addition, increasing temperature will increase the driving force for the reduction reaction (Fig. 2). Increasing hydrogen pressure also increases the driving force for the reaction (Figs. 1 and 2). The higher initial concentrations of ( NH4) $04 also has an adverse effect on the chemical driving force, because the ammonium ion acts as an acid (Fig. 1). 2.2. Magnesium-ammonia-water

system

The following analysis for the magnesium-ammonia-water system is similar to that described for the nickel-ammonia-water system. The parameters are listed in Table 2. The solution under consideration is magnesium chloride. In order to recover magnesium from this solution, ammonia is added to precipitate magnesium as Mg( OH),. The following assumptions are made in the calculations which follow: The initial solution concentration of MgCl, is ‘rn’M, where m is a variable. The equilibrium between the solution and Mg(OH)* is given by: [Mg2+] = [H+122 W

&,= Wg2+l[OH-12=

GiJ

Wg2+l [H+32

and the fraction precipitated is given by:

S. Zhong, MT. Hepworth/Hydrometallurgy Table 2 Thermodynamic

data for Mg-NH,-Hz0

38 (1995) 15-37

21

[9,10,13,14]

Reaction

Proper ty

Literature value

NH3+H+ =NH,’ H20=H+ +OHMg*++OH-=Mg(OH)+ Mg(OH),=Mg’+ +20HMg’+ +NH,=Mg (NH3)*+ Mg’+ +2NH,=Mg(NH&+ Mg2++3NH3=Mg(NH,):+ Mg2+ +4NH,=Mg(NH,):+ Mg’+ +5NH,=Mg(NH,):+ Mg*+ +6NH3=Mg(NH&+

r)=

m-[Mg*+l,+Wg*+l1. m

(5)

m

The mass balance for soluble Mg is given by:

~~g2+l,=~~g2+l+~~g~~~~+l+~~~g~~~,~~+l j=l _

=[Mg*+]{l+KJOH

I + iP,[NH,Ij) j=l

[Mg*+l,=[H+]*~(l+K,~

W+l

W

The

(6)

+ iPj[NH,ljI j=l

mass balance for NH3 is given by:

[N%lT= WM + [NH,+1+ &MNW;+l j=l 6

= [?JJH,I+KN[NH,I [H+l + [M$+] Cjpj[NH,]‘=

j=l

[NH,] +KN[NH3] [H+] + [H+,*%Kk.s.IPj[NH,lj= 6 ’

[NH31added

[NH,Ias~ti

J

The charge balance is given by:

2Wg*+l+ MdOW+l+2

; [MgWH,)f+l j=l + [H+] + [NH,+] = OH-] + [Cl-]

[Mg2+]{2+K,[OH-]

+2 tP,[NH,]‘)

j=l + [H+] +K,[H+]

[NH, = [OH-] +2m

(7)

22

S. Zhong, M. T. Hepworth / Hydrometallurgy 38 (1995) 15-37

[H+]‘${2+Kl~+2

j$[NH,,‘]

W

j=l

+[H+l

+KN[H+][NHJ

Kw

= [H+l +2m

2~11+~~j,NH,1’)[H+,3+~l W

j=l

+~+KN[NH1,~~H+]‘-2m[H+]

-K,=O

(8)

AW

The results’ are displayed in Figs. 3-5. Upon addition of ammonia, the solution pH increases rapidly at first then slowly (Fig. 3). The higher the initial concentrations of MgCl*, the greater the ammonia addition required to reach the same pH. For equivalent additions of ammonia, the pH values for solutions with higher initial concentration of MgCI, are lower. The plot of the total magnesium concentration against the added ammonia (Fig. 4) exhibits a minimum in the total magnesium ion concentration. Fig. 5 shows that there is a maximum percentage of magnesium precipitated. In magnesium chloride solutions, the added ammonia initially acts as a base to precipitate magnesium as Mg(OH)*. After the percent precipitated reaches a maximum value, an increase in the free ammonia concentration increases the complexation between magnesium ions and ammonia. At the same time the high pH increases the tendency toward complexation between Mg*+ and OH-. Therefore, the percent precipitated decreases with further additions of ammonia. Increases in initial MgC& concentration require more ammonia to precipitate Mg( OH)2. Therefore, the maximum point shifts to the right. Using the results shown in Fig. 5 and additional calculations, the maximum percentage precipitated is plotted against the different initial concentrations of magnesium chloride in Fig. 6. The higher the initial concentration of MgC12,the lower the maximum percentage of precipitated Mg( OH)*. A high solution pH improves Mg( OH) 2 precipitation. 2.3. Copper-ammonia-water

system

The following analysis for the copper-ammonia-water system is similar to that just described for the nickel and magnesium systems. The system parameters are listed in Table 3. Secondary copper-based materials may be dissolved in hot, oxygenated sulfuric acid; however, leaching using a combination of ammonia, ammonium carbonate and oxygen is often more effective and has the advantage of not dissolving iron [ 211. Although sulfuric acid leaching of chrysocolla (a hydrous copper silicate, CuO. SiOp . 2H20) is generally accepted as practical, the preponderance of carbonates in ’ The calculation procedure is as follows: ( 1) For a fixed initial concentration of MgCI, a series of values ranging from 10-O to 1 M is assigned to the concentration of free ammonia [NH,]. (2) The concentration of hydrogen ion, [H+ 1,is obtained by solving the cubic equation (8). The total soluble magnesium ion [ Mg*+ IT and the total added ammonia [NH,] addeQ are calculated via Eqs. (6) and ( 7). ( 3) The pH is acquired from [H+ I by introducing the activity coefficient and the fraction precipitated, 7).is obtained from Eq. (5).

S. Zhong, MT. Hepworth/Hydrotnetaliurgy 38 (1995) 15-37

10.0

23

(

9.5 -

z

9.0 -

d

iii

8.5 -I 7

$

ooooo m=0.5M

0~ m=l.OM be5ce m=2.OM

~~~~3, M~-NH~-HJ) sygem: the effect of ammonia addition on solution pH as a functionof MgCIZconcentration.

certain ore deposits can cause significant increases in acid consumption. Ammonia is an attractive reagent because it does not react with carbonates, and because of its ease of handling, low inventory cost and amenability to regeneration [ 221. The application of ammonia pressure leaching to the treatment of copper concentrates has been widely discussed and investigated [ 231. The overall reaction for the oxidation reaction of copper sulfide ore as chalcopyrite can be expressed by: 2CuFeSz + 12NH3 + 81 /202 + (n + 2)H,O = 2Cu( NH3) $04 + 2( NH4) $04 + Fe,O, . nH,O The thermodynamic analysis below is for the application copper, cuprous oxide and cupric oxide.

of ammonia leaching of metallic

2.3.1. Equilibria between the solution and copper metal The following assumptions are made in the calculations below: for all of these calculations, the initial concentration of CuSO, is 1.O M, and the concentration of ( NH4)$04 is ‘m’ M, where m is a variable. The solution is, for this section, equilibrated with metallic copper with the addition of ammonia to the solution. In addition to the complexing equilibria between ammonia and copper ions, other equilibria between the solution and copper are given below: Cu2+ + 2e = Cu cu2+ +cu=2cu+

EhcU2+,cU = 0.337 + O.O29510ga,,2+ [Cu’]

=Jfm

The mass balance for Cu is shown as:

(9)

S. Zhong, MT. Hepworth / Hydrometallurgy 38 (1995) 15-37

24

2.0

1.5

mwu

m=0.5M

00000

m=l.OM

*I

m=2.oM

=i \

75 _g l*O +&

“ul

E

0.5

0.0

Fig. 4. Mg-NH3-Hz0 system: the effect of ammonia addition on total magnesium ion concentration as a function of MgClz concentration.

[al],=

[Cu2’] +

L ,cu(NH,)f+,

+ [Cu’] + i [Cu(NH,)f]

i=l

j=l

=,C”2+]11+~~t[NH~,i~+[Cu+,~l+~~j’[NH i=l

{

1+ ~pi,NHs,y,cu2+ i=l

]2-

j=l

(2( 1+ i/3JNH31’) i=l

+K,(l+~P;[NH,1’)2)[cu2+,+1=0 (10) The mass balance for NH1 is shown as:

S. Zhong, M.T. Hepworth/Hydrometallurgy

38 (1995) 15-37

2.5

80

ooooo m=O.W ? -? m=l.OM AAAAA

m+&()M

0 Fig. 5. Mg-NH,-H,O system: the effect of ammonia addition on the percentage of Mg(OH), precipitation as a function of MgCI, concentration. 2

5

[NH,].=

[NH,1 + [NH:1

+ Ci[Cu(NH,)f+]

+ C j[Cu(NH,)f] j=l

i=l

+Kv[H+]

+

[ NH,

[NH,]

+

[cu2’]

t&[NH3]‘+ i=l

[Cu’]

1added

[NH,1 +fGdH+ 1[NH,1

i=l

The charge balance for the system is given by:

5 jfij[NH,]j=Zm j=l

j=l

= [NH,]

26

S. Zhong. MT. Hepworth/ Hydrometallurgy 38 (1995) 15-37

Fig. 6. Mg-NHS-Hz0 system: the effect of initial magnesium chloride concentration on the maximum percentage of Mg(OH)Z precipitation.

Table 3 Thermodynamic data for CU-NH,H,O

[ 1,9-II,15 ]

Reaction

property

NH3+H+ =Na HaO=H+ +OH1/202 + 2HC + 2e = Hz0 2H++2e=Hz cl?+ +cu=2cu+ 1/2CuzO+ I/2Hz0=Cu+ +OHCu’+ +H,O=CuO+2H+ 2CuZ++HzO+2e=Cu,0+Ht 2CuO + 2H+ + 2e = Cu,O + Hz0 CuZO+2H++2e=2Cu+H,0 Cu*+ +2e=Cu CU~++NH~=CU(NH~)~+ Cu*+ +2NHj=Cu(NH,);+ Cu2+ +3NHX=Cu(NH,):+ Cu’+ +4NH,=Cu(NH,):+ CL?+ +5NH,=Cu(NH,)j q+ Cu++NH,=Cu(NH,)+ Cu’ +2NH1=Cu(NH,):

KN KW Eh’ Eh” Ku K, KZ Eh’ Eh’ Eh’ Eh’ ;: P3 P4 Ps PI Pz’

Literature value ,po lo-

13.96

1.229 0 IO-0.” ,o- 14.7h ,0-T.% 0.203 0.673 0.474 0.337 104.= 107.83 1()‘O.R3 ,0’3.00 ,012.41 ,05.93 ,010.86

S. Zhong, M.T. Hepworth/Hydrometallurgy

00000 00000

38 (1995) 15-37

?I

m=lM m=2M

0 ;I,

,,,,,,,,,‘,,,,,,,,,,,,,,,,,,

2

4

6 PH

8

J

Fig. 7. Equilibria between solution and copper: the effect of ammonia addition on solution pH as a function of (NH,),SO, concentration.

2[cu*+]

+2

t

,Cu(NH,)j?‘,

+ [Cu’]

+

i=l

&Cu(NH,)f ] j=l

+[H+]+[NH,+]=[OH-]+2[SO:-] 2[CuZ+,~1+~~i,NH~l’)+~~~l+~P;~NH,l~l i=l

j=l

+

[H+] +K,[H+]

KW

[NH,] = ,H+l

+2(1

+m)

(1+K,[NH,]][H+]z+{2[Cu2+](l+~~i[NHs]i)+~~(1 i=l 2

+CP;[NH,]‘)-2(1+m)][H+]-K,=O j=L

(12)

28

S. Zhong, M.T. Hepworth / Hydrometallurgy 38 (I 995) 15-37

Solution

-0.21,~,,,,~,,,~~~~~~~~~,~~~~~‘~~~~~’~~~~~~~ 4 b 2 PH

B

13

Fig. 8. Eh-pH diagram for equilibria between solution and copper as a function of (NH&SO.,

concentration.

The relationship3 between pH and [ NH31addedis shown in Fig. 7. At the same time, the Eh-pH diagram (Fig. 8) for the equilibria between the solution and copper metal is plotted by considering the equilibrium between copper and cuprous oxide. The dashed lines in Fig. 8 show that the assumed equilibrium between the solution and copper is replaced by the equilibrium between copper and cuprous oxide: CuzO + 2H + + 2e = 2Cu + Hz0

EhcU~o,cu= 0.47 1 - 0.0591pH

The diagrams show that, for the leaching of secondary copper with an oxidant, the basic conditions required are approximately that the solution Eh must be greater than that of equilibrium line, the pH must be greater than 7.2-7.6 which is a function of the initial concentration of ( NH4) $04, and the ratio of [NH,] added/[ Cu] T must be greater than 3.1. Also, for pH values less than 7.2-7.6, the oxidation product is either cupric oxide or cuprous oxide rather than the leach solution. The initial concentration of ( NH&SO4 has an adverse effect on the stability of copper with a decrease in its range of stability. Increasing the concentration of ( NH4) *SO4 will improve the leaching of native copper ore since a higher concentration of (NH,) $04 causes more copper to dissolve in the form of copper amine. The leach solution contains Cu*+, Cu+, Cu(NH,)?+, Cu(NH,)f. The distribution of 3The calculation procedure is as follows: ( 1) For a fixed initial concentration of ( NH.,)2S04, a series of values ranging from 10S6 to 1 M is assigned to the concentration of free ammonia [NH,]. (2) The concentrations of free copper ion [Ct?+ ] and hydrogen ion [H+ ] are calculated by solving the quadratic equations ( 10) and ( 12). The total added ammonia [NH,] addedis obtained through Eq. ( 11) ( 3) The pH and Eh are obtained from 1H+ 1 and Eq. (9) by introducing the activity coefficients of hydrogen and copper ions.

S. Zhong, M.T. Hepworth/Hydrometallurgy

copper-bearing trical potential.

29

38 (1995) 15-37

species depends on the added amount of ammonia and the controlled

elec-

2.3.2. Equilibria between solution and cuprous oxide The calculations are, as previously stated, for 1 M CuS04 and ‘m’ M (NH4),S0,. However, in this case, the solution is equilibrated with cuprous oxide. In addition to the complexing equilibria between ammonia and copper ions, other equilibria between the solution and cuprous oxide are possible and are given below: 2Cu2++H,0+2e=Cu,0+2H+ Eh,,z + /cuzo = 0.203 + 0.059110ga,,2+ $~J~O+~H,O=CU+

[Cu’]

+0X

+ 0.059lpH

K,

=-

[OH-I

(13)

=K’,H’] Kw

The mass balance for Cu is shown as:

[a~],= [CU*+I+ i

[Cu(NH,)?+

3 + [CU+I+ t

[Cu(NH,)f

]

j=l

i=l

=[CU*+](~+~P,[NH,II)+[~U+](~+~~~~NH~~~)=~ j=l

i=l

[c~*+I(I+~~~~NH~~~~+~~H+~~I+~P;~NH~~~~=I W

i=l

(14)

j=l

The mass balance for NH3 is shown as:

[NH31T= [NH,1 + [NH21 + &IC~(NH,);+]

+ ~~[cu(NH,),+] j=l

i=l

=[WI =

[WI

2m

+

+KdH+l [NH,1 + [Cu*+] iipJNH,1’+ i=l

[ NH3

[Cu’]

‘j?jp;[NH,]j j=l

1added

+KiJH+l [NH,1 ‘&/3i[NH,]‘=Zm+

+ [Cu*+] tiB,[NH,l’+:[H+] i=l

The charge balance is shown as:

W

j=l

[NH,],,,

(IS)

S. Zhong, M.T. Hepworth / Hydrometallurgy 38 (1995) 15-37

30

2[Cu2’]

+2

;

,Cu(NH&+,

+ [Cu’]

i=l

+ i,Cu(NH,)f, j=l

+ [H+] + [NH,+] = [OH-] ~[cu~+I

{ 1+ ~~JNH,I’) i=l

+FlH+i W

11+ ~P;INH,]~) j=l

+ W+l +K,W+l [NH,1 = combining

+2[SOj-]

Kw ,H+l

+2(1

+m>

with Eq. ( 14) :

{l+K,[NH,I

+,l+

~~~,NH,]‘]J[H+]2-2m[H+] W

-K,=O

j=l

The relationship4 between pH and [ NH31 addedis shown in Fig. 9. The Eh-pH diagram (Fig. 10) for equilibria between solution and cuprous oxide is constructed by considering two pairs of equilibria: for copper and cuprous oxide and for cuprous oxide and cupric oxide. As previously stated, the dashed lines in Fig. 10 show that the assumed equilibrium between the solution and cuprous oxide is replaced by the equilibrium between cuprous oxide and cupric oxide: 2CuO + 2H + + 2e = Cu20 + H20 Cu,O + 2H+ + 2e = 2Cu + H20

EhCuo,cuzo = 0.669 - 0.0591 pH EhCuzo,cu = 0.47 1 - 0.0591pH

The following conclusions can be derived from Figs. 9 and 10: ( 1) For the leaching of cuprous oxide ore with oxidant, the basic conditions required, approximately, are that, in basic solutions, the solution Eh must be greater than that of the equilibrium line, the pH must be greater than 6.8-7.4, which is a function of the initial concentration of ( NH4) $04, and the ratio of [ NH31 added/[ Cu] T must range from 3.0 to 3.5. In acidic solutions, the Eh must be also be greater than that of the equilibrium line, the solution pH must be less than 4, and the ratio of [NH,] added/ [ Cu], must be less than 0.3. (2) For 3.8 < pH < 6.8-7.3, the oxidation product is cupric oxide instead of the leaching solution. (3) Under reducing conditions, cuprous oxide is reduced to copper. (4) The initial concentration of ( NH4)2S04 has an adverse effect on the stability of cuprous oxide with a decrease in its stability range. Increasing the concentration of (NH,) $0, improves the leaching of cuprous oxide ore.

4 The calculation procedure is as follows: ( 1) For a fixed initial concentration of (NH,) *Sod, a series of values ranging from 1O-6 to 1 M is assigned to the concentration of free ammonia [NH,] (2) The concentrations of free copper ion and hydrogen ion [H+ ] are calculated through Eqs. ( 14) and ( 16). and the total added ammonia, INWadded,is obtained from Eq. ( 15). (3) The pH and Eh are acquired from [H+ ] and Eq. ( 13) by introducing the activity coefficients of hydrogen and copper ions.

S. Zhong, MT. Hepworth / Hydromeiallurgy 38 (1995) 15-37

31

3

Fig. 9. Equilibria between solution and cuprous oxide: the effect of ammonia addition on solution pH as a function of (NH,),SO, concentration.

Solution

0.1

Fig. 10. Eh-pH concentration.

diagram

for the equilibria

between the solution and cuprous oxide as a function of (NH,),SO,

S. Zhong, M. T. Hepworth / Hydrometallurgy 38 (1995) 15-37

32

D

Fig. 11. Equilibria between solution and cupric oxide: the effect of ammonia addition on solution pH as a function of (N&) *SO4concentration.

The leaching solution contains Cu*+, Cu+, Cu(NH,)?+, Cu(NHs)‘. The distribution of copper-bearing species depends on the added amount of ammonia and the value of the electromotive potential of the solution. 2.3.3. Equilibrium between solution and cupric oxide The calculations are, as previously stated, for ‘m’A4 (NH,) $0,; however, in this case, the solutions are equilibrated with cupric oxide. In addition to the complexing equilibria between ammonia and copper ions, other equilibria between the solution and cupric oxide are given below: Cu*+ +H

20

=CuO+2H+

[cuZ+]

~L!gx 2

The mass balance for soluble Cu is given by:

[c~*+]~=[c~?+]+~[cu(NH~):+] i=l

(17)

5’. Zhong, M.T. Hepworth/Hydrometallur~y 38 (1995) 1.5-37

00000

m=lM

? ?nnno

m=2M

Fig. 12. Solubility diagram for the equilibria between solution and cupric oxide as a function of (NH,)$O, concentration.

The mass balance for NH, is given by: [NH31T= [NH,] + [NH,+] + &[Cu(NH,):+] i=l

=

[NH,] +K,[H+]

[NH,1 +GJH+l

[NH,] + [Cu2+] &&[NH,,‘=2m+ i=l

[NH,1 + ~.+NHJ=~~+ r=l

The charge balance is given by:

[NH,,,,,

[NHJadded

S. Zhong, M.T. Hepworth / Hydrornetallurgy 38 (1995) 15-37

34

0.75

0.50

2

0.25

4

0.00

-0.25

-0.50

Fig. 13. Eh-pH diagram for Cu-NH,-Ha0

2[Cu*+]

+2 5 [Cu(NH,);+] i=l

2?{,

+ ‘&,NH,,‘J 2

system, initial (NH&SO,,

+ [H+] + [NH,+] = [OH-]

+ ,H+,

+K,,H+]

[NH,,

concentration

1 M.

+2[SO:-]

=-+,+2(1

+m)

i=l

2~{~+~~i[NH,,i),H+,3+~~ 2

i=l

+K,.,[NH3]}[H+]2-2(1+m)[H+]-Kw=0

(19)

The calculated data5 are plotted in Figs. 11 and 12. Upon adding ammonia to copper solutions, the ammonia reacts with cupric ion, at first, to produce Cu( OH),, according to the reaction 2NH, + H,O + Cu*+ = Cu( OH), + 2NHz, and the solution pH increases slowly (Fig. 11). After about 2 mol/l of ammonia has been added to the initial solution and the above neutralization reaction is completed, the ammonia acts as a base to deliver ’ The calculation procedure is as follows: ( I ) For a fixed initial concentration of ( NH&S04, a series of values ranging from lo-’ to 1 M is assigned to the concentration of free ammonia, [NH, 1. (2) The concentration of hydrogen ion, [H+ 1,is calculated by solving the cubic equation ( 19). The total dissolved copper ion and the total added ammonia, [NH,],,,, are obtained from E?qs. ( 17) and ( 18). (3) The pH is acquired from [H+ ] by introducing the activity coefficient.

S. Zhong, M.T. Hepworth/Hydrometallurgy

38 (1995) 15-37

3.5

0.75

0.50

E L1

0.25

W

0.00

-0.25

-0.50

3

Fig. 14. Eh-pH diagram for Cu-NH,-HZ0

system, initial (NH,) 2S04 concentration

2 M.

hydroxyl ion directly, according to the reaction NH, + Hz0 = NH: + OH-. As a result, the solution pH rises rapidly. The added ammonia then acts mainly as a complexing agent to react with cupric oxide, yielding amine. The concentration of total dissolved copper increases rapidly. From Fig. 12, it is known that, for the leaching of cupric oxide ore with ammonia, the basic conditions required for 1 M of total dissolved copper are approximately that the solution pH must be greater than 7-7.6, which is a function of the initial concentration of ( NH4) $04, and the ratio of [NH,] added/ [ Cu] T must be greater than 3.7-3.9. These figures also indicate that higher initial concentrations of (NH,)$04 have adverse effects on the stability of cupric oxide. An increase in the concentration of (NH,)$O, improves the leaching of cupric oxide ore. The leaching solution contains cupric ion, Cu’+, and its complexes, Cu( NH,)?+. The distribution of copper-bearing species depends on the amount of ammonia added.

2.3.4. Eh-pH diagram for copper-ammonia-water system Figs. 13 and 14 are two Eh-pH diagrams, 1 M and 2 M are the initial concentrations of (NHJ $04 for the copper-ammonia-water system constructed by combining the Eh-pH relationships of equilibria among the solution and copper, cuprous oxide and cupric oxide. The stability region of water is also shown according to following equations:

S. Zhong, M.T. Hepworth/Hydrometallurgy

36

2H+ +2e=H, $+2H+

38 (1995) 15-37

Eh,+,,,=O-0.0591pH

+2e=H,O

Eho2,HZo= 1.229-0.0591pH

The two diagrams appear to be different from the traditional Eh-pH diagrams. There are no regions for individual ions instead of the region for the solution. Since ammonia rather than non-ammoniacal acid or base is added to adjust the solution pH, the solution pH is less than 9. From Figs. 13 and 14, it is known that, for the ammonia leaching processes of copper, cuprous oxide and cupric oxide, a pH range between 8 and 9 is required. Comparing Figs. 13 and 14, one should note that increasing the initial concentration of (NH&SO4 has an adverse effect on the stability range of Cu, CuzO and CuO. An increase in the initial (NH,) *SO4 concentration improves the leaching of copper, cuprous oxide, and cupric oxide.

3. Summary Based on both mass and charge balances, calculations have been made, with examples, in the Cu-, Mg- and Ni-ammonia-water systems to express equilibrium relations between solid phases and associated aqueous phases. More quantitative information on practical hydrometallurgical processes can be obtained from these calculations compared with the conventional calculation. This calculation method can readily be applied to other metalammonia-water systems.

References [ 11 Pourbaix, M., Atlas of Electrochemical Equilibria in Aqueous Solutions. Pergamon, London ( 1966). [2] Garrels, R.M. and Christ, C.L., Solutions, Minerals, and Equilibria. Harper and Row, New York ( 1965). [3] Osseo-Asare, K. and Fuerstenau, D.W., Application of activity-activity diagrams to ammonia hydrometallurgy. The copper-, nickel-, and cobalt-ammonia-water systems at 25°C. In: T.W. Chapman, L.L. Tavlarides, G.L. Hubred and R.M. Wellek (Editors), Fundamental Aspects of Hydrometallurgical Processes. AIChE Symp. Ser., 173(74) (1978). pp. 1-13. [4] Osseo-Asare, K., Application of activity-activity diagrams to ammonia hydrometallurgy. II, The copper-, nickel-, and cobalt-ammonia-water systems at elevated temperature. In: M.C. Kuhn (Editor), Process and Fundamental Considerations of Selected Hydrometallurgical Systems. AIME, New York ( 1981), pp. 360369. [5] Osseo-Asare, K., Application of activity-activity diagrams to ammonia hydrometallurgy: 3 - Mn-NH,H,O, Mn-NH,-H&CO,, and Mn-NH,--H*O-SO, systems at 25°C. Trans. Inst. Min. Metall. Sect. C: Mineral Process. Extr. Metall., 90 (1981): 152-158. [6] Osseo-Asare, K., Application of activity-activity diagrams to ammonia hydrometallurgy: 4 - Fe-NH,H20, Fe-NHr-H,O-CO,, and Fe-NH,-H20-SO4 systems at 25°C. Trans. Inst. Min. Metall. Sect. C: Mineral Process. Extr. Metall., 90 (1981): 159-163. [ 71 Luo Rutie, Overall equilibrium diagrams for hydrometallurgical systems: copper-ammonia-water system. Hydrometallurgy, 17 (1987): 177-199. [ 81 Butler, J.N., Ionic Equilibrium: A Mathematical Approach. Addison-Wesley, London (1964). [9] Sillen, L.G. and Martell, A.E., Stability constants of metal-ion complexes. Spec. Publ. Chem. Sot., 17 (1964) and Spec. Publ. Chem. Sot., 25 (1971).

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[ 141 [ 151 [ 161 [ 171

[ IS]

[ 191 [ 201 [21] [22] [23]

Smith, R.M. and Martell, A.E., Critical Stability Constants.

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and Vol. 6

(1989). Kotrly, S., Handbook of Chemical Equilibria in Analytical Chemistry. Ellis Horwood, Chichester, UK (1985). Yatsimirskii, K.B. and Vasilev, V.P., Instability Constants of Complex Compounds. Pergamon, New York (1960). Perrin, D.D., Ionisation Constants of Inorganic Acids and Bases in Aqueous Solution. Pergamon, New York (1982). Baes, C.F., Jr. and Mesmer, R.E., The Hydrolysis of Cations. Wiley, New York ( 1976). Latimer, W.M., The Oxidation States of the Elements and Their Potentials in Aqueous Solutions. PrenticeHall, Englewood Cliffs, N.J., 2nd ed. ( 1952). Zaytsev, I.D. and Aseyev, G.G., Properties of Aqueous Solutions of Electrolytes. CRC Press, Ann Arbor, Mich. ( 1992). Criss, C.N. and Cobble, J.W., The thermodynamic properties of high temperature aqueous solutions. IV. Entropies of the ions up to 200°C and the correspondence principle. J. Am. Chem. Sot., 86 (1964): 53855390. Criss, C.N., and Cobble, J.W., The thermodynamic properties of high temperature aqueous solutions. V. The calculation of ionic heat capacities up to 200°C. Entropies and heat capacities above 200°C. J. Am. Chem. sot., 86 (1964): 5390-5393. Boldt, J.R., The Winning of Nickel. Van Nostrand, Princeton, N.J. ( 1967). Evans, D.J.I., Advances in Extractive Metallurgy. Inst. Min. Metall., London (1968). Jackson, E., Hydrometallurgical Extraction and Reclamation. Ellis Horwood, Chichester, UK (1986). Mena, M. and Olson, F.A., Leaching of chrysocolla with ammonia-ammonium carbonate solutions. MetaIl. Trans. B, 16B(9) (1985): 441-448. Tozawa, K., Umetsu, Y. and Sato. K., On chemistry of ammonia leaching of copper concentrates. In: Extractive Metallurgy of Copper, Vol. II. Hydrometallurgy and Electrowinning. AIME, New York ( 1976).