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Mercury(lI) Oxide 307
COHPONENTS:
(1) dercury(II) oxide; HgO; [21908-53-2]
(2) Water, H 2
0; [7732-18-5]
EVALUATOR:
T. P. Dirkse Department of Chemistry Calvin College Grand Rapids, l1ichigan 49506, U.S.A. March 1983
CRITICAL EVALUATION:
Mercury is found in two oxidation states in its compounds: + and 2+. However, mercury does not show this dual behavior in its stable oxides. In the early literature Hg20 was considered to be one of the oxides of mercury and several studies were carried out to determine quantitatively some of its physical properties.
Hugarszky (1) made e.m.f. measurements of cells containing Hg20/Hg electrodes. Later, Brodsky (2) measured the temperature dependence of the activity coefficients of Ilg2(N03)2 in water, and the EO value for the Hg2+/Hg electrode and from these results and those of Bugarszky calculated the solubility product of Hg20 in water to be 1.8 x 10- 24 mo1 3 dm- 9 at 298 K.
Al1mand (3) attempted to repeat some of this work but found that the Hg20 electrodes he prepared gave the same values as did HgO electrodes. Identical values were obtained whether or not light was excluded. Therefore, Allmand concluded that all calculations based on the e.m.f. of Hg20 electrodes have dubious value because Hg20 is unstable in aqueous solution and these electrodes then are always contaminated with an unknown quantit of HgO. Later (4) it was shown by X-ray analysis that Hg20 is an intimate mixture of HgO and metallic mercury. Sanemasa (5) made a kinetic study of this disproportionation of Hg in water.
No further data on Hg20 are included because of this instability.
HgO crystals are red or yellow, depending on the method of preparation. When Hg(N03)2 is heated, red crystals of HgO are obtained, while when OH- ions are added to aqueous solutions of Hg2+ ions in the cold, yellow crystals of HgO precipitate. The difference between these two forms has been the subject of some disagreement. Cohen (6), as a result of some very careful work, obtained a value of 0.685 mV for the e.m.f. of the following cell:
Hg I HgO(yellow) I alkali I HgO(red) I Hg. (1)
This value remained constant for 3 days but only after a week had elapsed. From this Cohen concluded that the red and yellow forms were isomers and, at 298.2 K, the red form was the less stable and the less soluble. He stated that the red and yellow forms dissolved at different rates and he gave this as the reason it took about a week for the cell to reach equilibrium.
Ostwald (7) maintained that the difference between the red and yellow forms was merely one of particle size. This results in a difference in solubility, and Ostwald insisted that this was the cause of the e.m.f. of the cell represented in equation (1). As proof of this contention Ostwald ground up red crystals of HgO and obtained yellow ones. He also studied the extent to which the following reaction proceeded:
hgO(s) + H20 + 2 KBr'aq = 2 KOH'aq + HgBr2'aq. (2)
he found that the amount of KOH produced depended on the particle size of the HgO.
Hulett (8) pursued this matter further. He found that by digestion and decantation he could separate the most finely divided particles (which had a yellow color) from a sample of red HgO. Although he made only qualitative measurements he also found that by grinding red HgO he could increase its solubility in water at 298 K. He concluded from this that the red and yellow forms were not isomers but merely differed in particle size.
Schoch (9) made a microscopic examination of the two forms of HgO and concluded that there were two crystal forms: prismatic (red) and square tablets (yellow). He also found that the color did change with particle size and with heating, but
1108 V.23-k*
308 Mercury(lI)^ Oxide
COHPONENTS:
(1) Mercury(II) oxide; HgO; [21908-53-2]
(2) Water; H 2
0; [7732-18-5]
CRITICAL EVALUATION:
EVALUATOR:
T. P. Dirkse Department of Chemistry Calvin College Grand Rapids. Michigan 49506. U.S.A. March 1983
maintained that color is no criterion of crystal form because the finely ground red form changed color on heating but the finely ground yellow form did not. He concluded that the red form is probably the less soluble one at room temperature (this is not in agreement with solubility results determined later by others).
Allmand (3) measured the e.m.f. of the following cell:
Hg I HgO(s) I aq. alkali I calomel electrode. (3)
He used both red and yellow HgO. No solubility values are given in the paper but Allmand made the following observations: (a) in general. the red HgO gives the lower e.m.f. value and. consequently. is less soluble and less stable; (b) the e.m.f.'s generally decreased with time, and while this ordinarily indicates a phase change, this cannot be the explanation here because this decrease occurred with both forms of HgO; (c) the electrodes with the highest e.m.f. values had the smallest particle size when viewed microscopically; (d) the e.m.f. of a cell such as that of equation (1) was less than the experimental uncertainty of the measurement. Consequently. Allmand concluded also that the red and yellow forms were not isomers but differed only in particle size.
Ostwald (7) determined the amount of KOH produced by reaction (2). He used both forms of HgO and found that the differences were within the experimental uncertainty. He also found that by grinding the oxides more finely a larger amount of KOH was formed. Garrett and Hirschler (10) did observe differences in the solubility of the two forms but hesitated to conclude that this difference was significant enough to warrant the assumption that the two forms are isomers. The solubility of the two forms in HN03 (11) show no significant differences. Schick (12) observed that mixtures of the red RgO in water sometimes deposited a yellow form of the oxide on the walls of the container. He also found that mixtures of HgO with a small amount of water became redder in color when warmed and yellower when cooled.
Consequently. it appears that the differences that have been observed can be accounted for adequately by assuming that the yellow and red forms differ in particle size and the more finely divided the particles are (yellower) the larger the solubility value. This conclusion is confirmed by careful X-ray powder diffraction measurements. which are the same for both forms of RgO (13).
Solubility in vater.
Only five papers report a directly determined value for the solubility of HgO in water. Schick (12), Fuseya (14). Garrett and Hirschler (10), and Herz and ~iebenthal (15) made the determination at 298.2 K while Salem (16) made the determination at 308.2 K. Schick and Garrett and Hirschler used both red and yellow ~gO while the others used only the red form. In view of the relationship between solubility and particle size the best results are those in which the system was allowed to equilibrate the longest time. Schick does not state how many days were allowed for equilibration but indications are that it was a week or more. He found very little difference between the solubilities of the red and yellow forms. parrett and Hirschler found a larger difference but were reluctant to state the significance of this difference. The values they obtained with red HgO were less reproducible than those obtained with yellow HgO. probably because of non-uniformity of particle size. In view of all this the solubility of RgO in water is within 3% of 2.37 x 10- 4 mol dm- 3 at 298.2 K and within 3% of 3.47 x 10- 4 mol dm- 3 at 308.2 K.
Five papers report a value for the solubility product of Hg(OH)2' Labendzinski (17) gives a value of 1.5 x 10- 26 which was determined in his laboratory by Fulda. a colleague. However, there is no indication as to the experimental basis for this value, although it appears to be based on an e.m.f. measurement. No temperature is given. Grossmann (18) used the dissociation constants of HgBr 2 -and Rg(SCN)2-which 4 4
310 Mercury(1I)^ Oxide
COHPONENTS:
(1) i1ercury(II) oxide; IlgO; [2l908-S3-2]
(2) Watert H 2 0; [7732-l8-S]
CRITICAL EVALUATION:
EVALUATOR:
T. P. Dirkse Department of Chemistry Calvin College Grand Rapids, llichigan 49S06, U.S.A. :1arch 1983
The equilibrium constant expressions for the above reactions are:
and
(11)
(12)
and the total solubility, ~', may then be expressed as
(13)
Making substitutions from the equilibrium constant expressions and assuming the validity of the Debye-Huckel relationship for activity coefficients in these dilute solutions, equation (13), following a procedure described earlier (22), reduces to
mT = mUg(OU) + KO^ • rnou- + KO • (rnou-)2/YNaOH • aU 0 (14)
Fuseya (14) and Garrett and Uirschler (10) found a linear relationship between the solubility of HgO and molality of base up to a base molality of 1 mol kg-I. This indicates that the last term in equation (14) contributes less than the experimental uncertainty and the concentration of HgO~-in these dilute solutions is negligible. Equation (14) then, in effect, becomes
(1S)
A plot of fiT vs alkali concentration gives an intercept of 2.4 x 10- 4 mol Ug(OU) dm-3 • This is the solubility of Hg(OU)2 at 298.2 K and the value of the equilibrium quotient, K16, for reaction (16).
IIgO(s) + H20 .. Hg(OH)z(sln)
K16 .. mHg(OI1)
(16)
(17)
o The slope of the line is 3.1 x 10- S and this is the value of K9 at 298.2 K. From the values of KOg, K16' and the ion product constant of water, KS can be calculated and its value at 298.2 K is 1.4 x 10-lS mol dm- 3 •
The work of Garrett and Uirschler (10) shows that in alkalies the solubility of HgO is independent of the cation but does depend on the nature of the anion as well as on the total ionic strength. However, in the absence of alkalies, the solubility does appear to depend on the nature of the cation as well (IS). In chloride
solutions the increase in solubility due to the presence of the cation was Li+ > Na+
Kt. Hut the solubility is not a linear function of the concentration of the salts used. No work has been reported dealing more specifically with the effect of the total ionic strength in alkaline solutions.
In neutral and acid solutions Hg(OU)2 acts as a base and may ionize as follows:
Ug(OH)2 = HgO~ + OU-
Hg(OH)2 .. Ug2+ + 2 OU-
(18)
(19)
Mercury(ll) Oxide 311
COHPONENTS:
(1) Llercury(II) oxide; HgO; [21908-53-2]
(2) Water~ H 2 0; [7732-18-5]
CRITICAL EVALUATION:
or
EVALUATOR:
T. P. Dirkse Department of Chemistry Calvin College Grand Rapids, Nichigan 49506, U.S.A. llarcil 1983
HgO(s) + W ..
ItgO(s) + 2 W .. Hg2+ + H
The corresponding equilibrium quotients are:
KUi .. mhgOW • moh-/ mltg(OH)
K19 .. mltg 2 + • (moH-)2/ (^) IDHg (OH) 2 K20 .. mHgOlr+-/mW-
K21 .. mHg 2 +/(mw)
(21)
Evaluation of the equilibrium constants for reactions (18) -(21) is complicated by the fact that the solubility ot HgO in acid solutions is affected by the anions present because of the formation of complex ions involving these anions. This is particularly true of the halide ions. Another complication is that the solubility is large enough so that a significant, but unknown, amount of W- ions is used up in the dissolving process. Garrett and Howell (11) used a method of successive approximations to arrive at the following values: K20" 0.17 and K21 .. 53, both at 298.2 K. These values, together with the ion product constant of water and the solubility of undissociated Hg(OH)2 in water give K18 .. 7.1 x 10- 12 mol dm- 3 at 298.2 K and K19 .. 2.2 x 10- 23 mo1 2 dm- 6 at 298.2 K. The latter value is slightly less than that calculated from the solubility of HgO in alkali solutions. However, it is the better value because of the uncertainties in the K~O value for Hg(OH)2.
A summary of the variation of the solubility of HgO with pH is given on Figure 1
Oka (23) reports a value of 10- 22 mo1 2 dm-6^ at 298 K for K19. This was obviously calculated from data obtained by titrating a solution of Hg(N03)2 in HN with NaOH. However, no indications are given as to what the calculations were or how they were made. Nor is there any indication of the experimental uncertainty. ~onsequently, the value of K19 given above is to be preferred to that of Oka.
Bilinski, et al. (24) determined values for K18 and K19 in an indirect manner. The~ measured the solubility of HgC03·2HgO at a constant ionic strength of 0.5 mol dm- over a wide range of pH. At a pH above about 9.6 yellow HgO was the only solid phase appearing at 298 K. The d~ta are presented only in graphical form. However, the numerical data were kindly made available to us in a personal communication from Dr. H. Bilinski. The solubility of HgO at pH .. 10.4 is 4.64 x 10- 4 mol dm- 3 • It is independent of the carbonate ion concentration up to a carbonate ion concentration 0 about 0.01 mol dm- 3 • This solubility is slightly larger than that shown in Figure 1 at the corresponding pH. This is due to the presence of mercury-carbonato in addition to mercury-hydroxo solute species.
An equation for the total solubility of mercury in the form of the si~lest polynomial in lr+- ion concentration that would fit the data was developed(2). From this treatment a value of K18 .. 6.8 x 10- 12 mol dm- 3 was obtained. This is within the experimental uncertainty of the value given above.
A value of K19 .. 5 x 10- 22 mo1 2 dm-6 was also deduced from this treatment. This is about an order of magnitude larger than the value given above. However, because of the indirectness of the method used, the lower value is the preferred one.
Mercury(ll) Oxide 313
COHPONENTS:
(1) Hercury(II) oxide; HgO; [21908·'53-2]
(2) Water: H 2
0: [7732-18-5]
CRITICAL EVALUATION:
EVALUATOR:
T. P. Dirkse Department: of Chemistry Calvin College ' Grand Rapids, llichigan 49506, U. S.A. March 1983
A summary of the equilibrium constants for various reactions involving HgO and Hg(OH)2 is given in Table I.
Table I
Tentative equilibrium constant values at 298 K for HgO/Hg(OH)2 system, based on solubility measurements
Hg(OH)2(sln) .. l~ + H11g0- K5 .. 1.4 x^ 10-^15 mol dm- 3 2 0 ligO(s) + OH-^ ..^ HH (^) g 20- (^) K9 .. 3.1^ x^ 10-^5
HgO(s) + (^) H2 O ..^ Hg(OH)2(sln) (^) K16 .. 2.4 x^ 10-^ 4 mol dm- 3
Hg(OH)2(sln) .. HgOW + OH- (^) K18 .. 7.1 x 10- 12 mol dm- 3
Hg(OH)2(sln) ..^ lig^ 2+^ +^2 OH-^ K19 .. 2.2 x 10-^23 mo1^ 2 dm-
HgO(s) + (^) W ..^ HgOW K20 .. 0.
HgO(s) + 2W ..^ Hg2+^ + (^) H2 O (^) K21 .. 53^ mol-^1 dm^3
KOo for Hg(OH)2 is approximately 10- 26 mo1 3 dm- 9• s
The above values are considered as tentative because there is a lack of corroborating work reported.
Some solubility determinations have also been made in aqueous HC1 (25) and in aqueous HF (26, 27). The solubility values here tend to become erratic because of complexes formed with the acid anion. Thus, in HC1 solutions a black solid phase, 2ligO·HgC12, was formed in molalities of HCl as low as 0.007 mol/kg. Above molalities of 0.1 mol/kg it appeared to be the only solid phase preoent and its solubility in water at 298 K was measured to be 11.4 x 10- 4 mol/kg H20.
Likewise, the solubility of HgO in aqueous HF is accompanied by the formation of llgO·HF and HgF202H20. Jaeger (26) does not explicitly state the nature of the equilibrium solid phase but the implication is that it is HgO. Up to HF concen- trations of 4 mol dm- 3 the solubility of HgO in aqueous HF at 298 K (26, 27) can be expressed as
CHgO .. 0.103 CHF (26)
At liF concentrations greater than 4 mol dm- 3 the data available are erratic, perhaps due to the appearance of more than one solid phase.
A few data are available on the solubility of HgO in aqueous solutions of sodium sulfide (28) at 298 K. However, these data cannot be evaluated because of insufficient experimental detail in the article and the lack of other work dealing with the same system.
'314 Mercury(lI) Oxide
COHPONENTS:
(1) Mercury(II) oxide; HgO; [21908-53-2]
(2) Water; H 2
0; [7732-18-5]
CRITICAL EVALUATION:
EVALUATOR:
T. P. Dirkse Department of Chemistry Calvin College Grand Rapids, Michigan 49506, U.S.A. Harch 1983
IQ I
E
"'C
0
E
T/K 0 Ot :I: u^328
318
308
298
Figure 2. Solubility of HgO in aqueous KOH, ref. (35).
316 Mercury(ll)^ Oxide
COHPONENTS:
(1) Mercury(II) oxide; HgO; [21908-53-2]
(2) Water; H 2 0; [7732-18-5]
CRITICAL EVALUATION:
EVALUATOR:
T. P. Dirkse Department of Chemistry Calvin College Grand Rapids, Michigan 49506, U.S.A. March 1983
Solubility of HgO under hydrothermal conditions.
In recent years the solubility of HgO has been of interest to those studying various geochemical processes. As yet relatively few data have been reported in this area. Shlyapnikov and Shtern (33,34) measured the solubility of HgO in NaCI solutions exposed to a C02 atmosphere at varying pressures. At a pressure of 50 atm C02 the solubility of HgO in water at 293 K is reported as 0.0010 mol dm- 3 • This is larger than the solubility in water when exposed to ordinary atmospheres. It is also about twice as large as the value in carbonate solutions (24). However, no information is given as to the experimental uncertainty. Because of this and the fact that no other similar data have been reported, the solubility values given by Shlyapnikov and Shtern are tentative and must await further confirmation.
The solubility of HgO in solutions under a carbon dioxide atmosphere at an elevated pressure increases markedly as the concentration of NaCl in the water increases. This is due to the formation of mercury-chloro and basic mercury- carbonato solute species.
Mercury(lI) Oxide 317
COHPONENTS:
(1) l'1ercury(II) oxide; HgO; [21908-53-2]
(2) Water; H 2 0; [7732-18-5)
CRITICAL EVALUATION:
REFERENCES
EVALUATOR:
T. P. Dirkse Department of Chemistry Calvin College Grand :lapids, 11ichigan 49506, U.G.A. Ilardi 1983
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