High temperature potentiallph diagrams for the chlorine-water system BARBARA KOLODZIEJ' AND FATHI HABASHI. Received March 26.

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1 High temperature potentiallph diagrams for the chlorine-water system BARBARA KOLODZIEJ' AND FATHI HABASHI Department of Mining and Metallurgy, Lava1 University, Quebec City, P.Q., Canada GIK 7P4 Received March ' BARBARA KOLODZIEJ and FATHI HABASHI. Can. J. Chem. 63, 935 (1985). Thermodynamic evidence supports the view that the reaction 2 C1- + 2H+ + f 02+ Cl,(aq) + H20 takes place above 127 C at ph = 0. An increase in HCI concentration and/or oxygen partial pressure allow the reaction to proceed at lower temperature. BARBARA KOLODZIEJ et FATHI HABASHI. Can. J. Chem. 63, 935 (1985). Des donntes thermodynamiques suggtrent que la rtaction 2CI- + 2H+ + toz + Clz(aq) + HzO peut se produire a une temptrature suptrieure i 127 C et i un ph de 0. Si la concentration de HCI et/ou la pression partielle d'oxygkne est augmentte, la rtaction peut se produire 6 des temptratures plus basses. [Traduit par le journal] During the aqueous oxidation of some sulfide minerals at high temperature and pressure, it was observed that the rate of reaction in HC1 solution is higher than that in H,SO, solution of the same normality, all other factors being the same (1-4). This was explained to be due to the formation of chlorine according to the reaction 2C1- + 2H' + $0' + CI2(aq) + HzO Aqueous chlorine is known to be a strong oxidant for sulfides; that is why its formation would enhance the dissolution. 'The present report is a thermodynamic analysis of the above reaction using E-pH diagrams prepared especially for elevated temperatures. Method of calculation PotentiallpH diagrams for the chlorine-water system can be derived for temperatures up to 1 50 C by a method based largely on the Criss and Cobble Entropy Correspondence Principle (5). For representing equilibrium conditions in aqueous solutions we utilized the work by Pourbaix (6) which contains diagrams at 25 C. For easy comparison of the diagrams presented in the paper with those provided by Pourbaix we use his system of numbering of reactions. The diagrams presented in this paper were constructed using the Correspondence Principle together with various other data from sources shown in Table 1. The reactions considered in the chlorine-water systems are shown in Table 2. All reactions were considered in the general form [b] TABLE I. Thermodynamic data at 25'C Formula "Dissolved chlorine. "From Pourbaix (6). ' From Latimer (7). s:<)~ Kc (cal/mol deg) TABLE 2. Chlorine-water system Reaction involving the stability of water Oz(g) + 4H' + 4e- = 2Hz0(1) Two dissolved substances [3] 2CI- = Clz(aq) + 2e- [9] CI- + 4Hz0 = CIO, + 8H+ + 8e- [I51 Cl,(aq) + 8Hz0 = 2C10, + 16~' + 14e Two gaseous substances [25] 2HCl(g) = C12(g) + 2H+ + 2e One gaseous substance and one dissolved substance [3 I ] HCl(g) = CI- + H+ so that the reduction potential at temperature T becomes [35] 2~1--= Clz(g) + 2e- [41] Clz(g) + 8Hz0 = 2C10, + 16H+ + 14e-, RT aia.l+ [2] ET=ET+-In,, zf a~n~20 It the activity of water is assumed to be equal to unity, the term -log ah+ defined as ph, and with and 150 C was calculated from the equation with TI = 298 K, -0 7, [5] AGO,? = - (T, - T,)ASO,~ + (T2 - TI)ACp]T1 it follows that using the G:, and S,,, o values shown in Table 1 and heat capacity dattvshown Tables 3, 4. The ion heat-capacity functions have been presented by Criss and Cobble in the following., form A change in free energy for each reaction at temperatures 90 ' Present address: Institute of Inorganic Chemistry, Polytechnic University, WrocXaw, Poland. -0 7, [61 CpI2irx = a7? + bt2 s&x (absolute) where a7, and b, are temperature dependent and s:,, (absolute) ' Revision received October 16, is the entropy based on SO,+ = -5.0.

2 936 CAN. 1. CHEM. VOL. 63. I985 FIG. 1. Potential-pH equilibrium diagram for the system chlorine-water at 298 K, for solution containing 1 g-atom Cl/L. TABLE 3. Values of partial molar heat capacities for some ions and non-ionic species - n Formula C,Z9x C,O];z C,"]:; "From Criss and Cobble (5). ''From Helgeson (8). ' From Kelley (9). -0.L III L lb 1; 1; 1; 1; 1; ph FIG. 2. Potential-pH equilibrium diagram for the system chlorine-water at 363 K, for solutions containing 1 g-atom Cl/L. diagrams at various temperatures shows that an increase in temperature results in a considerable decrease in the reaction potentials denoted in Table 2 by [3], [25], [35]. Thus, with the increasing temperature the area of thermodynamical stability of gaseous chlorine in equilibrium with gaseous HC1 and C1- ions becomes enlarged and also the region of relative predominance of free chlorine Cl,(aq) contained in hydrochloric acid solution becomes enlarged. On the other hand, an increase in temperature results in an increase of the oxygen reduction potential which depends also on its partial pressure. By analyzing a possibility for reaction 2C1- = Cl?(aq) + 2e- (Table 4, reaction [3]) to occur, it was found on the grounds of the above mentioned relationships that at 25 C the potential of this reaction over the entire ph range lies above the reduction potential of O2 to H20 Only a slight potential difference for reaction 2C1- = 2Cl(g) + 2e- (Table 4, reaction [35]), and for reduction of oxygen of maximum 0.06 V for HC1 solutions at ph < 1.2 might indicate the oxidation reaction of C1- to C12(g) to occur. An increase in temperature up to 90 C favours the oxidation reaction of C1- to C12(g). At this temperature the difference in potentials between oxidation reaction of C1- ions to C12(g) and reduction reaction of oxygen to H20 in HCl solution at ph < 2.2 is max V. On the other hand, for the course of reaction Results and conclusions The equations for various reactions in the C1-H20 system presented in Table 2 were used in the construction of the potential/ph diagrams shown in Figs. 1, 2, 3, 4. 'The potential/ph equilibrium diagram for 25 C presented according to Pourbaix was supplemented with an expression related to equilibrium [311. 'The equation presenting equilibrium [31] (Table 2) allows for the partial pressure of HCl over the aqueous 1 M, 3 M hydrochloric acid solution at 298, 363, and 423 K. Partial of HCl in aqueous hydrochloric acid solutions at [g] 2~1- + $0, + 2H' = ~ l ~ + H,O ( ~ ~ ) various concentrations and temperatures (10) are provided in Table 5. The diagrams presented in Figs. 1, 2, 3, 4 provide a the temperature of 90 C is still too low. Though for ph < 1.4 suitable value of log p~,-, apart from the equilibrium line for the potential of oxidation reaction of C1- ions to C12(aq) is reaction [31]. In order to determine the effect of hydrochloric slightly higher than that for reduction of oxygen but in this ph acid concentration on the equilibria in question at 150 C also region, as shown by the diagram, one does not deal with the the potential/ph diagram for HCl was plotted (Fig. 4). HCI solution. Analysis of equilibria for the same reactions at Analysis of the equilibria presented in the potential/ph 150 C (Fig. 3) indicates that the course of oxidation reaction of

3 KOLODZIU AND HABASHI TABLE 4. Standard free energy AG" and equations relating ET, ph, and activity AGO, T Equation (K) (cal/mol) E,/pH relationships Volts/pH/activity [bl ET= pH log Po, = pH log Po? = pH log Po2 [31 [CIZl;lq ET = log -7 [CI-1- [CIZlaq = log 7 [cl-1- [CIZIaq = log - [CI-1' [ClO,l = pH log - Lc1-1 [ClO,l = pH log - LC1-1 [CIO,]' E,- = pH log - [Cl21., = pH log - [CIZI,, = pH log - [Clllaq PCI~[H+I' = pH log - PHCI pc,?[h+12 = pH log --- PHCI [ log [Cl-] = ph + log p~cl log [Cl-] = ph + log ph,, log [CI-] = ph + log p~cl PCl? = log - [cl-l2 PC]. = log - [Cl-l2 [c10,] = pH log = pH log - PC12 Limits of the domains of relative predominance of the dissolved substances [3'1 CI-/C12(aq) ET = log C* = log C 423 = log C

4 CAN. J. CHEM. VOL. 63, 1985 TABLE 4. (Cotzcluded) T AG: Equation (K) (cal/mol) ET/pH relationships Volts/pH/activity Er= pH log C = pH log C = pH log C Limits of the domains of relative predominance of the gaseous substances [25"]t 298 ET = pH log [Cl-] HCl(g)/ 363 = pH log [Cl-] C12(g),CI- 423 = pH log [Cl-] Solubility in g-atoms of chlorine per litre [35'I 298 E, = log C log pcll CI-/CI2(g) 363 = log C log 423 = I.I log C log p,,, *Concentration C calculated in moles of dissolved substance per litre t The relation ET is explained in the Appendix. FIG. 3. Potential-pH equilibrium diagram for the system FIG. 4. Potential-pH equilibrium diagram for the system chlorine-water at 423 K, for solutions containing I g-atom CI/L. chlorine-water at 423 K, for solutions containing 3 g-atom CI/L. C1- ions to C12(aq) is thermodynamically justified. If the partial pressure of oxygen is 20 atm, the maximum difference of potentials for oxidation of Cl- and reduction of oxygen in HCI solution at ph < 1 is 0.1 V for the reaction whose product is C12(g) and 0.37 V for the reaction which yields Cl,(aq), respectively. An increase in the hydrochloric acid concentration up to 3 M causes that under the same conditions the maximum potential difference increases up to 0.14 V and 0.43 V, respectively (Fig. 4). In order to estimate the temperature at which the course of oxidation reaction of C1- ions to Cl,(aq) becomes thermodynamically justified the following procedure was applied. The ET potential of oxygen reduction was determined at 25, 90, and 150 C, assuming Po, = 20 atm and ph of solution equals 0. It amounts to 1.25,- 1.24, 1.24 V, respectively. Because of such insignificant changes it was as- TABLE 5. Partial pressures (in atm) of HCI in aqueous hydrochloric acid solutions at various concentrations and temperatures Temperature (K) HCI "xtrapolated -- values (on the grounds of log ph,, = f(t)). sumed that its value does not depend on temperature and is 1.24 V. Then the relation of ET versus temperature for oxidation reaction of Cl- ions to C12(aq) was determined. On this basis it was found that the ET potential of this reaction assumes the

5 KOLODZIU AND HABASHI 939 value of 1.24 V at about 127 C. Thus one should infer that above this temperature the course of oxidation of the C1- ions to Cl,(aq) under oxygen pressure in HCI solution at ph = 0 is thermodynamically justified only at temperatures exceeding 127 C. Similar considerations carried out for solutions with varying ph and, hence, at various HC1 concentrations indicate that an increase in HCI concentration allows this reaction to proceed at correspondingly lower temperatures. An increase in partial pressure of oxygen has a similar effect on the course of oxidation reaction of C1- ions to Clz(aq): the higher the partial pressure of oxygen, the lower the temperature at which this reaction can take place. Recent experimental work by Kiwi and Gratzel (11) has shown that chloride ion in acid medium can be oxidized in presence of solid catalysts. Thus, at ph = 1, and lo-" Me4+, a solution of M NaCl generates Cl? at the rate of 60 km/l h in presence of ruthenium oxide, RuOz.xHzO. In view of the thermodynamic calculations presented here, the formation of C1, during the aqueous oxidation of sulfide minerals in HCI medium seems plausible. Acknowledgements The authors acknowledge with thanks the financial aid of the Institute of Inorganic Chemistry, Polytechnic University in Wrod'aw, Poland, and the Canadian Natural Sciences and En-, gineering Research Council. I. F. HABASHI and T. TOOR. Met. Trans. 10B, 49 (1979). 2. T. MlzOGUCHl and F. HABASHI. Trans. Inst. Min. Met. 92C, 14 (1983). 3. T. MlzOGUCHI and F. HABASHI. Intern. J. Min. Process. 8, 177 (1981). 4. K. NAITO and F. HABASHI. Trans. Inst. Min. Met. C93, 69 (1984). 5. C. M. CRISS and J. W. COBBLE. J. Am. Chem. Soc. 86, 5385 (1964). 6. M. POURBAIX. Atlas of electrochemical equilibria in aqueous solution. Pergamon, Oxford W. M. LATIMER. The oxidation states of the elements and their potentials in aqueous solution. Prentice-Hall, New York C. J. HELGESON. J. Phys. Chem. 71, 3121 (1967). 9. K. K. KELLEY. Contributions to the data on theoretical metallurgy. Bulletin 584, US Bureau of Mines, Washington, D.C., Perry's Chemical Engineer Handbook. 4th ed. McGraw-Hill, New York pp and J. KIWI and M. GRATZEL. Chem. Phys. Lett. 78 (2). 241 (1981). Appendix The equilibrium between the two gaseous substances HCl(g) and Cl,(g) should be referred to in the chlorine-water system under consideration which contains also the C1- ions derived from dissociation of aqueous hydrochloric acid solution. Hence, in order to determine the limits of the domain of relative predominance of HCl(g) and Cl?(g) in the presence of the C1- ions one should consider two reactions denoted by [25] and [31] in the Table 4. The equilibrium of reaction [25] is described by the Nernst equation: LI' p;,c, while that of reaction [31] by the equation log [Cl-] = log KT + ph + log p,,cl where KT = [CI-11 [H+]/pwcl The following combination of the above two relationships: RT log [Cl-] 2F yields an equation used for determination of the limits of the domains of relative predominance of the HCl(g) and Cl?(g) in the presence of the C1- ions. This equation is denoted by [25"] in Table 4 (according to the Pourbaix nomenclature) and is given for the respective temperatures 298, 363, and 423 K.

AP Questions: Thermodynamics

AP Questions: Thermodynamics 1970 Consider the first ionization of sulfurous acid: H2SO3(aq) H + (aq) + HSO3 - (aq) Certain related thermodynamic data are provided below: H2SO3(aq) H + (aq) HSO3 - (aq)