Chapter 2 Electrochemistry of the Chalcogens 2.1 General References Because of their multiple oxidation states, the chalcogens, particularly sulfur, can engage in numerous redox couples participating in acid–base, oxidation–reduction, precipitation, and complexation equilibria. In the anion electrochemical series, sulfur, being the less noble element com- pared to its heavier congeners, occupies an intermediate position between iodine and selenium [(+)F, Cl, Br, I, S, Se, Te(–)]. Selenium, regarded as a metalloid, is a relatively noble element. Tellurium is rather an amphoteric element: it can enter into solution in the form of both cations and anions. Regarded as a metal, i.e., with respect to its cations, tellurium occupies a position between copper and mercury. Regarded as a metalloid, i.e., with respect to its anions, it is located on the extreme right of the above series. A comprehensive survey of the classical electrochemical facts for sulfur, sele- nium, and tellurium, as documented until about 1970, can be found in the reviews of Zhdanov [1], wherefrom we cite the lists of standard and formal potentials for aqueous solutions, in Tables 2.1, 2.2, and 2.3, respectively. Many of these potentials have been calculated thermodynamically since the experimental determinations are few. The listed data are largely drawn from the monograph by Pourbaix (below), and the interested reader should validate the measurement conditions for the indicated potentials or the scatter in their values (not given here in detail). In these tables, the redox systems are assorted by decreasing formal valency of chalcogen in the oxi- dized state, while at a given valency of the oxidized state they appear in the order of decreasing valency of chalcogen in the reduced state. Latimer [2] has compiled useful aqueous redox transition potential diagrams (reproduced also in Zhdanov’s monographs) that are convenient as a quick guide in practical problems and for perceiving the oxidation–reduction properties of some chalcogen hydride and oxy- chalcogenide species. Standard potentials of chalcogens in non-aqueous media are generally not known, at least in a systematic manner. A standard approach for the theoretical presentation of electrochemical equilibria is the use of Pourbaix, or potential–pH predominance area diagrams, which incor- porate chemical and electrochemical thermodynamics simultaneously in a straight- forward manner. These diagrams comprise an extremely useful, yet fundamental, M. Bouroushian, Electrochemistry of Metal Chalcogenides, Monographs 57 in Electrochemistry, DOI 10.1007/978-3-642-03967-6_2, C Springer-Verlag Berlin Heidelberg 2010 58 2 Electrochemistry of the Chalcogens starting point for the study of electrochemical systems. Certainly, the ability to predict, understand, and ultimately control electrochemical reactions requires also knowledge of process kinetics. Reproduced below are the potential–pH diagrams for the chalcogen–water systems as originally derived and illustrated by Pourbaix [3], along with some accompanying data. These will serve as a guide to all subsequent discussion on aqueous systems used for electrochemical preparations of metal chalcogenides. The reader may take notice of the chalcogen-containing substances that are considered to be present in the electrolytes and their stability domains. The diagrams repre- sent, almost exclusively, the important “valencies” of the chalcogens, namely –2, +4, and +6 (and not effective oxidation states corresponding, e.g., to polysulfides and the like). Still, all known substances for each chalcogen, either stable in water or not, are registered here, in order to serve as a useful basis for the formulae and nomenclature used in the rest of the book. Note that the parts of the Pourbaix dia- grams outside the zero activity lines (log C = 0) regard ideal solutions containing 1 M of dissolved chalcogen in the forms considered, and the parts inside these lines regard solutions saturated with solid chalcogen (plus TeO2 in the case of tellurium). The diagrams are valid only in the absence of substances with which the respective chalcogen can form soluble complexes or insoluble salts. It is considered useful to include here the potential–pH diagram for some redox systems related to oxygen (Fig. 2.1) [4]. Lines 11 and 33 correspond to the (a) and (b) dashed lines bounding the stability region of water, as depicted in all the subsequent Pourbaix diagrams. Fig. 2.1 Potential–pH diagram for oxygen reactions. (L’ Her [4] Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission) 2.1 General References 59 2.1.1 Tables of Aqueous Standard and Formal Potentials Table 2.1 Sulfur reactions Standard or formal potential Half-reaction (V vs. SHE at 25 ◦C) 2− + − → 2− + S2O8 2e 2SO4 2.01 2− + + + − → − + S2O8 2H 2e 2HSO4 2.123 2− + + + − → 2− + − 2SO4 4H 2e S2O6 2H2O 0.22 2− + + + − → + + SO4 4H 2e H2SO3 H2O 0.17 2− + + + − → + + SO4 4H 2e SO2 2H2O 0.138 2− + + − → 2− + − − SO4 H2O 2e SO3 2OH 0.93 − + + + − → + + HSO4 7H 6e S(s) 4H2O 0.339 2− + + + − → + + SO4 8H 6e S(s) 4H2O 0.357 − + + + − → + + HSO4 9H 8e H2S(aq) 4H2O 0.289 2− + + + − → + + SO4 10H 8e H2S(aq) 4H2O 0.303 2− + + + − → − + + SO4 9H 8e HS 4H2O 0.252 2− + + + − → 2− + + SO4 8H 8e S 4H2O 0.149 2− + + + − → + + SO4 10H 8e H2S(g) 4H2O 0.311 2− + + + − → + S2O6 4H 2e 2H2SO3 0.57 2− + + + − → − + S2O6 2H 2e 2HSO3 0.455 2− + − → 2− + S2O6 2e 2SO3 0.026 + − → 2− + + 3H2SO3 2e S3O6 3H2O 0.30 + + + − → − + − − 2H2SO3 H 2e HS2O4 2H2O 0.08, 0.056 2− + + − → 2− + − − 2SO3 2H2O 2e S2O4 4OH 1.12 2− + + + − → 2− + + 2SO3 4H 2e S2O4 2H2O 0.416 − + + + − → − + + 2HSO3 3H 2e HS2O4 2H2O 0.060 − + + + − → 2− + − − 2HSO3 2H 2e S2O4 2H2O 0.013, 0.009 + + + − → 2− + + 4H2SO3 4H 6e S4O6 6H2O 0.51 − + + + − → 2− + + 4HSO3 8H 6e S4O6 6H2O 0.581 + + + − → −2 + + 4SO2(g) 4H 6e S4O6 2H2O 0.510 + + + − → 2− + + 2H2SO3 2H 4e S2O3 3H2O 0.40 2− + + − → 2− + − − 2SO3 3H2O 4e S2O3 6OH 0.58 2− + + + − → 2− + + 2SO3 6H 4e S2O3 3H2O 0.705 + − H2SO3 + 2H + 2e → H2SO2 + H2O<+0.4 − + + + − → 2− + + 2HSO3 4H 4e S2O3 3H2O 0.491 + + + − → 2− + + 5H2SO3 8H 10e S5O6 9H2O 0.41 + − H2SO3 + 4H + 4e → S(s) + 3H2O +0.45 2− − − SO + 3H2O + 4e → S(s) + 6OH −0.66 3 + − SO2(g) + 4H + 4e → S(s) + 2H2O +0.451, +0.470 2− + + + − → 2− + + SO3 6H 6e S 3H2O 0.231 2− + − → 2− + + − S4O6 2e 2S2O3 0.08, 0.219, 0.10 2− + + + − → + + S4O6 12H 10e 4S(s) 6H2O 0.416 60 2 Electrochemistry of the Chalcogens Table 2.1 (continued) Standard or formal potential Half-reaction (V vs. SHE at 25 ◦C) + − H2SO2 + 2H + 2e → S(s) + 2H2O>+0.5 2− + + + − → + + S2O3 6H 4e 2S(s) 3H2O 0.465 2− + + + − → + + S5O6 12H 10e 5S(s) 6H2O 0.484 + − SO(g) + 2H + 2e → S(s) + H2O +1.507 2− + + + − → 2− + + 5S2O3 30H 24e 2S5 15H2O 0.331 2− + + + − → − + + S2O3 8H 8e 2HS 3H2O 0.200 2− + + + − → 2− + − S2O3 6H 8e 2S 3H2O 0.006 − − S2Cl2 + 2e → 2S(s) + 2Cl +1.23 + − → 2− − − 5S(s) 2e S5 0.340, 0.315 + − → 2− − 4S(s) 2e S4 0.33 + − S(s) + 2H + 2e → H2S(aq) +0.141 + − S(s) + 2H + 2e → H2S(g) +0.171 S(s) + H+ + 2e− → HS− −0.065 − − − S(s) + H2O + 2e → HS + OH −0.52 S(s) + 2e− → S2− A large number of measurements and calculated values are available. These vary from −0.48 to −0.58 V, but most are closer to −0.48 V 2− + − → 2− − 4S5 2e 5S4 0.441 2− + + + − → − + S5 5H 8e 5HS 0.003 2− + + + − → + S5 10H 8e 5H2S(g) 0.299 2− + − → 2− − 3S4 2e 4S3 0.478 2− + − → 2− + 2− − S4 2e S S3 0.52 2− + + + − → − + S4 4H 6e 4HS 0.033 2− + − → 2− − 2S3 2e 3S2 0.506 2− + − → 2− + 2− − S3 2e S S2 0.49 2− + + + − → − + S3 3H 4e 3HS 0.097 2− + − → 2− − − S2 2e 2S 0.48, 0.524 2− + + + − → − + S2 2H 2e 2HS 0.298 Table 2.2 Selenium reactions Standard or formal potential Half-reaction (V vs. SHE at 25 ◦C) − + + + − → + + HSeO4 3H 2e H2SeO3 H2O 1.090 2− + + + − → + + SeO4 4H 2e H2SeO3 H2O 1.15 2− + + + − → − + + SeO4 3H 2e HSeO3 H2O 1.075 2− + + + − → 2− + + SeO4 2H 2e SeO3 H2O 0.880 2− + − + → 2− + − + SeO4 2e H2O SeO3 2OH 0.05 2.1 General References 61 Table 2.2 (continued) Standard or formal potential Half-reaction (V vs. SHE at 25 ◦C) + − H2SeO3 + 4H + 4e → Se(s) + 3H2O +0.740 − + + + − → + + HSeO3 5H 4e Se(s) 3H2O 0.778 2− + + + − → + + SeO3 6H 4e Se(s) 3H2O 0.875 2− + − + → + − − SeO3 4e 3H2O Se(s) 6OH 0.366 Se4+ + 4e− → Se(s) +0.846 + − H2SeO3 + 6H + 6e → H2Se + 3H2O +0.360 − + + + − → + + HSeO3 7H 6e H2Se 3H2O 0.386 − + + + − → − + + HSeO3 6H 6e HSe 3H2O 0.349 2− + + + − → − + + SeO3 7H 6e HSe 3H2O 0.414 2− + + + − → 2− + + SeO3 6H 6e Se 3H2O 0.276 − − Se2Cl2 + 2e → 2Se(s) + 2Cl +1.1 + − Se(s) + 2H + 2e → H2Se −0.40 Se(s) + H+ + 2e− → HSe− −0.510 Se(s) + 2e− → Se2− −0.92 + − Se(s) + 2H + 2e → H2Se(g) −0.369 Table 2.3 Tellurium reactions [TeO2aq(s) refers to hydrated TeO2 crystals] Standard or formal potential Half-reaction (V vs. SHE at 25 ◦C) + − 4+ H2TeO4 + 6H + 2e → Te + 4H2O +0.920 + + + − → ++ + H2TeO4 3H 2e HTeO2 2H2O 0.953 + + + − → −+ + H2TeO4 H 2e HTeO3 H2O 0.631 − + + + − → −+ + HTeO4 2H 2e HTeO3 H2O 0.813 − + + + − → 2−+ + HTeO4 H 2e TeO3 H2O 0.584 2− + + + − → 2−+ + TeO4 2H 2e TeO3 H2O 0.892 + − H2TeO4 + 2H + 2e → TeO2(s) + 2H2O +1.020 + − H2TeO4 + 2H + 2e → TeO2aq(s) + 2H2O +0.854 − + + + − → + + HTeO4 3H 2e TeO2(s) 2H2O 1.202 − + + + − → + + HTeO4 3H 2e TeO2aq(s) 2H2O 1.036 2− + + + − → + + TeO4 4H 2e TeO2(s) 2H2O 1.509 2− + + + − → + + TeO4 4H 2e TeO2aq(s) 2H2O 1.343 + − TeO3(s) + 2H + 2e → TeO2(s) + H2O +1.020 + − TeO3(s) + 2H + 2e → TeO2aq(s) + H2O +0.850 Te4+ + 4e− → Te(s) +0.584 (2–3 M HCI), +0.568, +0.556 + + + + − → + + HTeO2 3H 4e Te(s) 2H2O 0.551 + − H2TeO3 + 4H + 4e → Te(s) + 3H2O +0.589 − + + + − → + + HTeO3 5H 4e Te(s) 3H2O 0.713 2− + + + − → + + TeO3 6H 4e Te(s) 3H2O 0.827 + − TeO2(s) + 4H + 4e → Te(s) + 2H2O +0.521 + − TeO2aq(s) + 4H + 4e → Te(s) + 2H2O +0.604 62 2 Electrochemistry of the Chalcogens Table 2.3 (continued) Standard or formal potential Half-reaction (V vs.