Mercury(II) chloride nonoccurrence in nature: A molecular-level investigation

Abstract

Prepared by Chinese alchemists in the first millennium A.D., mercury(II) chloride (HgCl₂) has been known since the Middle Ages. However, as it has never been found in mercury mines, its nonoccurrence in nature appears to be a basic assumption. Electrochemical principles were applied to show that the secondary mineral calomel in mercury mines is partly converted to HgCl₂ in a multistep natural process. This showed how HgCl₂ does occur naturally, providing a theoretical basis for the detection and quantification of solid HgCl₂ as a mineral in mercury mines.

Keywords

Calomel complexation mercury mercury(II) chloride oxidation-reduction

Manufactured by the direct combination of the elements mercury and chlorine at > 300°C and condensation of the escaping sublimate vapor [1], mercury(II) chloride (HgCl₂) is used in disinfectants, fungicides, wood preservatives, and photography [2]. Obtained in the past by sublimation from mercury(II) sulfate (HgSO₄) and sodium chloride (NaCl), HgCl₂ has been known since the Middle Ages [3]. In the first millennium A.D., it was prepared by alchemists from China [4], which is the largest worldwide mercury emitter of today [5]. An account of its preparation was also given by Rhazes [4], whose name is widely known in Arabic alchemy [6]. According to Farrar and Williams [4], both from University of Manchester Institute of Science and Technology, “this is perhaps the first synthesis of a definite chemical compound that does not occur in nature.” Found along lines of previous volcanic activity, cinnabar (HgS) is the only important ore of mercury [3]. HgCl₂ does not have any “mineral name” [7], because mineralogists have not found it as a naturally occurring mercury compound. In other words, HgCl₂ is not one of the minerals recognized by the International Mineralogical Association (IMA) [8]. However, not finding a compound in nature is not a sufficient reason to conclude that it does not occur there. Is it true that HgCl₂ is not a naturally occurring compound? This position paper is a theoretical investigation to answer this question.

There are two types of inorganic mercury ions: mercury(I) ( ${Hg}_{2}^{2 +}$ ) and mercury(II) (Hg²⁺). Mercury(I) chloride (Hg₂Cl₂) is found as the secondary mineral calomel in mercury-bearing deposits [9]. An example of the natural occurrence of calomel is its presence in Moschellandsberg, Rhineland-Palatinate, Germany [10]. Another example is a study in which mine water samples around the abandoned Haliköy mercury mine in Turkey were found oversaturated with calomel [11]. The slight solubility of Hg₂Cl₂ in water, shown by “Equation 1” and its K_sp = 1.43×10^-18 [12], results in the low concentration of aqueous ${Hg}_{2}^{2 +}$ of [ ${Hg}_{2}^{2 +}$ ] = 7.10×10^-7 M in a saturated solution of Hg₂Cl₂ in water. ${Hg}_{2} {Cl}_{2} (s) ⇄ {Hg}_{2}^{2 +} (aq) + 2 {Cl}^{-} (aq)$ (1) $K_{sp} = [{Hg}_{2}^{2 +} (aq)] {[{Cl}^{-} (aq)]}^{2}$ (2) $1.43 \times 10^{- 18} = s \times {(2 s)}^{2}$ (3)

s = 7.10×10^-7

The Hg₂Cl₂ that occurs naturally in mercury mines is partly exposed to and dissolved in natural waters or precipitation, according to “Equation 1”; however, the question is whether any of the aqueous ${Hg}_{2}^{2 +}$ ions produced in the dissolution process are converted to Hg²⁺.

Mousavi [13] shows in detail that in an acidic environment in the presence of gaseous oxygen (O₂) in air, aqueous ${Hg}_{2}^{2 +}$ is almost completely oxidized to aqueous Hg²⁺ at 25°C (with the equilibrium constant K_Mousavi = 9.4×10²⁰). However, in the field of the chemistry of mercury, there is a similar reaction that requires neither an acidic environment nor the presence of O₂: the reaction for the disproportionation of aqueous ${Hg}_{2}^{2 +}$ ions, leading to the formation of Hg²⁺ ions, shown by “Equation 4” [3]. ${Hg}_{2}^{2 +} (aq) ⇄ {Hg}^{2 +} (aq) + Hg (l)$ (4)

Greenwood and Earnshaw [3] note that the measured cell potentials “of concentration cells of mercury(I) salts are only explicable on the assumption that a 2-electron transfer is involved.” In other words, they acknowledge that n = 2 mol e^-/mol rxn. Therefore, the two half-reactions that result in the disproportionation of ${Hg}_{2}^{2 +}$ according to “Equation 4” are shown by “Equation 5” and “Equation 6” as follows: ${Hg}_{2}^{2 +} (aq) \to 2 {Hg}^{2 +} (aq) + 2 e^{-}$ (5) ${Hg}^{2 +} (aq) + 2 e^{-} \to Hg (l)$ (6)

Greenwood and Earnshaw [3] further report the standard reduction potential (E°) values of + 0.920 V for the oxidation half-reaction and + 0.8545 V for the reduction half-reaction. Therefore, the standard cell potential of the reaction shown by “Equation 4” is $E_{cell}^{\circ}$ = $E_{cathode}^{\circ}$ - $E_{anode}^{\circ}$ = + 0.8545 V – (+ 0.920 V) = –0.0655 V. Our knowledge of this value and “Equation 7” allow us to calculate the equilibrium constant (K_eq) of the reaction shown by “Equation 4” at 25°C. $E_{cell}^{o} = (RT / nF) ln K_{eq}$ (7)

With the values for the gas constant R = 8.314 J/mol rxn.K and Faraday’s constant F = 96485 C/mol e^-, the temperature T = 298.15 K, n = 2 mol e^-/mol rxn, and $E_{cell}^{\circ}$ = –0.0655 V, we have K_eq = 6.1×10^-3. It is worth noting that Greenwood and Earnshaw [3] also report this number (0.0061) as the value of the equilibrium constant for the reaction at 25°C.

It is, from a mineralogical perspective, important to know what percentage of the aqueous ${Hg}_{2}^{2 +}$ ions from the dissolved calomel in mines is oxidized to Hg² ⁺ . The following equilibrium calculations shown by “Equation 8” and “Equation 9” lead to determining that percentage: $K_{eq} = [{Hg}^{2 +} (aq)] / [{Hg}_{2}^{2 +} (aq)]$ (8) $6.1 \times 10^{- 3} = x / (100 - x)$ (9)

Based on the calculated value of x = 6.0×10^-1, when the reaction shown by “Equation 4” is at equilibrium, 0.60% of the aqueous mercury ions in the reaction mixture are aqueous Hg²⁺ ions. Greenwood and Earnshaw [3] state this theoretical finding from an experimentally verifiable perspective: “at equilibrium, aqueous solutions of mercury(I) salts will contain around 0.6% of mercury(II).” This occurs not only in the laboratory, but also when calomel is dissolved in water in mercury mines, because scientific laws describe nature in general, including in mines. Further, “Equation 4” is consistent with the findings that “relatively small amounts” of elemental mercury occur as droplets (Hg(l)) in mercury mines, such as those in Spain and California [14], and that mineral calomel is “often found associated with droplets” of elemental mercury (Hg(l)) [10].

That some of the mineral calomel in mines is naturally converted to a mixture of Hg²⁺ and Cl^- aqueous ions in a 1 : 2 molar ratio is indisputable, shown by “Equation 10”. ${Hg}_{2}^{2 +} (aq) + 2 {Cl}^{-} (aq) ⇄ {Hg}^{2 +} (aq) + 2 {Cl}^{-} (aq) + Hg (l)$ (10)

No doubt, each of the two processes of mercury(II) complexation by hydroxide (OH^-) [15] and mercury(II) complexation by chloride (Cl^-) [15] will decrease the concentration of aqueous Hg² ⁺ . When the total mercury(II) concentration is less than 10^-4 M, mercury(II) may occur primarily as Hg²⁺, [HgCl]⁺, [HgCl₂], [HgCl₃]^-, [HgCl₄]^2-, or [Hg(OH)₂], depending on pH and Cl^- concentration [15]. From the perspective of mercury speciation, it is worth noting that when the total mercury(II) concentration is higher than 10^-4 M, the precipitate HgO might be produced at high pH values [15]; however, that is unexpected if mineral calomel is the mercury source (considering its K_sp = 1.43×10^-18 [12]).

In light of the corresponding formation constants reported by Sawyer et al. [15], the complexation reactions are shown by “Equations 11–16”. These formation constants are also listed as β values (K values from Hg²⁺ and the ligands Cl^- and OH^- only, not incremental additions of the ligands). ${Hg}^{2 +} (aq) + {Cl}^{-} (aq) ⇄ {[HgCl]}^{+} (aq)$ (11)

K_Cl1 = 5.2×10⁶ [15]

= > β_Cl1 = 5.2×10⁶ ${[HgCl]}^{+} (aq) + {Cl}^{-} (aq) ⇄ [{HgCl}_{2}] (aq)$ (12)

K_Cl2 = 3.2×10⁶ [15]

= > β_Cl2 = 1.7×10¹³ $[{HgCl}_{2}] (aq) + {Cl}^{-} (aq) ⇄ {[{HgCl}_{3}]}^{-} (aq)$ (13)

K_Cl3 = 10 [15]

= > β_Cl3 = 1.7×10¹⁴ ${[{HgCl}_{3}]}^{-} (aq) + {Cl}^{-} (aq) ⇄ {[{HgCl}_{4}]}^{2 -} (aq)$ (14)

K_Cl4 = 9.3 [15]

= > β_Cl4 = 1.6×10¹⁵ ${Hg}^{2 +} (aq) + {OH}^{-} (aq) ⇄ {[HgOH]}^{+} (aq)$ (15)

K_OH1 = 2.0×10¹⁰ [15]

= > β_OH1 = 2.0×10¹⁰

${[HgOH]}^{+} (aq) + {OH}^{-} (aq) ⇄ [Hg {(OH)}_{2}] (aq)$ (16)

K_OH2 = 2.5×10¹¹ [15]

= > β_OH2 = 5.0×10²¹

The equilibrium constant for the net reaction shown by “Equation 17” can be calculated using “Equation 18” and is K_net1 = 1.3×10^-5, and the equilibrium constant for the net reaction shown by “Equation 19” can be calculated using “Equation 20” and is K_net2 = 44. ${Hg}_{2} {Cl}_{2} (s) + 2 {Cl}^{-} (aq) ⇄ {[{HgCl}_{4}]}^{2 -} (aq) + Hg (l)$ (17) $K_{net 1} = K_{sp} \times K_{eq} \times K_{Cl 1} \times K_{Cl 2} \times K_{Cl 3} \times K_{Cl 4}$ (18)

K_net1 = (1.43×10^-18)^×(6.1×10^-3)×(5.2×10⁶)×(3.2×10⁶)×(10)×(9.3)=1.3×10^-5 ${Hg}_{2} {Cl}_{2} (s) + 2 {OH}^{-} (aq) ⇄ [Hg {(OH)}_{2}] (aq) + 2 {Cl}^{-} (aq) + Hg (l)$ (19) $K_{net 2} = K_{sp} \times K_{eq} \times K_{OH 1} \times K_{OH 2}$ (20)

K_net2 = (1.43×10^-18)^×(6.1×10^-3)×(2.0×10¹⁰)×(2.5×10¹¹) = 44

That K_net2 is almost three-million times larger than K_net1 shows that the conversion of mineral calomel to [Hg(OH)₂] is much more favorable than the conversion of mineral calomel to [HgCl₄]^2-. Sawyer et al. [15] use a predominance area diagram to illustrate that with the total mercury(II) concentration less than 10^-4 M and the concentration of aqueous chloride between 10^-4 M and 10^-2 M, the predominant species at pH greater than 7 is [Hg(OH)₂], while at pH lower than 7 the predominant species is [HgCl₂].

Still, if mineral calomel is dissolved in water, a mixture of Hg²⁺ and Cl^- aqueous ions in a 1 : 2 molar ratio, produced in the reaction shown by “Equation 10”, will always exist. It is also worth noting that, based on Le Ch â telier’s Principle, the equilibrium shown by “Equation 10” will be displaced to the right if mercury(II) undergoes complexation. The aqueous mixture produced in “Equation 10” may be represented as HgCl₂(aq) (Hg²⁺(aq) being the complex [Hg(H₂O)₆]²⁺ [3]). This mixture will, upon the natural evaporation of water, the solvent in “Equation 10”, in mines, certainly yield solid HgCl₂, with the crystalline structure composed of linear Cl – Hg – Cl molecules [3].

The dissolution of mineral calomel in water (“Equation 1”), the conversion of aqueous ${Hg}_{2}^{2 +}$ to aqueous Hg²⁺ (“Equation 4” and “Equation 10”), and the formation of solid HgCl₂ through water evaporation are all processes that occur in mercury mines due to ordinary natural (that is, not anthropogenic) circumstances. Therefore, HgCl₂ is a naturally occurring mercury compound, whether in the aqueous phase or in the solid phase. The discovery of HgCl₂ in nature would point at the necessity to reconsider the chemical role of oxidation in the discovery of new minerals from already discovered ones. It is especially worth noting that of the other zinc family halides, mercury(I) iodide (Hg₂I₂) and mercury(II) iodide (HgI₂) are found as the minerals moschelite and coccinite, respectively [8]. The detection and quantification of solid HgCl₂ as a mineral in mercury mines will be an appropriate mineralogical study in the future.

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