Formation constants of some chloride complexes of mercury(II) in methanol-water mixtures

.I illorA', tu4~l ( h~,m., t973, Voh 36, pp. 1837 1839. l'ergamon Press. Printed in Great Britain

FORMATION CONSTANTS OF SOME CHLORIDE COMPLEXES OF MERCURY(II) IN M E T H A N O L - W A T E R MIXTURES K. P. A N D E R S O N .

A. L. C U M M I N G S . . I . L .

BILLS and K. J. W A L K E R . Jr.

D e p a r t m e n t o f C h e m i s t r y , Brigham Y o u n g University. Prow~. Utah 84602

(R~'ceivcd 15 October 1972) Abstract The solubility of m e r c u r y ( l l ) chloride in a q u e o u s m i x t u r e s of 0 . 8 and 16 per cent m e t h a n o l and various c o n c e n t r a t i o n s of hydrochloric acid was studied. Values of c o n s t a n t s for the f o r m a t i o n o1" the m e r c u r y - c o n t a i n i n g species were determined. A g r e e m e n t with literature values ['or aqueous solutions is good and indicated the prolitabilily of further studies using this a p p r o a c h in mixed solvent i n v e s u g a u o n s .

INTRODUCTION

EXPERIMENTAi~

THIz INCRKASED solubility of lnercury{ll) chloride in aqueous solutions containing excess chloride ion has been known for many years[1 3]. Its solubility has also been investigated in pure methanol[4]. Formationconstant data have been collected for aqueous solutions (see Table 21, but little work has been done in methanolwater mixtures. This study is an investigation of saturated mercury(lI) chloride dissolved in aqueous mixtures of 0. 8 and 16 per cent methanol and various quantities of hydrochloric acid at 25°C.

All chemicals \~erc reagent grade. H y d r o c h l o r i c acid solutiot>, \\ere s t a n d a r d i z e d using trislhydroxymcthyl}aminonaeHaanc. The initial mobil(lies o[ tile HCI (moles H('t kg H_,O -> ( ' H 3 ( ) H ) tire li~ted in Table I. The ~,olution> \'.crc equilibrated x;ith excess H g ( ' l f l s t [or a nainhnum of >even days in a x~ater bath controlled tit 25 + 0.01°( ". Each ~,ample \\a,, c o n t a i n e d m a settled vial lixed to at rotatillg drulu m o u n t e d in the ,aatcr bath. F o l l o w i n g equilibration, the samples \\ere a[Io\\cd to settle in the bath for several ]1o1.11>, before analysis \\as begun. The chloride-ion act(\it\ \\as nlca~,Ul'ed \\ ith tin Orion 98-17 c o l n b i n a t i o n chloride specilic-

Table I. Molal concclltrations and act(\it> coct}icient,, 8.(/6 '3( M e t h a n o l

(! '),'i M e l h u n o l

It ('1~,,~, 0.0 0.00849 0.0170 0.0255 0.034() 0.O425 I).0510 ().0595 I) 11681 11.()7(~7 0-087,2 ~).102 (b.NI5 (~.31)9 ((.414 0.519 ().625 0.73! (1.84(I 0.949 I.I!57

16-5(1'){I McthanoI

)qu,

('l

ttg,,~

tt(l~,,,

Ym~

('1

It g,,,,

H('li,,~,

lq,<1

1.000 0.911 0.883 0.864 0.850 (1.838 (I-829 ~.X20 uSl4 (I.8(18 O.SIB I).794 0-764 0.754 0.754 0.758 ().766 1!.776 0.7,'<9 0.802 (1.815

0.000153 0.00174 0.00368 0.00569 O-0O8] 7 0.00966 (I-0124 0.0149 /).0IS7 0.0186 0.0231 O.0238 04)518 0.0S00 0.109 0.112 (!.145 0.179 O. I'),~ 1).196 (I.233

0.263 0.264 0.266 II.279 0.29O 0-!98 0.326 0.324 0.331 (t.352 O.352 0.352 ().505 C).633 (b761 0-814 11.85(I 0.973 1.018 1.162 1.239

i).o 0.00814 0.0163 0.0244 0.0326 (I.0408 04)571 0.0653 1/4)898 (I.104 0.106 0.2(/5 0.2(17 0.316 0.423 0.530 (I.639 0.748 (1.857 0.970 14182

1.000 (I.908 0.878 0.859 0.844 O.832 [).813 0.807 0.789 (i.7,";1 i1.780 o.751t 0.751 ().741 0.741 0.745 0.753 0.761 0,770 0.779 (I-788

0.000127 0.00175 0.0(t320 (1.00597 0.O0673 0-00782 0.00906 0.0119 (I.0160 0.0186 ~1.0162 ().036t~ 0.0419 (b0663 0.0877 0.110 0.129 (i.154 O. ]75 0.20(1 0.223

(1.352 (I.362 0.365 0.375 0-383 0.37i 0.396 0.423 (!.446 0.480 ~)431 O-525 O-542 0.626 0-740 0.873 0.914 !).988 1.029 1.192 1.218

0.0 (1.011833 0.01(~7 ()-(1250 0.0417 0.0500 0.0919 0. I00 0./09 0.215 0.323 0.433 (t.542 0.654 0.766 (I-878 0.993 1.11)7

1.000 (!.902 0.868 0-840 (i.818 0-808 (/.773 0.770 0.764 0.733 0.720 0.718 0.721 0.727 0.737 0-748 (I.760 ()771

I837

(1 ()I)O0101 04)(1124 0.00259 0.00~44 O-00521 ().00724 0.133 0.0106 0.01(~5 11.0345 0.0517 !l I}741 0.0930 0.110 0.127 I).145 (!.162 0.177

It g~,,, 0.286

0.293 (I-295 (1293 (t 3()(1 1)-307 (1.]37 0-336 ~)368 0.54l 0.646 0.734 (Ix)()4 1-065 1.324 145 v 1.597 1:,36

1838

K.P. ANDERSON, A. L. CU?,~,IINOS, J. L. BILLSand K. J. WALKER, Jr.

ion electrode and an Orion 801 digital vohmeter. Standard chloride solutions were used to prepare calibration curves for the three solvent systems, plotting log (~:hloride activityt versus millivolt reading. The linear portion of the calibration curve was extrapolated for use on solutions of low ( < l0 *) chloride activity. Values of the mean ionic activity coefficient for HC1 in the various methanol water mixtures were determined by interpolation of literature valuesf5-7] measured in 0, I0 and 20 per cent methanol. These values are listed in Table l, along with the final molalities of free chloride ion. Aliquots of the saturated solutions were diluted. and the total mercury concentration was determined with a Perkin-Elmer 303 atomic absorption spectrophotometer. The final molalities of total mercury in the saturated solutions are also listed in Table 1. Formation constants for the various mercury-containing species in each solvent mixture were obtained by a relativedeviation least-squares fittingf8, 9] of the data in Table 1 to the mass-balance equation for total mercury. The species required for a satisfactory fit were HgCI +, HgCI 2, HgCIf. and HgCI~.. {The HgCI 2 was required only in 16 per cent CH~OH.) The mass-balance equation is fHg.,,~,,j = [HgCl+~ + [HgCI2] + [HgCI3~ + [ngCI 2 ](1) where the square brackets denote molalities, Equation (1) can be rewritten in terms of formation constants. [Hg'°'"'] = K';?-{(21:] + g'z +

K'2KSCI-]

]2 y2...........

+ K~2K3K~[CI

K,, iS the equilibrium constant for the formation of HgCI~-" from HgCt~5'~' and CI-. K~2 is the constant for HgC12(s) going to HgCtz(dissolved). In the derivation of Eqn (2), the activity coefficient for HCI Ip from Table 1) was assumed to apply also to HgCV and HgCI;. The activity coefficient of HgCI 2 was taken as unity, and that of HgCt~ was assumed to be ?,a. The addition of HCI to most of the solutions decreased the concentrations of HgC1 + to such an extent that K 2 could not be determined very accurately from Eqn (2) alone. In the solutions with no initial HCL where the [HgCI*] was highest, hydrolysis of the HgCI z was significant. The mercury-containing hydrolysis products proposed by Luther [10] and by Ciavatta and Grimaldi[I 11 are all uncharged species. Therefore, the charge-balance equation is !H+2 + UHgCI+] = [Cl-] + UHgCI~-]

{3)

where HgC1] was omitted as negligible in the absence of additional HC1. The same pH of 3.5 was measured for saturated HgCI 2 in 0, 8 and 16 per cent CH3OH, giving [H + ] as 3.2 x I0 -'~. The UCt-] was taken from Table 1, and the [HgCI~-] was calculated as the next-to-last term in Eqn {2). Equation 13) was solved for [HgCI+]. and an improved value of K 2 was calculated from the relationship: K 2 = K~e/;2[HgCI+][CI ]. This new value of K z was used in a second fit to Eqn (2~. and the calculation was repeated until the values of all constants did not change with succeeding iterations. Since the total mercury present in hydrolytically produced species was never greater than 3 x 10-'* molal, its exclusion from the mass-balance equation did not affect the calculated values of the constants within experimental error. Equation 121 was found to reproduce the values of Table 1 for Hgtou~ with and average deviation of 3.1, 2.9 and 3,2 per cent in 0. 8.06 and 16.50 per cent methanol, respectively.

RESULTS AND D I S C U S S I O N

Logarithms of the formation constants determined in this work, as well as pertinent values from the literature, are listed in Table 2. The uncertainties to be attached to our values are +0.2 for log K 2 and +0.05 for the other log values. The literature values are not strictly comparable to ours because most of the previous work was done in solutions containing NaC104. The molar concentrations of NaC104 (or NaC104 + HC104) were 1 and 3 in [12] and [13], respectively, and 0.5 in [14], [151, and [17]. In those reports, equilibrium constants were calculated from concentrations instead of activities. C o m p a r i s o n s within the present study are also approximate because each solvent mixture requires a separate standard state. However, the increase in K~2 on going from 0 to 8 ~ C H 3 O H is consistent with the general disruption of the ordering in pure w a t e r [ D ] by the addition of a small a m o u n t of an organic component. Curiously, a further increase in the C H 3 O H concentration from 8 to 16-5~o reverses the change in Ks2, decreasing K~2 almost to its value in pure water. Studies of the solubilities of silver salts in methanol water and other aqueous-organic mixtures indicate that general trends in changes in formation constants with solvent dielectric constant do not become observable until the initial violent disruptive effect is over-

Table 2. Logarithms of mercury(II)--chloride formation constants Percentage of CHaOH 0

8.06 16.50

log K,2

log K 1

-0.59

-0-57[2, 18] - 0.44 -0.56

log K 2

6.72112] 7,07[13]

6.53 6,5t [12] 6-91113]

6.621141] 6.80116] 6-741t7]

6.361141] 6.60[16] 6,48117] 6.93 6.90

tog K 3

log K4

1.22 1.00112] 0.97112] 0.75113] 1-38!13] 1.08113] 1.09[133 0-95115] 1.05115] 0 . 5 7 1 1 6 ] 1.46!16] 0 . 8 5 1 1 7 ] 1.00~17] 1-8712, 18] 1.06 1,30 0.22

Mercury(ll) chloride in methanol ~ater come as more and more organic component is added to the mixture. The present method did not yield a value for log K ~, but the literature values were averaged and the result {6.79) was added to log K 2 to give a value of 13.32 for log f12, the formation constant of HgCI 2 from Hg 2+ and 2C1- in 0°/o CH3OH. This log/~z was then subtracted from log K~2 to give a value of - 13.91 for log K o, the solubility-product constant in water. These da~ia indicate that this approach to the study of mercury(ll) chloride complexes in various solvent mixtures is valid and thal similar studies can profitably be performed in an extended range of solvent mixture composition. REFERENCES

1..I. Davy, Phi/. Trans. 112, 359 (1822). 2. H. C. Thomas, J. Am. chem. Soc. 61,920 (1939). 3. H. Stephen and T. Stephen, In Binary S.vslems, Vol, 1. Part 1. Pergamon Press. Oxford (19631. 4. M. Z. Dukelski, Z. anor~] allg. Chem. 53, 327 (19071. 5. M. Randall and L, Young, J. Am. chem. Soc. 50, 995 (19281.

1839

6. H. S. Harned and H. C. Thomas. J. Am. chem. So(. 58, 763 (19361. 7. M. Paabo, R. Bates and R. Robinson, Anah't. Chem. 37, 463 (19651. 8. K. P. Anderson and R. L. Snow, J. chem. Educ. 44, 756 (19671. 9. T. P. Kohman, J. chem. Educ. 47, 65'7 (19701. 10. R. Luther, Z. phys. Chem. 47, 107 (-1904). 11. L. Ciavatta and M. Grimaldi. J. inor~, nucl. Chem. 30. 563 (19681. 12. L. Ciavatta and M. Grimaldi, .I. inor~,,, nm/, Chem. 30, 197 (19681. 13. R. Arnek, Ark. Kemi24, 531 (1965). 14. L. D. Hansen, R. M. lzatt and J. J. Christensen, lnorg. ('hem. 2, 1243 (19631. 15. Y. Marcus, Aeta chem. seand. 11,599 (19571. 16. K. Damm and A. Weiss, Z. Natuti/i 10b, 534 (19551. 17. L. G. Sillcn. Acta chem. stand. 3, 539 (1949). 18. A. B. Garrett. J. Am. chem. Soc. 61. 2744(19391. 19. F. Franks, Phv.~ico-Chemica/ Fro~e~w.~' in Mi\cd Soli'cnts IEdited by F. Franks), pp. 50 75. American Elsevier. New York, 1967. 20. K. P. Anderson, E, A. Butler and E. M. Woolley, .I. phy.~. Chem. 71, 3566 (1967). 21. K. P. Anderson, E, A. Butler and E. M. Woolley, J. ptly~. Chem. 75, 93 ( 19"71).