Mercury equilibrium constants

Use the information in Appendix 2B to determine the equilibrium constant for the disproportionation of mercury(I)... 

Mercury-chloride complexes in dilute solutions. This slightly more difficult example will be useful in showing how to handle poorly conditioned systems of equations. It is assumed that mercury chloride HgCl2 is dissolved in pure water with a molality m = 10 5 mol kg-1. Given the equilibrium constants for chloride complex formation... 

From this voltage, calculate the equilibrium constant for the reaction Hg2+ + 41 - lignin 0.5 M KI, virtually all the mercury is present as Hgl. ... 

We write the usual equilibrium constant expressions but omit the pure solid carbon in part (a) and the pure liquid mercury in part (b) because their concentrations are constants that are incorporated into the equilibrium constant Kc. 

To apply equation (3) for calculation of the equilibrium constant K waves Ia and ic must both be limited by diffusion. To prove this the current is measured under conditions when it is 15% or less of the total limiting current and its dependence on the mercury pressure is followed. A diffusion current must, under these conditions, show a linear dependence on the square root of the height of the mercury column. Whenever possible, polarographic dissociation curves should be compared with data on dissociation obtained by other methods, e.g. potentiometry, N.M.R. or spectrophotometry. In the latter case it is important to show that the species responsible for a given polarographic wave is identical with that responsible for the observed absorption peak. 

Literally hundreds of complex equilibria like this can be combined to model what happens to metals in aqueous systems. Numerous speciation models exist for this application that include all of the necessary equilibrium constants. Several of these models include surface complexation reactions that take place at the particle-water interface. Unlike the partitioning of hydrophobic organic contaminants into organic carbon, metals actually form ionic and covalent bonds with surface ligands such as sulfhydryl groups on metal sulfides and oxide groups on the hydrous oxides of manganese and iron. Metals also can be biotransformed to more toxic species (e.g., conversion of elemental mercury to methyl-mercury by anaerobic bacteria), less toxic species (oxidation of tributyl tin to elemental tin), or temporarily immobilized (e.g., via microbial reduction of sulfate to sulfide, which then precipitates as an insoluble metal sulfide mineral). 

A similar pH dependence is operative when the metal is covered by a hydroxide layer, or an oxide hydrate layer, provided that the appropriate equilibrium constants are considered. Usually, these electrodes need calibration (the mercury-mercury oxide electrode is an exception and it is a common reference electrode in strongly alkaline systems). Metal-metal oxide electrodes are applicable in a pH range that is limited to low pH values due to the oxide (or hydroxide) dissolution. 

The value of the equilibrium constant for the reaction indicates that the favored direction of the reaction is actually from right to left. It is driven to the right by the strong complexation of Hg + with any of a large number of hgands. The third common form of mercury is as organic mercurials such as methyl mercury, CH3Hg+. 

The equilibrium constant of reaction (1), K = [Cu ][Cu ]/[Cu ], is of the order of 10 thus, only vanishingly small concentrations of aquo-copper(I) species can exist at equilibrium. However, in the absence of catalysts for the disproportionation—such as glass surfaces, mercury, red copper(I) oxide (7), or alkali (311)—equilibrium is only slowly attained. Metastable solutions of aquocopper(I) complexes may be generated by reducing copper(II) salts with europium(II) (113), chromium(II), vanadium(II) (113, 274), or tin(II) chloride in acid solution (264). The employment of chromium(II) as reducing agent is best (113), since in most other cases further reduction to copper metal is competitive with the initial reduction (274). 

Table 2. Equilibrium Constants for the Oxymercuration of Alkenes with Mercury(II) T rifluoroacetate...

Alkenes react rapidly and reversibly with mercury(II) trifluoroacetate, which is readily soluble in nonpolar solvents. The equilibrium constants, K = [adduct]/ [Hg(TFA)2][alkene], for the reaction of various alkenes with mercuryfll) trifluoroacetate in THF, and for cyclohexene in a variety of solvents, are given in Table 2. Norbornene will displace any of the other alkenes from the adducts formed with Hg(TFA)2. 

These equilibrium constant expressions indicate that equilibrium exists at a given temperature for one and only one concentration and one partial pressure of oxygen in contact with liquid mercury and solid mercury(ll) oxide. 

The three equilibrium constants are obtained directly from the experimental data and may be combined with other constants to furnish the stability constants of the mercury(ll)-selenide complexes. 

The increase in the solubility of HgSe(s) with increasing hydrogen selenide and selenide activities is well documented at high pH and the third equilibrium constant is accepted by the review. The first constant, which equals an intrinsic solubility of HgSe(s) is most likely too large due to the rather simple method used for the separation of solid material from the aqueous phase. It is tentatively included due to its potential importance in connection with the mobility of mercury from repositories. Further investigations with a more efficient phase separation technique would be welcome here. 

The formation of Hg(SeCN)4 is well established by the potentiometric work of Toropova [56TOR], while her experimental data pertaining to the formation and the formation constant of Hg(SeCN)3 only comprise a few points. In their polarographic work Murayama and Takayanagi [72MUR/TAK] studied the anodic mercury wave in the presence of 0.001 to 0.003 M SeCN . The electrode process was assumed to comprise the charge transfer Hg(l) Hg + 2e combined with the formation of Hg(SeCN)2(aq) and Hg(SeCN)3. No primary data are provided and the evaluation procedure is rather involved, which makes the assessment difficult. The results are mixed equilibrium constants, since an activity coefficient correction was applied to the Hg ion. The following complexes are thus proposed to prevail in the Hg -SeCN system ... 

A non-dependence of the thermodynamic equilibrium constant on the solvent for two different types of diols was found 34>, which indicated that Ag+ as well as undissociated AgN03 formed complexes with olefins, comparable with mercury salt-olefin complexes 35>. Further formation constant investigations 36> by gas chromatography of silver complexes of cyclo-olefins had shown that methyl substitution at the double bond markedly reduced the stability and... 

The same is true for pure liquids (undissolved) in equilibrium, such as mercury. The standard state of water is taken as unity in dilute aqueous solutions, and water does not appear in equilibrium constant expressions. 

TABLE 6.13 Equilibrium Constants for Exchange Reactions of Mercury Complexes ... 

Mercury(I) oxide decomposes into elemental mercury and elemental oxygen 2 Hg20(s) v 4 Hg(/) + 02(g). (a) Write the equilibrium-constant expression for this reaction in terms of partial pressures, (b) Suppose you run this reaction in a solvent that dissolves elemental mercury and elemental oxygen. Rewrite the equilibrium-constant expression in terms of molarities for the reaction, using (sohr) to indicate solvation. 

Lipoi( acid is rediunblo at the dropping mercury electrode (Reed et al., 1953a Ke, 1957). The half-wave potential at pH 7.0 is —0.567 volt versus the saturated calomel electrode (Ke, 1957). This value corresponds to a reduction potential of —0.325 volt with respect to the standard hydrogen electrode. The reduction potential of the dihydrolipoic acid-lipoic acid system has been calculated from the equilibrium constant of the dihydro-... 

The decrease in free energy of the system in a spontaneous redox reaction is equal to the electrical work done by the system on the surroundings, or AG = nFE. The equilibrium constant for a redox reaction can be found from the standard electromotive force of a cell. 10. The Nernst equation gives the relationship between the cell emf and the concentrations of the reactants and products under non-standard-state conditions. Batteries, which consist of one or more galvanic cells, are used widely as self-contained power sources. Some of the better-known batteries are the dry cell, such as the Leclanche cell, the mercury battery, and the lead storage battery used in automobiles. Fuel cells produce electrical energy from a continuous supply of reactants. 

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