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Diagrams of electrochemical equilibria of the system copper-potassium ethylxanthate-water at 25°C
International Journal of Mineral Processing, 4 (1977) 345--361 345 © Elsevier Scientific Publishing Company, Amsterdam -- Printed in The Netherlands
DIAGRAMS OF ELECTROCHEMICAL EQUILIBRIA OF THE SYSTEM COPPER--POTASSIUM ETHYLXANTHATE--WATER A T 25 ° C
TADEUSZ HEPEL and ANDRZEJ POMIANOWSKI Department of Hydrometallurgy, Institute of Chemistry, JagieUonian University (Poland) Research Laboratories of Catalysis and Surface Chemistry, Polish Academy of Sciences, Krakow (Poland) (Received June 2, 1976; revised and accepted July 11, 1977)
ABSTRACT Hepel, T., and Pomianowski, A., 1977. Diagrams of electrochemical equilibria of the system copper--potassium ethylxanthate--water at 25 ° C. Int. J. Miner. Process., 4: 345--361. On the basis of thermodynamic calculations, diagrams of metastable electrochemical equilibria of the system copper--potassium ethylxanthate--water, at 25°C were constructed. The equilibria equations and diagrams for the total activity of [EtX-] + [HEtX] + 2 [(EtX)2 ] (equalling 10 -1, 10 -4 and 10 -~ tool/dinS), are presented. The practical conclusions resulting from these diagrams for the flotation of native copper are discussed.
INTRODUCTION A c c o r d i n g t o the generally a c c e p t e d opinions o f Taggart and G a u d i n (1923), it is believed t h a t t o float the grains o f n o n - h y d r o p h o b i c minerals it is necessary and sufficient t o f o r m o n t h e i r surface a c o l l e c t o r coating with a thickness o f a b o u t one m o n o l a y e r . At the same t i m e the m e c h a n i s m o f the interaction o f t h e c o l l e c t o r with the surface o f t h e mineral is e x p l a i n e d in t e r m s o f an e x c h a n g e r e a c t i o n , a d s o r p t i o n o f ions or c o l l e c t o r molecules, or the oxid a t i o n - r e d u c t i o n processes involving t h e o x i d i z e d surface c o m p o u n d s or t h e o x y g e n dissolved in the solution. In some r e c e n t papers the influence o f semic o n d u c t i v e p r o p e r t i e s o f minerals has also been c o n s i d e r e d (Plaksin, 1957; Plaksin and Shafeev, 1959, 1960). It is characteristic t h a t , in spite o f t h e ass u m p t i o n t h a t a m o n o m o l e c u l a r c o l l e c t o r coating is f o r m e d , t h e t h e r m o d y namic d a t a referring t o the m a c r o s c o p i c phases with a definite crystallographic s t r u c t u r e are used in m a n y papers to describe the system. On t h e o t h e r h a n d , it is k n o w n f r o m t h e f o r m e r investigations o f G a u d i n and S c h u h m a n n ( 1 9 3 6 ) t h a t at higher c o l l e c t o r c o n c e n t r a t i o n s several m o l e c u l e coatings can be f o r m e d (the s y s t e m a m y l x a n t h a t e - - c h a l c o c i t e was e x a m i n e d ) . In his electrochemical investigations Woods ( 1 9 7 1 ) p r o v e d t h a t the e t h y l x a n t h a t e ions on p l a t i n u m , gold, c o p p e r and galena e l e c t r o d e s u n d e r g o r e a c t i o n s which are ac-
346
companied by collector adsorption -desorption processes. Similar investigations on mercury were done earlier by Pomianowski (1967) (cf. also Hepel et al., 1973). Recently (Pomianowski and Kowal, 1973) he extended them to copper and chalcocite. The thermodynamic approach to the equilibria attained in flotation systems includes, especially in recent papers (Abramov, 1968, 1970a, b; Majima and Takeda, 1968; Fuerstenau et al., 1968; Toperi and Tolun, 1969; Majima, 1969), not only the "phase" of the c o m p o u n d of the collector with the particular metal, (eg. Pb(C2HsOCS2)2, Cu2(C2HsOCS2h, etc.), but also the surface phases formed by the mineral's oxidation (eg. Cu(OH)2, Cu2(OH)2CO3, Fe(OH)3, etc.) which take place when the mineral reacts with the collector solution, or earlier in the geochemical processes or during the extracting of the ore. In the present paper (cf. also: Hepel, 1973) the authors try to describe the system metallic copper--potassium xanthate aqueous solutions. Because of the complicated behaviour of the sulphide minerals in the flotation process conditions, a complete description of sulphide systems at this stage of the investigations was given up and a more careful analysis of a simpler case -native copper flotation -- was undertaken. Only the phases, the properties of which correspond to the macroscopic phases, were considered. THE CONSTRUCTION OF DIAGRAMS
The approach offered in the present paper is based on the application of electrochemical equilibria diagrams to investigate flotation systems (Fuerstenau et al., 1968; Hepel, 1973). The constructed diagrams refer to metastable equilibria. Since for the system native copper--ethylxanthate solution there only exist comparative thermodynamic data, a calculation method enabling clear and easy operations was used; it allowed the equilibria to be calculated for more than fifty reactions, occurring in the Cu--HEtX--H20 system, which have to be considered to construct the electrochemical equilibria diagram. To estimate the values of individual standard potentials and equilibrium constants of chemical reactions, pairs of substances were compiled, as it is impossible to calculate the free energy of formation at 25 ° C (298 ° K) of substances containing the E t X - group, and on the basis of accessible information the differences in changes of the free energy of formation for each of these pairs was estimated (Hepel, 1973). Some values of the differences Ag ° (298) for some of the formed pairs are presented in Table I as examples. After having analyzed the results of potentiometric investigations, and after having calculated them, a v a l u e - 3 . 2 2 9 kcal/mot was accepted for further calculations for the pair (EtX)2(s)--2 EtX-. This corresponds to the normal oxidation potential of the ethyIxanthate ions to dixanthogen in a solution saturated with the oxidation product. For the normal potential of this reaction, a value - 0 . 0 7 0 V with respect to the standard hydrogen electrode was assumed. According to experimental data, its value ranges within rather large limits: Stepanov et al. (1959) and Kakovskii and Arashkevich (1969) give the
347
TABLE I The thermodynamic data for the Cu--KEtX--H20 system Compounds pairs
Difference AUo (29.) kcal mo1-1
J moi-' "10 -3
(EtX)~(s)--2 EtX(EtX)2(aq)--2 EtX-
- 3.229 + 3.454
- 13.52 + 14.46
(EtX)2(s)--(EtX)2(aq)
- 6.683
- 27.97
CuEtX--EtXCu(EtX)2--2 EtX-
-14.28 -15.82
- 59.78 - 66.22
Cu(EtX)2--(EtX)2(s)
-12.59
- 52.70
Cu(EtX)~--(EtX)2(aq )
-19.27
- 80.66
2 CuEtX--(EtX)2(s)
-25.33
-106.0
2 CuEtX--(EtX)2(aq)
-32.01
-134.00
Cu(EtX)2--2 CuEtX
+12.74
+ 53.33
HEtX--EtX-
- 2.076
-
8.69
Source
see text Leja (1973); Tipman and Leja (1975)*, ** Leja (1973); Tipman and Leja (1975)* Sheka and Kriss (1959)* Kakovskii and Arashkevich (1969)*, ** Kakovskii and Arashkevich (1969)*, 8 ** Kakovskii and Arashkevich (1969); Leja (1973); Tipman and Leja (1975)*, ** Sheka and Kriss (1959); Abramov (1968)*, ** Abramov (1968); Leja (1973); Tipman and Leja (1975)*, ** Abramov (1968); Kakovskii and Arashkevich (1969)* Kakovskii and Silina (1962)*
*Computed from data. **For the difference (EtX)2(s)-2 EtX-, a value of ~uo(29s): -3.229 kcal/mol assumed in this work was used for calculations. value - 0 . 0 3 7 V, Majima a n d T a k e d a ( 1 9 6 8 ) t h e value - 0 . 0 4 9 V, G o l d s t i c k (1959); - 0 . 0 5 3 V, Du Rietz ( 1 9 5 7 ) : - 0 . 0 6 7 V a n d T o l u n a n d K i t c h e n e r (1963): - 0 . 0 8 1 V; finally t h e e s t i m a t i o n b y F i n k e l s t e i n ( 1 9 6 7 ) f r o m spectrop h o t o m e t r i c investigations is t h a t t h e s t a o d a r d p o t e n t i a l o f t h e ( E t X ) 2 - - E t X s y s t e m ranges f r o m - 0 . 0 8 0 t o - 0 . 1 4 0 V. We a s s u m e d a r a t h e r l o w e r value o f t h e s t a n d a r d p o t e n t i a l because: (1) activity o f e l e c t r o c h e m i c a l l y d e p o s i t e d (EtX)2 on t h e s u r f a c e o f t h e e l e c t r o d e is r a t h e r greater t h a n u n i t y ; (2) activity of E t X - ions is l o w e r t h a n t h e i r c o n c e n t r a t i o n in t h e s o l u t i o n s applying. To construct the diagrams for a Cu--KEtX--H20 system, the new solutions f o u n d f o r a s y s t e m C u - - H ~ O 4 - - H 2 0 ( H e p e l a n d P o m i a n o w s k i , 1 9 7 4 ) were used; t h e y d i f f e r f r o m t h e o n e s o b t a i n e d b y De Z o u b o v et al. {1966) in considering t h e stable c o m p l e x Cu(OH)2 (aq). This caused a change in t h e s y s t e m o f equilibria b e t w e e n t h e solid phases CuO, C u 2 0 a n d Cu a n d t h e s o l u t i o n o f dissolved c o p p e r . T h e e q u a t i o n o f t h e c h e m i c a l a n d e l e c t r o c h e m i c a l reactions, c o n s i d e r e d during c o n s t r u c t i n g t h e d i a g r a m o f p o t e n t i a l vs p H f o r C u - - K E t X - - H ~ O s y s t e m at 25 ° C, a n d t h e equilibria e q u a t i o n s are r e p r e s e n t e d b e l o w .
348
Equations o f chemical and electrochemical reactions and equilibria equations for a Cu - K E t X - H 2 0 system at 25 ° C (a) T h e r e a c t i o n s n o t i n v o l v i n g t h e E t X - - s p e c i e s : Cu2+ + e ~ Cu+
[Cu 2÷] E, = 0 . 1 5 3 + 0 . 0 5 9 1 l o g
(1)
.....
[Cu +1
Cu(OH)2(aq) + 2H + + e~Cu + + 2 H20
[Cu (OH):] ....... - 0.1182 pH
E2 = 0 . 9 6 2 + 0 . 0 5 9 1 l o g
(2)
[Cu+l
Cu2+ + 2 H 2 0 #Cu (OH)2(aq) + 2 H+
pH = 6.84 + 0.5 log
(3)
[Cu (OH)2] ..............
[Cu2+1
Cu (OH)2(aq) ~- HCuOz- + H +
(4)
[HCuO2-] pH = 13.03 + log [Cu (OHhl HCuO2- ~ CuO: 2- + H +
pH = 13.14 + log
(5)
[CuOz2-1 ........
[Cu (OH)2]
Cu (OH)g(aq) + 2 H ÷ + 2 e ~ C u + 2 H 2 0
(6) E6 = 0 . 7 4 1 + 0 . 0 2 9 5 l o g [ C u ( O H ) 2 ] - 0 . 0 5 9 1 p H HCuO2- + 3 H + + 2 e-~ Cu + 2 H20
(7) E7 = 1 . 1 2 7 + 0 . 0 2 9 5 l o g [ H C u O 2 - ] CuO2 g - + 4 H + + 2 e ~ C u
- 0.0886 pH
+ 2 H20
(8) Es = 1 . 5 1 5 + 0 . 0 2 9 5 l o g [ C u O 2 2 - ] - 0 . 1 1 8 2 p H 1/2 C u 2 0 + H
÷+e~Cu+'~
H20
(9) E9 = 0 . 4 7 1 - 0 . 0 5 9 1 p H
349
CuO + H + + e # 1A C u 2 0 + '/~ H 2 0
(10) E,o = 0 . 6 9 7 - 0 . 0 5 9 1 p H 2 Cu (OH)2 (aq) + 2 H + + 2 e ~ C u 2 0 + 3 H 2 0 (11) El, = 1 . 0 1 2 + 0 . 0 5 9 1 log [Cu ( O H ) 2 ] - 0 . 0 5 9 1 pH 2 HCuO2- + 4 H + + 2 e # Cu20 + 3 H20 (12) E,2 = 1 . 7 8 2 + 0 . 0 5 9 1 log [ H C u O 2 - ] - 0 . 1 1 8 2 p H 2 CuO22- + 6 H ÷ + 2 e # C u 2 0 + 3 H20 (13) E,3 = 2 . 5 5 9 + 0 . 0 5 9 1 log [CuO22-] - 0 . 1 7 7 3 p H CuO22- + 2 H + # CuO + H 2 0 (14) p H = 1 5 . 7 5 + 0.5 log [CuO22-] CuO + 2 H ÷~ Cu2+ +
H20 (15)
pH = 4 . 1 8 - 0.5 log [Cu 2+] Cu ÷ + e ~- Cu (16)
E,6 = 0.520 + 0.0591 log [Cu+] (b) The reactions of KEtX: (EtX)2(s) + 2 e ~ 2 EtX(17)
E,7 = - 0 . 0 7 0 - 0.0591 log [EtX-] (EtX)2(aq) + 2 e ~ 2 E t X E,8 = 0 . 0 7 5 + 0 . 0 2 9 5 log
(18)
[(EtX)21 [EtX-] 2
HEtX ~ EtX- + H ÷ p H = 1.52 + log
(19)
[EtX-] [SStX]
(EtX)2(aq) + 2 H ÷ + 2 e ~ 2 H E t X E2o = 0 . 1 6 5 - 0 . 0 2 9 5 log
[(EtX)2I [HEtX] 2
(20) - 0.0591 pH
350
(EtX)2(s) + 2 H + + 2 e ¢ 2 H E t X
(21) E2, = 0 . 0 2 0 -
0.0591 log [HEtX]-
0.0591 pH
(c) T h e r e a c t i o n s i n v o l v i n g c o p p e r a n d x a n t h a t e : CuEtX + e ~ Cu + EtX(22) E2: -- - 0 . 6 1 9
- 0.0591 log [EtX-]
CuEtX + H ÷ + e ~Cu + HEtX (23) E23 = - 0 . 5 2 9
- 0.0591 log [HEtX]-
0.0591 pH
Cu(EtX)2 + 2 e ~-Cu + 2 EtX(24) E24 = - 0 . 3 4 3
- 0.0591 log [EtX-]
Cu(EtX)2 + H ÷ + 2 e ~ Cu + 2 HEtX (25) E2s = - 0 . 2 5 3
- 0.0591 log [HEtX]-
0.0591 pH
Cu ( E t X ) 2 + e ~ C u E t X + E t X (26) E26 = - 0 . 0 6 7
-
0.0591 log [EtX-]
Cu ( E t X ) 2 + H + + e # C u E t X + H E t X (27) E27 = 0 . 0 2 3 - 0 . 0 5 9 1 l o g [ H E t X ] 2CuEtX+H20
- 0.0591 pH
~Cu20+2EtX-+2H
+ (28)
pH = 18.44 + log [EtX-] CuO + EtX- + 2 H + + e ~CuEtX
+ H20
(29) E2~ = 1 . 7 8 7 + 0 . 0 5 9 1 l o g [ E t X - ] -
0.1182 pH
(EtX)2(s) + C u O + 2 H ÷ + 2 e ~ - C u ( E t X ) 2 + H 2 0 (30) E30 = 0 . 8 5 7 - 0 . 0 5 9 1 p H (EtX)2(s) + C u 2+ 2 e ~ C u ( E t X ) 2 (31)
E3, = 0 . 6 1 0 + 0 . 0 2 9 5 l o g [ C u :+]
351
(EtX)2(s) + Cu (OH)2(aq) + 2 H ÷ + 2 e ~ C u ( E t X ) 2 + 2 H 2 0
(32)
E32 = 1 . 0 1 4 + 0 . 0 2 9 5 log [Cu ( O H ) 2 ] - 0 . 0 5 9 1 p H
Cu(OH)2(aq) + E t X - + 2 H + + e ~- C u E t X + 2 H 2 0 (34) E33 = 2 . 1 0 2 + 0 . 0 5 9 1 log [ E t X - ] + 0 . 0 5 9 1 log [ C u ( O H ) 2 ] - 0 . 1 1 8 2 p H Cu :+ + E t X - + e ~ C u E t X (34) E34 = 1 . 2 9 3 + 0 . 0 5 9 1 log [Cu2÷]+ 0 . 0 5 9 1 log [ E t X - ] Cu 2+ + H E t X + e ~ C u E t X + I-I+
(35) E3s = 1 . 2 0 2 + 0 . 0 5 9 1 log [Cu 2+] + 0 . 0 5 9 1 log [HEtX]+ 0 . 0 5 9 1 p H (EtX)2(aq) + 2 C u O + 4 H + + 4 e ~ 2 C u E t X + 2 H 2 0 (36) E36 = 0 . 9 3 1 + 0 . 0 1 4 7 log [ ( E t X ) 2 ] - 0 . 0 5 9 1 p H Cu + + (EtX)2(s) + e ~ Cu (EtX)2 (37) E37 = 1 . 0 6 6 + 0 . 0 5 9 1 log [Cu +] 2 H C u O 2 - + (EtX)2(s) + 6 H + + 4 e # 2 C u E t X + 4 H 2 0
(38) E38 = 1.401 + 0 . 0 2 9 5 log [ H C u O 2 - ] - 0 . 0 8 8 6 p H 2 CuO22- + (EtX)2(s) + 8 H ÷ + 4 e ~ 2 C u E t X + 4 H 2 0
(39) Ea9 = 1 . 7 9 0 + 0 . 0 2 9 5 log [ C U O 2 2 - ] - 0 . 1 1 8 2 p H HCuO2- + EtX- + 3 H ÷ + e # CuEtX + 2 H20
(40) F~0 = 2 . 8 7 2 + 0 . 0 5 9 1 log [ E t X - ] + 0 . 0 5 9 1 log [ H C u O : - ] - 0 . 1 7 7 3 p H CuO22- + E t X - + 4 H + + e ~ C u E t X + 2 H 2 0 (41) E4, = 3 . 6 4 9 + 0 . 0 5 9 1 log [CuO22-] + 0 . 0 5 9 1 log [ E t X - ] - 0 . 2 3 6 4
pH
Cu + + H E t X ~ C u E t X + H +
(42) p H = - 1 7 . 7 5 - log [ H E t X ] - log [Cu +]
352
2 Cu 2+ + (EtX)2(aq) + 4 e ~ 2 C u E t X ( 43 i E43 = 0 . 6 8 4 + 0 . 0 2 9 5 l o g l C u > ] +
0.0147 log [(EtX)2]
2 Cu(OH)2(aq) + (EtX)2(aq) + 4 H ÷ + 4 e = 2 C u E t X + 4 H 2 0 {44) E44 = 1 . 0 8 8 + 0 . 0 2 9 5 l o g [ C u ( O H ) 2 ] + 0 . 0 1 4 7 l o g [ ( E t X ) 2 ] - 0 . 0 5 9 1 p H
~
q
600-
400
CuO (EtX)2
CufEtX)2 @~
.
(EfX)2 rE, X)2,.:.,,
':
200 ~
~(? "t "
E CO
~_
1~
CuE~X
~ -200~
%
Q~ -LO0~
-500,
Cu 2
4
6
8
10
12
1L
pH
16
Fig.1. Diagram of the metastable electrochemical equilibria for the system Cu--KEtX--H20 at 25 ° C. The initial general activity [ E t X - ] + [ H E t X ] + 2 [(EtX)2(aq) ] equals 0.1 tool/din 3. (The additional lines (24) and (25) in the diagram are shown only for comparison with results of some experiments. )
DISCUSSION OF RESULTS AND CONCLUSIONS
The most developed system for a thermodynamic approach to the flotation process was recently presented by Abramov (1968, 1970a, b). The thermodynamic equilibria considered by this author are not stable. This fact is noteworthy, as in describing the metastable equilibrium state it is necessary to regard and confirm the assumption concerning the kinetic nature of the system, and especially the restrictions of the processes leading to the establishment of stable equilibria. Simultaneously, the considered metastable
353
[ fEfXJzCu:"
~
600 F
S CuO -'EtX).
Cu (EtX)2 ~ {E~X)2
400 ~HEIX
"
~,
..b =oo
,=
E
,,,,=
,'
CU +
"tu3
~p.~
CuO
c
o -b.. 4
.._r
"'2
CuEtX
i~
~-200
%a
~Cu20 -400'-
@
""9..
Cu i
i
.:
(~)
-600 0
2
4
6
8
10
12
14
pH
i6
Fig.2. Potential--pH metastable equilibria diagram for the system Cu--KEtX--H20 at 25 ° C. The initial general activity [ E t X - ] + [HEtX] + 2 [(EtX)2(aq)] equals 10 -4 tool/din 3. (The additional lines (24) and (25) in the diagram are shown only for comparison with results of some experiments. )
equilibria must be chosen based on experimental indications, proving that they are of practical importance, i.e. that the forward and backward processes run with sufficient velocity. Another procedure suggested (Hepel, 1973} is also possible; it assumes the course of particular reactions and then by means of thermodynamic analysis estimates the parameters describing the particular states of metastable and stable equilibria and compares the results with the experimental observations. This analysis is best demonstrated when represented in graphic form. Abramov (1968, 1970a, b) solved a system of equations describing chemical and electrochemical equilibria reactions and balanced the equations, thus arriving at a number of interesting conclusions concerning the influence of different reagents competition on the flotation range of lead, copper and iron sulphide minerals. The authors of the present paper believe that the way in which Abramov selected the chemical and electrochemical equations describing the processes in real systems is too arbitrary. In another paper, H e p e l e t al. (1973) tried to represent on the Lange--Nagel diagram the surface adsorption phases formed in the aqueous mercury--ethylxanthate system. The results of earlier experimental investigations, carried
3 5 `4,
+,,
t,'J"
_
600:
.
....
_
CuO &O0:
(EfX~
, e
HEfX -~gf io
.
Cu{Of#.
EtX -
CU"
2 0 0 '~
,~
HCuO~
CuEfX
-.\
\23 +"
.
Cu
-400
'"
- 600 0
2
4
6
8
10
12
14
pH
15
Fig. 3. Diagram of the metastable electrochemical equilibria for the system Cu--KEtX--H20 at 25 ° C. The initial general activity [ E t X - ] + [ H E t X ] + 2[(EtX)2(aq ) ]equals 10 -7 tool/din 3. (The additional lines (24) and (25) in the diagram are shown only for comparison with results of some experiments. )
out by means of polarography and voltammetry (Pomianowski, 1963, 1967), were used to construct the diagram. The solution of the system of the adsorption phases equilibria was only a first attempt to describe the flotation system and so it does not pretend to be a definite one. This is due to the lack of some experimental data and on the other hand to interpretative troubles, especially in the systems having low xanthate ion concentrations. In the present thermodynamic description of the systems occurring in the flotation process, the adsorption phenomena were not taken into account because of the lack of data concerning the free adsorption energy of individual substances existing in the system. Accordingly, the results presented above may be directly applied to the flotation process only if it is known that on the surface of the flotation minerals coatings are formed of hydrophobic compounds such as: xanthate--the mineral involved (here copper), characterized by the values corresponding to the macroscopic phases of these compounds. On the other hand, when the data concerning the surface adsorption phases are obtained, the diagrams presented here may be completed and adapted for a direct interpretation of the flotation process. The diagrams of potential vs pH worked out on the basis of the above
355
600~-
~ ~9
qb
~3,
(EtX)2
' E
-"-"
~ -- - - - 4 7
i
CuO
,, Cu {EfX)2
(EtX)2
~ .I
HCuO= >
~ "~,
CuEtX
CuO
",,~
I
5
. -200
.
. . . . .
i
-4oo~
@
..Cu..
~)
-600 L ~ L I ! ~ 0
\ ...
2
X I~
\
x,
_~
~
1
4
6
~
t 8
,
• 12
1[,
pH
16
Fig.4. Potential--pH plot of the metastable electrochemical equilibria for the system Cu--KEtX--H~O at 25 ° C. General activity [ EtX- ] + [ HEtX ] equals 10-7 tool/din 3 below lines (17) and (21) and decreases at higher potential values. Activity of (EtX)2(aq) is constant above lines (17) and (21) and equals 1.25"10-s tool/din 3 and decreases at lower potential values.
equilibria reactions are presented in Figs.l--4. The first three plots correspond to the total xanthate forms activity in the solutions: 10 -1, 10 TM, 10 -7 mol/ dm 3, respectively. On all diagrams a dash-dotted line marks the areas of the relative predominance of the individual species of ethylxanthate dissolved in water. They are: E t X - ions dominating in the whole range of pH above the value 1.52 and at sufficiently low potentials, then the undissociated ethylxanthic acid HEtX existing in relative predominance in strongly acid solutions*, and dixanthogen, towards which the solution quickly becomes saturated because of its low solubility (equal to 7.9.10 -6 mol/dm 3 (Stepanov, 1960), 1.10 -s mol/dm 3 (Kakovskii et al., 1969), 1.3.10 -s mol/dm 3 (Pomianowski and Leja, 1963) and 1.25-10 -s mol/dm 3 (Leja, 1973; Tipman and Leja, 1975)}. The (EtX)~ relative predominance area occurs at higher potential values. The stability of aqueous xanthate solutions is limited by a number of conditions: the solution decomposes irreversibly at pH values lower than about 6 and greater than about 10.5, the solution is likely to be oxidized *cf. lower: the remarks on the instability of KEtX solutions.
356
(Pomianowski and Leja, 1963; Finkelstein, 1967; Shulman and Larionov, 1967; Garbacik et al., 1972) and the limiting solution concentrations is about, 0.1 tool KEtX/dm 3. in acid solutions there occurs a decomposition caused by a large shift of the equilibrium of reaction (19) towards the formation of undissociated HEtX, which undergoes an irreversible reaction, forming carbon disulphide and ethyl alcohol (Pomianowski and Leja, 1963b; Shulman and Larionov, 1967). In the alkaline environment, the decomposition of EtX-ions leads to a thiocarbonic acid salt, sulphides, disulphides, carbonates and others (Pomianowski, 1963; Finkelstein, 1967). Simultaneously, due to the --SH functional group, strong reduction abilities are characteristic for xanthates. Under these conditions it was necessary to introduce some new designations on the Pourbaix diagrams, worked out for the metastable equilibria. All the lines, concerning the chemical and electrochemical equilibria reactions occurring in the area relatively dominated by an unstable ethyloxanthate acid and "near" to this area are drawn as broken lines. The pH value of 3 is conventionally accepted as the value limiting the metastability range of E t X ions from the acid part of the pH scale, at which (as can be seen, among others, from Finkelstein, 1967), the half-life of xanthate is 5 minutes. This is sufficiently long and a flotation experiment under such conditions may be considered. The unstability of xanthate solutions in strongly alkali environment having a pH higher than about 12.5 is also marked by a broken line (Leja, 1973). Generally speaking, we assume that all substances containing the xanthate group are not stable in the absence of CS2 and EtOH (Leja, 1973; Tipman and Leja, 1975) in proper concentrations and carbonate species (at the domain of thermodynamic stability of water) that could guarantee attaining the stable equilibrium. The rate of irreversible decomposition of particular xanthate forms varies, however, in dependence on the pH value of the solution. It is conventionally assumed that in the pH range from a b o u t 3 to about 12.5 the rate of irreversible decomposition is considerably lower than the rate of processes proceeding among the solution components: EtX-, (EtX)2(aq) and solids: CuEtX, Cu(EtX)2, (EtX)2, and is n o t associated with the xanthate group decomposition. This permits consideration of metastable equilibria connected with these processes. Since a slow xanthate decomposition process proceeds parallelly, the given metastabte equilibrium can only be spoken of at the time period when the decrement of xanthate activity caused by decomposition is negligibly small. The higher the rate of the irreversible decomposition process, the shorter the duration time of metastable equilibrium. Outside the pH range from 3 to 12.5 it is not certain whether the decomposition rate is lower than the metastable equilibrium stabilization rate. Simultaneously, with regard to the high decomposition rate the metastable equilibrium duration time (if it stabilizes in those conditions) should be very short (e.g. shorter than the time of flotation experiment). The ranges of pH values, lower than 3 or greater than 12.5, shown in the diagrams with broken lines are merely a prolongation of diagrams of proper metastable equilibria
357 included in the pH range from 3 to 12.5. They were drawn for the sake of a precise formulation of the problem in works of investigators dealing with kinetics, who could solve the problem if irreversible decomposition is associated with a series of metastable equilibria states or not and possibly determine the duration time of metastable equilibrium.(i.e, the time when the values A [ E t X - ] / [EtX-] in, A [(EtX)2] / [(EtX)2]in, etc., differ only slightly from zero; in the above expressions the activity differences mean the decrement of the substance caused by the irreversible decomposition). If we fix the total activity in a solution [ E t X - ] + [HEtX] + 2 [(EtX)2(aq)] at a value 10 TM mol/dm 3, as was done in the case of the diagram in Fig.l, then the thermodynamically stable solid phases in a system Cu--HEtX--H20, depending on the potential value and the environment pH will be: native copper, cuprous and cupric ethylxanthate, dixanthogen and cupric oxide. The metallic copper already passes into insoluble CuEtX at a potential of - 5 6 0 mV, the plot not being a function of pH, except in the system with a pH lower than 1.52 (cf. above: the remarks on the instability of xanthate solutions}. Only from the potential of - 8 mV does the CuEtX precipitate vanish and in its place a xanthate Cu(EtX)2 with a higher oxidation level is formed according to reaction (26). At the same time dixanthogen, maintaining an oxidation--reduction equilibrium with E t X - ions, reaches its maximum value, corresponding to a saturated solution, and separates itself as a separate phase at the potential of - 1 1 mY, which is very near to the equilibrium potential of CuEtX--Cu(EtX)2. Because of the small difference between the values of potentials, representing the equilibrium state of reactions (26) and (17), both equilibria are represented by one line. The coexistence of substances, taking part in reaction (17), refers to the whole pH range in which the equilibrium line was drawn, i.e. from 1.52 to 16. The equilibrium of reaction (26), however, concerns a slightly smaller range of the hydrogen ion concentration: from pH values of 1.52 to 14.7 at higher potentials in the inert or basic environment, cupric oxide and dixanthogen coexist. The most important pH range from the practical flotation point of view is that of 8 to 12. Essential for the flotation process (as can be seen from the plots) may be the system consisting of phases Cu(EtX)2 -(EtX)2 and CuO--(EtX)2. On the diagrams the line of Cu--Cu(EtX)2 is also marked. This reaction of the oxidation of metallic copper to cupric ethylxanthate precipitating on the electrode surface was observed, among others, in the experiments of Pomianowski and Kowal (1973) if the electrode processes were investigated by those authors in concentrated solution of potassium xanthate 0.1 mol/dm 3. But already at the concentration 10 -3 m01/dm 3 it was no longer possible to isolate the reaction Cu-* Cu(EtX)2. According to the above-quoted paper, the reaction Cu -~ CuEtX, designated as (22) in diagram 1 may be isolated both for the KEtX concentration 0.1 mol/dm 3, and for the concentration 10 -3 mol/dm 3. At the same time, in 0.1M solution, the shape of the potentiometric curve indicates (after Pomia-
358 nowski and Kowal, 1973) a transformation E t X - - ~ (EtX)2. The plot in Fig.1 indicates a possibility of the occurrence of the reaction CuEtX -* (~(EtX):, the equilibrium potential of which is only about 3 mV higher than the redox potential of E t X - / ( E t X ) : (at the accepted thermodynamic data, given in Table I). The electrochemical equilibria diagram presented in Fig.2 concerns the total activity [ E t X - ] + [HEtX]+ 2 [(EtX)2] equal to 10 -4 mol/dm 3. Except for some details and a respective shift of particular lines, the character of the diagram does not differ much from the former one. At very high pH values a new solid phase appears; it is cuprous oxide. The range of its occurrence is limited from three sides by coexistence with native copper, cuprous ethylxanthate (see above: remarks about the instability of water solutions of xanthates) and tenorite. The domain of thermodynamic CuEtX metastability was diminished (the stability areas of CuO and Cu:O were enlarged to its cost) and shifted towards higher potentials, but the domain of Cu(EtX)2 metastability was diminished because of a shift of CuEtX--(EtX)2 equilibrium towards more positive potential values. It is also noteworthy that the equilibrium line of reactions (17) and (26) now got into the domain of copper (II) ions in the whole pH range considered, whereas in the case of solutions with the general activity of the xanthate species equal to 1 0 - ' mol/dm 3 (diagram 1) only above the pH values about 8.5 do the lines of these equilibria get into the domain of relative predominance of cupric complexes: Cu(OH)2, H C u O : - and CuO::-. The lines of equilibria (17) and (18) also came nearer to each other. If we lower still further the activity assumed for all xanthate species in the system, the lines (17) and (18) will cover each other and at still lower total activities of the xanthate species, the phases (EtX)2 and Cu(EtX)2 will no longer form the areas of thermodynamic metastability on the diagram. This situation is represented in a plot (Fig.3). The general activity of dissolved ethylxanthate species, assumed here, is 10 -7 mol/dm a. It is easily seen that in this case there are only four solid phases: Cu, CuEtX, CuO, Cu20. The cupricous xanthate metastability area was shifted towards still more positive potentials, enclosing almost the whole former Cu(EtX)~ area. The oxide phases occupy larger pH and potential areas. The investigations of Pomianowski and Czubak-Pawlikowska (1973b) imply that the metallic copper powder flotation still occurs in solutions of (EtX)2 concentration equal to 2.7.10 -6 mol/dm 3 , which, if the consumption of collector forming a surface c o m p o u n d is taken into account, allegedly confirms the conjecture that the cuprous xanthate formed on the surface of floated grains plays a main role in the flotation process. The adsorption of dixanthogen on copper and other surface phenomena are still an open question. The last diagram in Fig.4 is not a diagram of the Pourbaix type since in its construction it was not assumed that the initial general activity of dissolved xanthate forms is constant. The diagram was made for the following conditions: (1) in the range of potential values lower than those denoted with the equilibrium line (17), (21) the general activity [ EtX-]+ [ HEtX ]is stabi-
359
lized at the level 10 -7 mol/dm 3, whereas the (EtXh(aq) activity varies from lowest values (close to zero) to the saturation of the solution (this being attained at potentials corresponding to the lines (17), (21); (2) in the range of potential values higher than denoted with the equilibrium line (17), (21) the activity of (EtX)2(aq) is stabilized and corresponds to the saturation of the solution, whereas the general activity [ E t X - ] + [HEtX] decreases with the potential increase, in agreement with the reduction--oxidation equilibrium (18), (20). A comparison of plot 4 with diagram 3 shows that at the same activity of E t X - the phase Cu(EtX)2 can be formed (plot 4) or not (diagram 3). This depends on the equilibrium activity (EtX)2(aq). A comparison of the diagrams permits of formulating the problem of experimental solution if a layer of the c o m p o u n d CuEtX formed on the surface of native copper suffices for its floatability. In the case of this diagram, as well as in all those discussed, it was proved that the formation of the metastable areas of Cu(EtX)2 and (EtX)2 coexistence is a thermodynamic property. At the same time the newly formed CuEtX phase should be homogeneous outside the small domain of the metastable coexistence of CuEtX + (EtX)~. These phases are metastable in the potential values region from - 7 0 to - 6 7 mV vs SHE at the standard conditions. It seems that the value of the standard potential of reaction (17) does not exceed the value of - 6 7 mV, because in any other case we could not explain the occurrence of the disproportionation reaction Cu(EtX)2 -~ CuEtX + 1~ (EtX)2. Kakovskii and Arashkevich (1969) give for this reaction the value of AG ° equal to +0.16 kcal/mol. It results from our calculation that this value is negative and e q u a l s - 0 . 0 7 5 kcal/mol. The possibility of the flotation of copper covered with tenorite coexisting with an (EtX): phase in a large range of potentials and environment pH values should also be mentioned. It seems that in the area below line (17) it is possible to check dixanthogen adsorption. We think that the poor results of the oxidated mineral flotation are connected with rather slight adsorption of (EtX)2 on the mineral surface. The phase Cu(EtX)2 formed on mineral surface is rather weakly bonded or does not suffice for flotation. A discussion of the problem of floatability of oxidated minerals is given by Laskowski (1972). Many problems of an experimental nature are connected with the solution for the flotation system Cu--K EtX--H20 presented above. These questions will be discussed in subsequent papers. REFERENCES Abramov, A.A., 1968. Thermodynamic analysis of the interaction between xanthate or dixanthogen and the surface of galena. Tr. Nauchno-Tekh. Konf. Inst. Mekhanobr., 1: 279---293. Abramov, A.A., 1970a. The relationships of the flotation of the sulphide minerals of lead, copper and iron in the presence of the sulphide ion. Obogashch. Rud, 1/2: 66--72. Abramov, A.A., 1970b. On the necessary xanthate concentration in the flotation of copper sulphides. Obogashch. Rud, 3: 26--33.
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Du Rietz, C., 1957. Xanthate analysis by means of potentiometric titration: some chemical properties of the xanthates. Svensk. Kern. Tidskr., 69: 310--327. Finkelstein, N.P., 1967. Kinetic and thermodynamic aspects of the interaction between potassium ethyl xanthate and oxygen in aqueous solution. Trans. Inst. Min. Metall., 76: 51--59. Fuerstenau, M.C., Kuhn, M.C. and Elgillani, D.A., 1968. The role of dixanthogen in xanthate flotation of pyrite. Trans. AIME, 2 4 1 : 1 4 8 156. Garbacik, J., Najbar, J. and Pomianowski, A., 1972. Kinetics of the reaction of xanthate with hydrogen peroxide. Roczn. Chem., 46: 85--97. Gaudin, A.M. and Schuhmann Jr., R., 1936. The action of potassium n-amyt xanthate on chalcocite. J. Phys. Chem., 40: 257--275. Goldstick, T.K., 1959. Electrochemical Investigation of Ethyl Xanthate. Thesis, Massachusetts Inst. of Technology (after Finkelstein, 1967). Granville, A., Finkelstein, N.P. and Allison, S.A., 1972. Review of reactions in the flotation system galena--xanthate--oxygen. Trans. Inst. Min. Metail, Sect. C, 81: 1--30. Hepel, T., 1973. Hydrometallurgical Electroextraction of Copper from Sulphide Minerals with Peculiar Consideration of the Electrochemical Equilibria. Thesis, Jagiellonian University, Krakow (in Polish). Hepel, T. and Pomianowski, A., 1974. On the solubility of cupric oxide and hydroxide in aqueous solutions. Zesz. Nauk. Uniw. Jagiellon., Pr. Chem., 19: 251--261. Hepel, T., Hepel, M. and Pomianowski, A., 1973. Thermodynamics of the flotation systems. Part I: An attempt to the description of the adsorption phases in the system Hg--KEtX--H:O. Physicochem. Probl. Miner. Process., 7: 23--41. Kakovskii, I.A. and Arashkevich, V.M., 1969. The investigation of the properties of anorganic disulphides. In: 8th Int. Congr. Miner. Process, 2. Nauka, Leningrad, p p . 3 0 0 - 3 1 4 Kakovskii, I.A. and Silina, E.I., 1962. Thermodynamic method of investigation of the flotation reagents. Tr. Inst. Uralmekhanobr., 9: 3--47. Laskowski, J., 1972. Flotation of non-sulphide copper minerals. Physicochem. Probl. Miner. Process., 6: 57--63. Leja, J., 1973. Some electrochemical and chemical studies related to froth flotation with xanthates. Miner. Sci. Eng., 5: 278--286. Majima, H., 1969. How oxidation affects selective flotation of complex sulphide ores. Can. Metall. Q., 8: 2 6 9 - 2 7 3 . Majima, H. and Takeda, M., 1968. Electrochemical studies of the xanthate--dixanthogen system on pyrite. Trans. AIME, 241: 431--436. Plaksin, I.N., 1957. Proc. Int. Congr. Surface Activ., 2nd ed. 3 : 3 5 5 (after Woods, 1971). Plaksin, I.N. and Shafeev, R.S., 1959. On the question of the quantitative value xanthates interaction at the dependence surface property of the sulphide minerals. Dokl. Akad. Nauk SSSR, 128: 777--780. Plaksin, I.N. and Shafeev, R.S., 1960. On properties of distribution of xanthate on the surface of sulphide minerals. J. Mines Met. Fuels, 8: 1--8. Pomianowski, A., 1963. Physical chemistry of model flotation of mercury. II. Zesz. Nauk. Uniw. Jagiellon., Pr. Chem., 8: 87--113. Pomianowski, A., 1967. Electrical and surface characteristics in mercury--xanthate--air system. Thesis, Institute of Physical Chemistry, Polish Academy of Science, Krakow. Pomianowski ,A. and Czubak-Pawlikowska, J., 1973. Unpublished data. Pomianowski, A. and Kowal, A., 1973. Electrochemistry of the flotation systems. I. Chronopotentiometric characteristics of mercury, platinum and copper electrodes in an aqueous solutions of potassium ethylxanthate. Zesz. Nauk. Uniw. Jagiellon., Pr. Chem., 18: 2 4 9 - 2 5 9 . Pomianowski, A. and Leja, J., 1963. Spectrophotometric study of xanthate and dixanthogen solutions. Can. J. Chem., 41: 2219--2230.
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Pomianowski, A. and Leja, J., 1964. Equimolar solutions of xanthate and akryl trimethyl ammonium bromide adsorption on copper, nickel and sphalerite powders. Trans. AIME, 229-- 307--312. Sheka, Z.A. and Kriss, E.E., 1959. Metal xanthates. Rab. Khim. Rastvorov Kompleksn. Soedin. Akad. Nauk. Ukr. SSR, 2: 135--162. Shulman, V.M. and Larionov, S.V., 1967. Xanthate oxidation and redox potential of the xanthate--dixanthogen couple. Izv. Sibir. Otd. Akad. Nauk SSSR, SER. Khim. Nauk, 12 35--43. Stepanov. B.A., 1960. Tr. Uralmekhanobra, 7: 33. (after Shuiman, 1967.) Stepanov, B.A. and Zubkov, A.A., 1968. Effect of oxygen on floatability of the metallic mercury. Tsvetn. Met., 8: 30--31. Stepanov, B.A., Kakovskii, I.A. and Serebryakova, N.V., 1959. Nauchn. Dokl. Vyssh. Shk., Khim. Khim. Tekhnol., 2: 277--279. Taggart, A.F. and Gaudin, A.M., 1923. Surface tension and adsorption phenomena in flotation. Trans. AIME, 68: 479--530. Tipman, N.R. and Leja, J., 1975. Reactivity of xanthate and dixanthogen in aqueous solutions of different pH. Colloid. Polym. Sci., 253: 4--10. Tolun, R. and Kitchener, J.A., 1963-64. Electrochemical study of the galena--xanthate-oxygen flotation system. Trans. Inst. Min. Metall., 73: 313--322. Toperi, D. and Tolun, R., 1969. Electrochemical study and thermodynamic equilibria of the galena--oxygen--xanthates flotation system. Trans. Inst. Min. Metall., 78: 191--197. Woods, R., 1971. The oxidation of ethyl xanthate on platinum, gold, copper and galena electrodes. Relation to the meehanizm of mineral flotation. J. Phys. Chem., 75: 354-362. De Zoubov, N., Vanleugenhaghe, C. and Pourbaix, M., 1966. Copper. In: M. Pourbaix (Editor), Atlas of Electrochemical Equilibria in Aqueous Solutions. Pergamon, Cebelcor. pp. 384--392.
DIAGRAMS OF ELECTROCHEMICAL EQUILIBRIA OF THE SYSTEM COPPER--POTASSIUM ETHYLXANTHATE--WATER A T 25 ° C
TADEUSZ HEPEL and ANDRZEJ POMIANOWSKI Department of Hydrometallurgy, Institute of Chemistry, JagieUonian University (Poland) Research Laboratories of Catalysis and Surface Chemistry, Polish Academy of Sciences, Krakow (Poland) (Received June 2, 1976; revised and accepted July 11, 1977)
ABSTRACT Hepel, T., and Pomianowski, A., 1977. Diagrams of electrochemical equilibria of the system copper--potassium ethylxanthate--water at 25 ° C. Int. J. Miner. Process., 4: 345--361. On the basis of thermodynamic calculations, diagrams of metastable electrochemical equilibria of the system copper--potassium ethylxanthate--water, at 25°C were constructed. The equilibria equations and diagrams for the total activity of [EtX-] + [HEtX] + 2 [(EtX)2 ] (equalling 10 -1, 10 -4 and 10 -~ tool/dinS), are presented. The practical conclusions resulting from these diagrams for the flotation of native copper are discussed.
INTRODUCTION A c c o r d i n g t o the generally a c c e p t e d opinions o f Taggart and G a u d i n (1923), it is believed t h a t t o float the grains o f n o n - h y d r o p h o b i c minerals it is necessary and sufficient t o f o r m o n t h e i r surface a c o l l e c t o r coating with a thickness o f a b o u t one m o n o l a y e r . At the same t i m e the m e c h a n i s m o f the interaction o f t h e c o l l e c t o r with the surface o f t h e mineral is e x p l a i n e d in t e r m s o f an e x c h a n g e r e a c t i o n , a d s o r p t i o n o f ions or c o l l e c t o r molecules, or the oxid a t i o n - r e d u c t i o n processes involving t h e o x i d i z e d surface c o m p o u n d s or t h e o x y g e n dissolved in the solution. In some r e c e n t papers the influence o f semic o n d u c t i v e p r o p e r t i e s o f minerals has also been c o n s i d e r e d (Plaksin, 1957; Plaksin and Shafeev, 1959, 1960). It is characteristic t h a t , in spite o f t h e ass u m p t i o n t h a t a m o n o m o l e c u l a r c o l l e c t o r coating is f o r m e d , t h e t h e r m o d y namic d a t a referring t o the m a c r o s c o p i c phases with a definite crystallographic s t r u c t u r e are used in m a n y papers to describe the system. On t h e o t h e r h a n d , it is k n o w n f r o m t h e f o r m e r investigations o f G a u d i n and S c h u h m a n n ( 1 9 3 6 ) t h a t at higher c o l l e c t o r c o n c e n t r a t i o n s several m o l e c u l e coatings can be f o r m e d (the s y s t e m a m y l x a n t h a t e - - c h a l c o c i t e was e x a m i n e d ) . In his electrochemical investigations Woods ( 1 9 7 1 ) p r o v e d t h a t the e t h y l x a n t h a t e ions on p l a t i n u m , gold, c o p p e r and galena e l e c t r o d e s u n d e r g o r e a c t i o n s which are ac-
346
companied by collector adsorption -desorption processes. Similar investigations on mercury were done earlier by Pomianowski (1967) (cf. also Hepel et al., 1973). Recently (Pomianowski and Kowal, 1973) he extended them to copper and chalcocite. The thermodynamic approach to the equilibria attained in flotation systems includes, especially in recent papers (Abramov, 1968, 1970a, b; Majima and Takeda, 1968; Fuerstenau et al., 1968; Toperi and Tolun, 1969; Majima, 1969), not only the "phase" of the c o m p o u n d of the collector with the particular metal, (eg. Pb(C2HsOCS2)2, Cu2(C2HsOCS2h, etc.), but also the surface phases formed by the mineral's oxidation (eg. Cu(OH)2, Cu2(OH)2CO3, Fe(OH)3, etc.) which take place when the mineral reacts with the collector solution, or earlier in the geochemical processes or during the extracting of the ore. In the present paper (cf. also: Hepel, 1973) the authors try to describe the system metallic copper--potassium xanthate aqueous solutions. Because of the complicated behaviour of the sulphide minerals in the flotation process conditions, a complete description of sulphide systems at this stage of the investigations was given up and a more careful analysis of a simpler case -native copper flotation -- was undertaken. Only the phases, the properties of which correspond to the macroscopic phases, were considered. THE CONSTRUCTION OF DIAGRAMS
The approach offered in the present paper is based on the application of electrochemical equilibria diagrams to investigate flotation systems (Fuerstenau et al., 1968; Hepel, 1973). The constructed diagrams refer to metastable equilibria. Since for the system native copper--ethylxanthate solution there only exist comparative thermodynamic data, a calculation method enabling clear and easy operations was used; it allowed the equilibria to be calculated for more than fifty reactions, occurring in the Cu--HEtX--H20 system, which have to be considered to construct the electrochemical equilibria diagram. To estimate the values of individual standard potentials and equilibrium constants of chemical reactions, pairs of substances were compiled, as it is impossible to calculate the free energy of formation at 25 ° C (298 ° K) of substances containing the E t X - group, and on the basis of accessible information the differences in changes of the free energy of formation for each of these pairs was estimated (Hepel, 1973). Some values of the differences Ag ° (298) for some of the formed pairs are presented in Table I as examples. After having analyzed the results of potentiometric investigations, and after having calculated them, a v a l u e - 3 . 2 2 9 kcal/mot was accepted for further calculations for the pair (EtX)2(s)--2 EtX-. This corresponds to the normal oxidation potential of the ethyIxanthate ions to dixanthogen in a solution saturated with the oxidation product. For the normal potential of this reaction, a value - 0 . 0 7 0 V with respect to the standard hydrogen electrode was assumed. According to experimental data, its value ranges within rather large limits: Stepanov et al. (1959) and Kakovskii and Arashkevich (1969) give the
347
TABLE I The thermodynamic data for the Cu--KEtX--H20 system Compounds pairs
Difference AUo (29.) kcal mo1-1
J moi-' "10 -3
(EtX)~(s)--2 EtX(EtX)2(aq)--2 EtX-
- 3.229 + 3.454
- 13.52 + 14.46
(EtX)2(s)--(EtX)2(aq)
- 6.683
- 27.97
CuEtX--EtXCu(EtX)2--2 EtX-
-14.28 -15.82
- 59.78 - 66.22
Cu(EtX)2--(EtX)2(s)
-12.59
- 52.70
Cu(EtX)~--(EtX)2(aq )
-19.27
- 80.66
2 CuEtX--(EtX)2(s)
-25.33
-106.0
2 CuEtX--(EtX)2(aq)
-32.01
-134.00
Cu(EtX)2--2 CuEtX
+12.74
+ 53.33
HEtX--EtX-
- 2.076
-
8.69
Source
see text Leja (1973); Tipman and Leja (1975)*, ** Leja (1973); Tipman and Leja (1975)* Sheka and Kriss (1959)* Kakovskii and Arashkevich (1969)*, ** Kakovskii and Arashkevich (1969)*, 8 ** Kakovskii and Arashkevich (1969); Leja (1973); Tipman and Leja (1975)*, ** Sheka and Kriss (1959); Abramov (1968)*, ** Abramov (1968); Leja (1973); Tipman and Leja (1975)*, ** Abramov (1968); Kakovskii and Arashkevich (1969)* Kakovskii and Silina (1962)*
*Computed from data. **For the difference (EtX)2(s)-2 EtX-, a value of ~uo(29s): -3.229 kcal/mol assumed in this work was used for calculations. value - 0 . 0 3 7 V, Majima a n d T a k e d a ( 1 9 6 8 ) t h e value - 0 . 0 4 9 V, G o l d s t i c k (1959); - 0 . 0 5 3 V, Du Rietz ( 1 9 5 7 ) : - 0 . 0 6 7 V a n d T o l u n a n d K i t c h e n e r (1963): - 0 . 0 8 1 V; finally t h e e s t i m a t i o n b y F i n k e l s t e i n ( 1 9 6 7 ) f r o m spectrop h o t o m e t r i c investigations is t h a t t h e s t a o d a r d p o t e n t i a l o f t h e ( E t X ) 2 - - E t X s y s t e m ranges f r o m - 0 . 0 8 0 t o - 0 . 1 4 0 V. We a s s u m e d a r a t h e r l o w e r value o f t h e s t a n d a r d p o t e n t i a l because: (1) activity o f e l e c t r o c h e m i c a l l y d e p o s i t e d (EtX)2 on t h e s u r f a c e o f t h e e l e c t r o d e is r a t h e r greater t h a n u n i t y ; (2) activity of E t X - ions is l o w e r t h a n t h e i r c o n c e n t r a t i o n in t h e s o l u t i o n s applying. To construct the diagrams for a Cu--KEtX--H20 system, the new solutions f o u n d f o r a s y s t e m C u - - H ~ O 4 - - H 2 0 ( H e p e l a n d P o m i a n o w s k i , 1 9 7 4 ) were used; t h e y d i f f e r f r o m t h e o n e s o b t a i n e d b y De Z o u b o v et al. {1966) in considering t h e stable c o m p l e x Cu(OH)2 (aq). This caused a change in t h e s y s t e m o f equilibria b e t w e e n t h e solid phases CuO, C u 2 0 a n d Cu a n d t h e s o l u t i o n o f dissolved c o p p e r . T h e e q u a t i o n o f t h e c h e m i c a l a n d e l e c t r o c h e m i c a l reactions, c o n s i d e r e d during c o n s t r u c t i n g t h e d i a g r a m o f p o t e n t i a l vs p H f o r C u - - K E t X - - H ~ O s y s t e m at 25 ° C, a n d t h e equilibria e q u a t i o n s are r e p r e s e n t e d b e l o w .
348
Equations o f chemical and electrochemical reactions and equilibria equations for a Cu - K E t X - H 2 0 system at 25 ° C (a) T h e r e a c t i o n s n o t i n v o l v i n g t h e E t X - - s p e c i e s : Cu2+ + e ~ Cu+
[Cu 2÷] E, = 0 . 1 5 3 + 0 . 0 5 9 1 l o g
(1)
.....
[Cu +1
Cu(OH)2(aq) + 2H + + e~Cu + + 2 H20
[Cu (OH):] ....... - 0.1182 pH
E2 = 0 . 9 6 2 + 0 . 0 5 9 1 l o g
(2)
[Cu+l
Cu2+ + 2 H 2 0 #Cu (OH)2(aq) + 2 H+
pH = 6.84 + 0.5 log
(3)
[Cu (OH)2] ..............
[Cu2+1
Cu (OH)2(aq) ~- HCuOz- + H +
(4)
[HCuO2-] pH = 13.03 + log [Cu (OHhl HCuO2- ~ CuO: 2- + H +
pH = 13.14 + log
(5)
[CuOz2-1 ........
[Cu (OH)2]
Cu (OH)g(aq) + 2 H ÷ + 2 e ~ C u + 2 H 2 0
(6) E6 = 0 . 7 4 1 + 0 . 0 2 9 5 l o g [ C u ( O H ) 2 ] - 0 . 0 5 9 1 p H HCuO2- + 3 H + + 2 e-~ Cu + 2 H20
(7) E7 = 1 . 1 2 7 + 0 . 0 2 9 5 l o g [ H C u O 2 - ] CuO2 g - + 4 H + + 2 e ~ C u
- 0.0886 pH
+ 2 H20
(8) Es = 1 . 5 1 5 + 0 . 0 2 9 5 l o g [ C u O 2 2 - ] - 0 . 1 1 8 2 p H 1/2 C u 2 0 + H
÷+e~Cu+'~
H20
(9) E9 = 0 . 4 7 1 - 0 . 0 5 9 1 p H
349
CuO + H + + e # 1A C u 2 0 + '/~ H 2 0
(10) E,o = 0 . 6 9 7 - 0 . 0 5 9 1 p H 2 Cu (OH)2 (aq) + 2 H + + 2 e ~ C u 2 0 + 3 H 2 0 (11) El, = 1 . 0 1 2 + 0 . 0 5 9 1 log [Cu ( O H ) 2 ] - 0 . 0 5 9 1 pH 2 HCuO2- + 4 H + + 2 e # Cu20 + 3 H20 (12) E,2 = 1 . 7 8 2 + 0 . 0 5 9 1 log [ H C u O 2 - ] - 0 . 1 1 8 2 p H 2 CuO22- + 6 H ÷ + 2 e # C u 2 0 + 3 H20 (13) E,3 = 2 . 5 5 9 + 0 . 0 5 9 1 log [CuO22-] - 0 . 1 7 7 3 p H CuO22- + 2 H + # CuO + H 2 0 (14) p H = 1 5 . 7 5 + 0.5 log [CuO22-] CuO + 2 H ÷~ Cu2+ +
H20 (15)
pH = 4 . 1 8 - 0.5 log [Cu 2+] Cu ÷ + e ~- Cu (16)
E,6 = 0.520 + 0.0591 log [Cu+] (b) The reactions of KEtX: (EtX)2(s) + 2 e ~ 2 EtX(17)
E,7 = - 0 . 0 7 0 - 0.0591 log [EtX-] (EtX)2(aq) + 2 e ~ 2 E t X E,8 = 0 . 0 7 5 + 0 . 0 2 9 5 log
(18)
[(EtX)21 [EtX-] 2
HEtX ~ EtX- + H ÷ p H = 1.52 + log
(19)
[EtX-] [SStX]
(EtX)2(aq) + 2 H ÷ + 2 e ~ 2 H E t X E2o = 0 . 1 6 5 - 0 . 0 2 9 5 log
[(EtX)2I [HEtX] 2
(20) - 0.0591 pH
350
(EtX)2(s) + 2 H + + 2 e ¢ 2 H E t X
(21) E2, = 0 . 0 2 0 -
0.0591 log [HEtX]-
0.0591 pH
(c) T h e r e a c t i o n s i n v o l v i n g c o p p e r a n d x a n t h a t e : CuEtX + e ~ Cu + EtX(22) E2: -- - 0 . 6 1 9
- 0.0591 log [EtX-]
CuEtX + H ÷ + e ~Cu + HEtX (23) E23 = - 0 . 5 2 9
- 0.0591 log [HEtX]-
0.0591 pH
Cu(EtX)2 + 2 e ~-Cu + 2 EtX(24) E24 = - 0 . 3 4 3
- 0.0591 log [EtX-]
Cu(EtX)2 + H ÷ + 2 e ~ Cu + 2 HEtX (25) E2s = - 0 . 2 5 3
- 0.0591 log [HEtX]-
0.0591 pH
Cu ( E t X ) 2 + e ~ C u E t X + E t X (26) E26 = - 0 . 0 6 7
-
0.0591 log [EtX-]
Cu ( E t X ) 2 + H + + e # C u E t X + H E t X (27) E27 = 0 . 0 2 3 - 0 . 0 5 9 1 l o g [ H E t X ] 2CuEtX+H20
- 0.0591 pH
~Cu20+2EtX-+2H
+ (28)
pH = 18.44 + log [EtX-] CuO + EtX- + 2 H + + e ~CuEtX
+ H20
(29) E2~ = 1 . 7 8 7 + 0 . 0 5 9 1 l o g [ E t X - ] -
0.1182 pH
(EtX)2(s) + C u O + 2 H ÷ + 2 e ~ - C u ( E t X ) 2 + H 2 0 (30) E30 = 0 . 8 5 7 - 0 . 0 5 9 1 p H (EtX)2(s) + C u 2+ 2 e ~ C u ( E t X ) 2 (31)
E3, = 0 . 6 1 0 + 0 . 0 2 9 5 l o g [ C u :+]
351
(EtX)2(s) + Cu (OH)2(aq) + 2 H ÷ + 2 e ~ C u ( E t X ) 2 + 2 H 2 0
(32)
E32 = 1 . 0 1 4 + 0 . 0 2 9 5 log [Cu ( O H ) 2 ] - 0 . 0 5 9 1 p H
Cu(OH)2(aq) + E t X - + 2 H + + e ~- C u E t X + 2 H 2 0 (34) E33 = 2 . 1 0 2 + 0 . 0 5 9 1 log [ E t X - ] + 0 . 0 5 9 1 log [ C u ( O H ) 2 ] - 0 . 1 1 8 2 p H Cu :+ + E t X - + e ~ C u E t X (34) E34 = 1 . 2 9 3 + 0 . 0 5 9 1 log [Cu2÷]+ 0 . 0 5 9 1 log [ E t X - ] Cu 2+ + H E t X + e ~ C u E t X + I-I+
(35) E3s = 1 . 2 0 2 + 0 . 0 5 9 1 log [Cu 2+] + 0 . 0 5 9 1 log [HEtX]+ 0 . 0 5 9 1 p H (EtX)2(aq) + 2 C u O + 4 H + + 4 e ~ 2 C u E t X + 2 H 2 0 (36) E36 = 0 . 9 3 1 + 0 . 0 1 4 7 log [ ( E t X ) 2 ] - 0 . 0 5 9 1 p H Cu + + (EtX)2(s) + e ~ Cu (EtX)2 (37) E37 = 1 . 0 6 6 + 0 . 0 5 9 1 log [Cu +] 2 H C u O 2 - + (EtX)2(s) + 6 H + + 4 e # 2 C u E t X + 4 H 2 0
(38) E38 = 1.401 + 0 . 0 2 9 5 log [ H C u O 2 - ] - 0 . 0 8 8 6 p H 2 CuO22- + (EtX)2(s) + 8 H ÷ + 4 e ~ 2 C u E t X + 4 H 2 0
(39) Ea9 = 1 . 7 9 0 + 0 . 0 2 9 5 log [ C U O 2 2 - ] - 0 . 1 1 8 2 p H HCuO2- + EtX- + 3 H ÷ + e # CuEtX + 2 H20
(40) F~0 = 2 . 8 7 2 + 0 . 0 5 9 1 log [ E t X - ] + 0 . 0 5 9 1 log [ H C u O : - ] - 0 . 1 7 7 3 p H CuO22- + E t X - + 4 H + + e ~ C u E t X + 2 H 2 0 (41) E4, = 3 . 6 4 9 + 0 . 0 5 9 1 log [CuO22-] + 0 . 0 5 9 1 log [ E t X - ] - 0 . 2 3 6 4
pH
Cu + + H E t X ~ C u E t X + H +
(42) p H = - 1 7 . 7 5 - log [ H E t X ] - log [Cu +]
352
2 Cu 2+ + (EtX)2(aq) + 4 e ~ 2 C u E t X ( 43 i E43 = 0 . 6 8 4 + 0 . 0 2 9 5 l o g l C u > ] +
0.0147 log [(EtX)2]
2 Cu(OH)2(aq) + (EtX)2(aq) + 4 H ÷ + 4 e = 2 C u E t X + 4 H 2 0 {44) E44 = 1 . 0 8 8 + 0 . 0 2 9 5 l o g [ C u ( O H ) 2 ] + 0 . 0 1 4 7 l o g [ ( E t X ) 2 ] - 0 . 0 5 9 1 p H
~
q
600-
400
CuO (EtX)2
CufEtX)2 @~
.
(EfX)2 rE, X)2,.:.,,
':
200 ~
~(? "t "
E CO
~_
1~
CuE~X
~ -200~
%
Q~ -LO0~
-500,
Cu 2
4
6
8
10
12
1L
pH
16
Fig.1. Diagram of the metastable electrochemical equilibria for the system Cu--KEtX--H20 at 25 ° C. The initial general activity [ E t X - ] + [ H E t X ] + 2 [(EtX)2(aq) ] equals 0.1 tool/din 3. (The additional lines (24) and (25) in the diagram are shown only for comparison with results of some experiments. )
DISCUSSION OF RESULTS AND CONCLUSIONS
The most developed system for a thermodynamic approach to the flotation process was recently presented by Abramov (1968, 1970a, b). The thermodynamic equilibria considered by this author are not stable. This fact is noteworthy, as in describing the metastable equilibrium state it is necessary to regard and confirm the assumption concerning the kinetic nature of the system, and especially the restrictions of the processes leading to the establishment of stable equilibria. Simultaneously, the considered metastable
353
[ fEfXJzCu:"
~
600 F
S CuO -'EtX).
Cu (EtX)2 ~ {E~X)2
400 ~HEIX
"
~,
..b =oo
,=
E
,,,,=
,'
CU +
"tu3
~p.~
CuO
c
o -b.. 4
.._r
"'2
CuEtX
i~
~-200
%a
~Cu20 -400'-
@
""9..
Cu i
i
.:
(~)
-600 0
2
4
6
8
10
12
14
pH
i6
Fig.2. Potential--pH metastable equilibria diagram for the system Cu--KEtX--H20 at 25 ° C. The initial general activity [ E t X - ] + [HEtX] + 2 [(EtX)2(aq)] equals 10 -4 tool/din 3. (The additional lines (24) and (25) in the diagram are shown only for comparison with results of some experiments. )
equilibria must be chosen based on experimental indications, proving that they are of practical importance, i.e. that the forward and backward processes run with sufficient velocity. Another procedure suggested (Hepel, 1973} is also possible; it assumes the course of particular reactions and then by means of thermodynamic analysis estimates the parameters describing the particular states of metastable and stable equilibria and compares the results with the experimental observations. This analysis is best demonstrated when represented in graphic form. Abramov (1968, 1970a, b) solved a system of equations describing chemical and electrochemical equilibria reactions and balanced the equations, thus arriving at a number of interesting conclusions concerning the influence of different reagents competition on the flotation range of lead, copper and iron sulphide minerals. The authors of the present paper believe that the way in which Abramov selected the chemical and electrochemical equations describing the processes in real systems is too arbitrary. In another paper, H e p e l e t al. (1973) tried to represent on the Lange--Nagel diagram the surface adsorption phases formed in the aqueous mercury--ethylxanthate system. The results of earlier experimental investigations, carried
3 5 `4,
+,,
t,'J"
_
600:
.
....
_
CuO &O0:
(EfX~
, e
HEfX -~gf io
.
Cu{Of#.
EtX -
CU"
2 0 0 '~
,~
HCuO~
CuEfX
-.\
\23 +"
.
Cu
-400
'"
- 600 0
2
4
6
8
10
12
14
pH
15
Fig. 3. Diagram of the metastable electrochemical equilibria for the system Cu--KEtX--H20 at 25 ° C. The initial general activity [ E t X - ] + [ H E t X ] + 2[(EtX)2(aq ) ]equals 10 -7 tool/din 3. (The additional lines (24) and (25) in the diagram are shown only for comparison with results of some experiments. )
out by means of polarography and voltammetry (Pomianowski, 1963, 1967), were used to construct the diagram. The solution of the system of the adsorption phases equilibria was only a first attempt to describe the flotation system and so it does not pretend to be a definite one. This is due to the lack of some experimental data and on the other hand to interpretative troubles, especially in the systems having low xanthate ion concentrations. In the present thermodynamic description of the systems occurring in the flotation process, the adsorption phenomena were not taken into account because of the lack of data concerning the free adsorption energy of individual substances existing in the system. Accordingly, the results presented above may be directly applied to the flotation process only if it is known that on the surface of the flotation minerals coatings are formed of hydrophobic compounds such as: xanthate--the mineral involved (here copper), characterized by the values corresponding to the macroscopic phases of these compounds. On the other hand, when the data concerning the surface adsorption phases are obtained, the diagrams presented here may be completed and adapted for a direct interpretation of the flotation process. The diagrams of potential vs pH worked out on the basis of the above
355
600~-
~ ~9
qb
~3,
(EtX)2
' E
-"-"
~ -- - - - 4 7
i
CuO
,, Cu {EfX)2
(EtX)2
~ .I
HCuO= >
~ "~,
CuEtX
CuO
",,~
I
5
. -200
.
. . . . .
i
-4oo~
@
..Cu..
~)
-600 L ~ L I ! ~ 0
\ ...
2
X I~
\
x,
_~
~
1
4
6
~
t 8
,
• 12
1[,
pH
16
Fig.4. Potential--pH plot of the metastable electrochemical equilibria for the system Cu--KEtX--H~O at 25 ° C. General activity [ EtX- ] + [ HEtX ] equals 10-7 tool/din 3 below lines (17) and (21) and decreases at higher potential values. Activity of (EtX)2(aq) is constant above lines (17) and (21) and equals 1.25"10-s tool/din 3 and decreases at lower potential values.
equilibria reactions are presented in Figs.l--4. The first three plots correspond to the total xanthate forms activity in the solutions: 10 -1, 10 TM, 10 -7 mol/ dm 3, respectively. On all diagrams a dash-dotted line marks the areas of the relative predominance of the individual species of ethylxanthate dissolved in water. They are: E t X - ions dominating in the whole range of pH above the value 1.52 and at sufficiently low potentials, then the undissociated ethylxanthic acid HEtX existing in relative predominance in strongly acid solutions*, and dixanthogen, towards which the solution quickly becomes saturated because of its low solubility (equal to 7.9.10 -6 mol/dm 3 (Stepanov, 1960), 1.10 -s mol/dm 3 (Kakovskii et al., 1969), 1.3.10 -s mol/dm 3 (Pomianowski and Leja, 1963) and 1.25-10 -s mol/dm 3 (Leja, 1973; Tipman and Leja, 1975)}. The (EtX)~ relative predominance area occurs at higher potential values. The stability of aqueous xanthate solutions is limited by a number of conditions: the solution decomposes irreversibly at pH values lower than about 6 and greater than about 10.5, the solution is likely to be oxidized *cf. lower: the remarks on the instability of KEtX solutions.
356
(Pomianowski and Leja, 1963; Finkelstein, 1967; Shulman and Larionov, 1967; Garbacik et al., 1972) and the limiting solution concentrations is about, 0.1 tool KEtX/dm 3. in acid solutions there occurs a decomposition caused by a large shift of the equilibrium of reaction (19) towards the formation of undissociated HEtX, which undergoes an irreversible reaction, forming carbon disulphide and ethyl alcohol (Pomianowski and Leja, 1963b; Shulman and Larionov, 1967). In the alkaline environment, the decomposition of EtX-ions leads to a thiocarbonic acid salt, sulphides, disulphides, carbonates and others (Pomianowski, 1963; Finkelstein, 1967). Simultaneously, due to the --SH functional group, strong reduction abilities are characteristic for xanthates. Under these conditions it was necessary to introduce some new designations on the Pourbaix diagrams, worked out for the metastable equilibria. All the lines, concerning the chemical and electrochemical equilibria reactions occurring in the area relatively dominated by an unstable ethyloxanthate acid and "near" to this area are drawn as broken lines. The pH value of 3 is conventionally accepted as the value limiting the metastability range of E t X ions from the acid part of the pH scale, at which (as can be seen, among others, from Finkelstein, 1967), the half-life of xanthate is 5 minutes. This is sufficiently long and a flotation experiment under such conditions may be considered. The unstability of xanthate solutions in strongly alkali environment having a pH higher than about 12.5 is also marked by a broken line (Leja, 1973). Generally speaking, we assume that all substances containing the xanthate group are not stable in the absence of CS2 and EtOH (Leja, 1973; Tipman and Leja, 1975) in proper concentrations and carbonate species (at the domain of thermodynamic stability of water) that could guarantee attaining the stable equilibrium. The rate of irreversible decomposition of particular xanthate forms varies, however, in dependence on the pH value of the solution. It is conventionally assumed that in the pH range from a b o u t 3 to about 12.5 the rate of irreversible decomposition is considerably lower than the rate of processes proceeding among the solution components: EtX-, (EtX)2(aq) and solids: CuEtX, Cu(EtX)2, (EtX)2, and is n o t associated with the xanthate group decomposition. This permits consideration of metastable equilibria connected with these processes. Since a slow xanthate decomposition process proceeds parallelly, the given metastabte equilibrium can only be spoken of at the time period when the decrement of xanthate activity caused by decomposition is negligibly small. The higher the rate of the irreversible decomposition process, the shorter the duration time of metastable equilibrium. Outside the pH range from 3 to 12.5 it is not certain whether the decomposition rate is lower than the metastable equilibrium stabilization rate. Simultaneously, with regard to the high decomposition rate the metastable equilibrium duration time (if it stabilizes in those conditions) should be very short (e.g. shorter than the time of flotation experiment). The ranges of pH values, lower than 3 or greater than 12.5, shown in the diagrams with broken lines are merely a prolongation of diagrams of proper metastable equilibria
357 included in the pH range from 3 to 12.5. They were drawn for the sake of a precise formulation of the problem in works of investigators dealing with kinetics, who could solve the problem if irreversible decomposition is associated with a series of metastable equilibria states or not and possibly determine the duration time of metastable equilibrium.(i.e, the time when the values A [ E t X - ] / [EtX-] in, A [(EtX)2] / [(EtX)2]in, etc., differ only slightly from zero; in the above expressions the activity differences mean the decrement of the substance caused by the irreversible decomposition). If we fix the total activity in a solution [ E t X - ] + [HEtX] + 2 [(EtX)2(aq)] at a value 10 TM mol/dm 3, as was done in the case of the diagram in Fig.l, then the thermodynamically stable solid phases in a system Cu--HEtX--H20, depending on the potential value and the environment pH will be: native copper, cuprous and cupric ethylxanthate, dixanthogen and cupric oxide. The metallic copper already passes into insoluble CuEtX at a potential of - 5 6 0 mV, the plot not being a function of pH, except in the system with a pH lower than 1.52 (cf. above: the remarks on the instability of xanthate solutions}. Only from the potential of - 8 mV does the CuEtX precipitate vanish and in its place a xanthate Cu(EtX)2 with a higher oxidation level is formed according to reaction (26). At the same time dixanthogen, maintaining an oxidation--reduction equilibrium with E t X - ions, reaches its maximum value, corresponding to a saturated solution, and separates itself as a separate phase at the potential of - 1 1 mY, which is very near to the equilibrium potential of CuEtX--Cu(EtX)2. Because of the small difference between the values of potentials, representing the equilibrium state of reactions (26) and (17), both equilibria are represented by one line. The coexistence of substances, taking part in reaction (17), refers to the whole pH range in which the equilibrium line was drawn, i.e. from 1.52 to 16. The equilibrium of reaction (26), however, concerns a slightly smaller range of the hydrogen ion concentration: from pH values of 1.52 to 14.7 at higher potentials in the inert or basic environment, cupric oxide and dixanthogen coexist. The most important pH range from the practical flotation point of view is that of 8 to 12. Essential for the flotation process (as can be seen from the plots) may be the system consisting of phases Cu(EtX)2 -(EtX)2 and CuO--(EtX)2. On the diagrams the line of Cu--Cu(EtX)2 is also marked. This reaction of the oxidation of metallic copper to cupric ethylxanthate precipitating on the electrode surface was observed, among others, in the experiments of Pomianowski and Kowal (1973) if the electrode processes were investigated by those authors in concentrated solution of potassium xanthate 0.1 mol/dm 3. But already at the concentration 10 -3 m01/dm 3 it was no longer possible to isolate the reaction Cu-* Cu(EtX)2. According to the above-quoted paper, the reaction Cu -~ CuEtX, designated as (22) in diagram 1 may be isolated both for the KEtX concentration 0.1 mol/dm 3, and for the concentration 10 -3 mol/dm 3. At the same time, in 0.1M solution, the shape of the potentiometric curve indicates (after Pomia-
358 nowski and Kowal, 1973) a transformation E t X - - ~ (EtX)2. The plot in Fig.1 indicates a possibility of the occurrence of the reaction CuEtX -* (~(EtX):, the equilibrium potential of which is only about 3 mV higher than the redox potential of E t X - / ( E t X ) : (at the accepted thermodynamic data, given in Table I). The electrochemical equilibria diagram presented in Fig.2 concerns the total activity [ E t X - ] + [HEtX]+ 2 [(EtX)2] equal to 10 -4 mol/dm 3. Except for some details and a respective shift of particular lines, the character of the diagram does not differ much from the former one. At very high pH values a new solid phase appears; it is cuprous oxide. The range of its occurrence is limited from three sides by coexistence with native copper, cuprous ethylxanthate (see above: remarks about the instability of water solutions of xanthates) and tenorite. The domain of thermodynamic CuEtX metastability was diminished (the stability areas of CuO and Cu:O were enlarged to its cost) and shifted towards higher potentials, but the domain of Cu(EtX)2 metastability was diminished because of a shift of CuEtX--(EtX)2 equilibrium towards more positive potential values. It is also noteworthy that the equilibrium line of reactions (17) and (26) now got into the domain of copper (II) ions in the whole pH range considered, whereas in the case of solutions with the general activity of the xanthate species equal to 1 0 - ' mol/dm 3 (diagram 1) only above the pH values about 8.5 do the lines of these equilibria get into the domain of relative predominance of cupric complexes: Cu(OH)2, H C u O : - and CuO::-. The lines of equilibria (17) and (18) also came nearer to each other. If we lower still further the activity assumed for all xanthate species in the system, the lines (17) and (18) will cover each other and at still lower total activities of the xanthate species, the phases (EtX)2 and Cu(EtX)2 will no longer form the areas of thermodynamic metastability on the diagram. This situation is represented in a plot (Fig.3). The general activity of dissolved ethylxanthate species, assumed here, is 10 -7 mol/dm a. It is easily seen that in this case there are only four solid phases: Cu, CuEtX, CuO, Cu20. The cupricous xanthate metastability area was shifted towards still more positive potentials, enclosing almost the whole former Cu(EtX)~ area. The oxide phases occupy larger pH and potential areas. The investigations of Pomianowski and Czubak-Pawlikowska (1973b) imply that the metallic copper powder flotation still occurs in solutions of (EtX)2 concentration equal to 2.7.10 -6 mol/dm 3 , which, if the consumption of collector forming a surface c o m p o u n d is taken into account, allegedly confirms the conjecture that the cuprous xanthate formed on the surface of floated grains plays a main role in the flotation process. The adsorption of dixanthogen on copper and other surface phenomena are still an open question. The last diagram in Fig.4 is not a diagram of the Pourbaix type since in its construction it was not assumed that the initial general activity of dissolved xanthate forms is constant. The diagram was made for the following conditions: (1) in the range of potential values lower than those denoted with the equilibrium line (17), (21) the general activity [ EtX-]+ [ HEtX ]is stabi-
359
lized at the level 10 -7 mol/dm 3, whereas the (EtXh(aq) activity varies from lowest values (close to zero) to the saturation of the solution (this being attained at potentials corresponding to the lines (17), (21); (2) in the range of potential values higher than denoted with the equilibrium line (17), (21) the activity of (EtX)2(aq) is stabilized and corresponds to the saturation of the solution, whereas the general activity [ E t X - ] + [HEtX] decreases with the potential increase, in agreement with the reduction--oxidation equilibrium (18), (20). A comparison of plot 4 with diagram 3 shows that at the same activity of E t X - the phase Cu(EtX)2 can be formed (plot 4) or not (diagram 3). This depends on the equilibrium activity (EtX)2(aq). A comparison of the diagrams permits of formulating the problem of experimental solution if a layer of the c o m p o u n d CuEtX formed on the surface of native copper suffices for its floatability. In the case of this diagram, as well as in all those discussed, it was proved that the formation of the metastable areas of Cu(EtX)2 and (EtX)2 coexistence is a thermodynamic property. At the same time the newly formed CuEtX phase should be homogeneous outside the small domain of the metastable coexistence of CuEtX + (EtX)~. These phases are metastable in the potential values region from - 7 0 to - 6 7 mV vs SHE at the standard conditions. It seems that the value of the standard potential of reaction (17) does not exceed the value of - 6 7 mV, because in any other case we could not explain the occurrence of the disproportionation reaction Cu(EtX)2 -~ CuEtX + 1~ (EtX)2. Kakovskii and Arashkevich (1969) give for this reaction the value of AG ° equal to +0.16 kcal/mol. It results from our calculation that this value is negative and e q u a l s - 0 . 0 7 5 kcal/mol. The possibility of the flotation of copper covered with tenorite coexisting with an (EtX): phase in a large range of potentials and environment pH values should also be mentioned. It seems that in the area below line (17) it is possible to check dixanthogen adsorption. We think that the poor results of the oxidated mineral flotation are connected with rather slight adsorption of (EtX)2 on the mineral surface. The phase Cu(EtX)2 formed on mineral surface is rather weakly bonded or does not suffice for flotation. A discussion of the problem of floatability of oxidated minerals is given by Laskowski (1972). Many problems of an experimental nature are connected with the solution for the flotation system Cu--K EtX--H20 presented above. These questions will be discussed in subsequent papers. REFERENCES Abramov, A.A., 1968. Thermodynamic analysis of the interaction between xanthate or dixanthogen and the surface of galena. Tr. Nauchno-Tekh. Konf. Inst. Mekhanobr., 1: 279---293. Abramov, A.A., 1970a. The relationships of the flotation of the sulphide minerals of lead, copper and iron in the presence of the sulphide ion. Obogashch. Rud, 1/2: 66--72. Abramov, A.A., 1970b. On the necessary xanthate concentration in the flotation of copper sulphides. Obogashch. Rud, 3: 26--33.
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Du Rietz, C., 1957. Xanthate analysis by means of potentiometric titration: some chemical properties of the xanthates. Svensk. Kern. Tidskr., 69: 310--327. Finkelstein, N.P., 1967. Kinetic and thermodynamic aspects of the interaction between potassium ethyl xanthate and oxygen in aqueous solution. Trans. Inst. Min. Metall., 76: 51--59. Fuerstenau, M.C., Kuhn, M.C. and Elgillani, D.A., 1968. The role of dixanthogen in xanthate flotation of pyrite. Trans. AIME, 2 4 1 : 1 4 8 156. Garbacik, J., Najbar, J. and Pomianowski, A., 1972. Kinetics of the reaction of xanthate with hydrogen peroxide. Roczn. Chem., 46: 85--97. Gaudin, A.M. and Schuhmann Jr., R., 1936. The action of potassium n-amyt xanthate on chalcocite. J. Phys. Chem., 40: 257--275. Goldstick, T.K., 1959. Electrochemical Investigation of Ethyl Xanthate. Thesis, Massachusetts Inst. of Technology (after Finkelstein, 1967). Granville, A., Finkelstein, N.P. and Allison, S.A., 1972. Review of reactions in the flotation system galena--xanthate--oxygen. Trans. Inst. Min. Metail, Sect. C, 81: 1--30. Hepel, T., 1973. Hydrometallurgical Electroextraction of Copper from Sulphide Minerals with Peculiar Consideration of the Electrochemical Equilibria. Thesis, Jagiellonian University, Krakow (in Polish). Hepel, T. and Pomianowski, A., 1974. On the solubility of cupric oxide and hydroxide in aqueous solutions. Zesz. Nauk. Uniw. Jagiellon., Pr. Chem., 19: 251--261. Hepel, T., Hepel, M. and Pomianowski, A., 1973. Thermodynamics of the flotation systems. Part I: An attempt to the description of the adsorption phases in the system Hg--KEtX--H:O. Physicochem. Probl. Miner. Process., 7: 23--41. Kakovskii, I.A. and Arashkevich, V.M., 1969. The investigation of the properties of anorganic disulphides. In: 8th Int. Congr. Miner. Process, 2. Nauka, Leningrad, p p . 3 0 0 - 3 1 4 Kakovskii, I.A. and Silina, E.I., 1962. Thermodynamic method of investigation of the flotation reagents. Tr. Inst. Uralmekhanobr., 9: 3--47. Laskowski, J., 1972. Flotation of non-sulphide copper minerals. Physicochem. Probl. Miner. Process., 6: 57--63. Leja, J., 1973. Some electrochemical and chemical studies related to froth flotation with xanthates. Miner. Sci. Eng., 5: 278--286. Majima, H., 1969. How oxidation affects selective flotation of complex sulphide ores. Can. Metall. Q., 8: 2 6 9 - 2 7 3 . Majima, H. and Takeda, M., 1968. Electrochemical studies of the xanthate--dixanthogen system on pyrite. Trans. AIME, 241: 431--436. Plaksin, I.N., 1957. Proc. Int. Congr. Surface Activ., 2nd ed. 3 : 3 5 5 (after Woods, 1971). Plaksin, I.N. and Shafeev, R.S., 1959. On the question of the quantitative value xanthates interaction at the dependence surface property of the sulphide minerals. Dokl. Akad. Nauk SSSR, 128: 777--780. Plaksin, I.N. and Shafeev, R.S., 1960. On properties of distribution of xanthate on the surface of sulphide minerals. J. Mines Met. Fuels, 8: 1--8. Pomianowski, A., 1963. Physical chemistry of model flotation of mercury. II. Zesz. Nauk. Uniw. Jagiellon., Pr. Chem., 8: 87--113. Pomianowski, A., 1967. Electrical and surface characteristics in mercury--xanthate--air system. Thesis, Institute of Physical Chemistry, Polish Academy of Science, Krakow. Pomianowski ,A. and Czubak-Pawlikowska, J., 1973. Unpublished data. Pomianowski, A. and Kowal, A., 1973. Electrochemistry of the flotation systems. I. Chronopotentiometric characteristics of mercury, platinum and copper electrodes in an aqueous solutions of potassium ethylxanthate. Zesz. Nauk. Uniw. Jagiellon., Pr. Chem., 18: 2 4 9 - 2 5 9 . Pomianowski, A. and Leja, J., 1963. Spectrophotometric study of xanthate and dixanthogen solutions. Can. J. Chem., 41: 2219--2230.
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Pomianowski, A. and Leja, J., 1964. Equimolar solutions of xanthate and akryl trimethyl ammonium bromide adsorption on copper, nickel and sphalerite powders. Trans. AIME, 229-- 307--312. Sheka, Z.A. and Kriss, E.E., 1959. Metal xanthates. Rab. Khim. Rastvorov Kompleksn. Soedin. Akad. Nauk. Ukr. SSR, 2: 135--162. Shulman, V.M. and Larionov, S.V., 1967. Xanthate oxidation and redox potential of the xanthate--dixanthogen couple. Izv. Sibir. Otd. Akad. Nauk SSSR, SER. Khim. Nauk, 12 35--43. Stepanov. B.A., 1960. Tr. Uralmekhanobra, 7: 33. (after Shuiman, 1967.) Stepanov, B.A. and Zubkov, A.A., 1968. Effect of oxygen on floatability of the metallic mercury. Tsvetn. Met., 8: 30--31. Stepanov, B.A., Kakovskii, I.A. and Serebryakova, N.V., 1959. Nauchn. Dokl. Vyssh. Shk., Khim. Khim. Tekhnol., 2: 277--279. Taggart, A.F. and Gaudin, A.M., 1923. Surface tension and adsorption phenomena in flotation. Trans. AIME, 68: 479--530. Tipman, N.R. and Leja, J., 1975. Reactivity of xanthate and dixanthogen in aqueous solutions of different pH. Colloid. Polym. Sci., 253: 4--10. Tolun, R. and Kitchener, J.A., 1963-64. Electrochemical study of the galena--xanthate-oxygen flotation system. Trans. Inst. Min. Metall., 73: 313--322. Toperi, D. and Tolun, R., 1969. Electrochemical study and thermodynamic equilibria of the galena--oxygen--xanthates flotation system. Trans. Inst. Min. Metall., 78: 191--197. Woods, R., 1971. The oxidation of ethyl xanthate on platinum, gold, copper and galena electrodes. Relation to the meehanizm of mineral flotation. J. Phys. Chem., 75: 354-362. De Zoubov, N., Vanleugenhaghe, C. and Pourbaix, M., 1966. Copper. In: M. Pourbaix (Editor), Atlas of Electrochemical Equilibria in Aqueous Solutions. Pergamon, Cebelcor. pp. 384--392.