Water standard reduction potentials

Fig. 1. Reduction potential E (referenced to the standard hydrogen electrode) versus pH for various species of vanadium. Boundary lines correspond to E, pH values where the species in adjacent regions are present in equal concentrations. The short dashed lines indicate uncertainty in the location of the boundary. The upper and lower long dashed lines correspond to the upper and lower limits of stability of water. Standard reduction potentials are given by the intersections of horizontal lines with the abscissa pH = 0. The half reactions are 02 + 4H+ + 4e = 2H20, E° = 1.23V V02+ + 2H+ + e = V02+ + H20, E° = 1.0V V02+ + 2H+ + e = V3+ + H,0, E° = 0.36V 2H+ + 2e = H2, E° = 0.0V and V3+ + e = V2+, E° = -0.25V. V2+ is therefore a strong reductant. Air oxidation of V02+ presumably proceeds by the reaction 4V02t + 02 + 2H20 = 4VOJ + 4H+, E° = 0.23V which is favored at higher pH. Not all known species are represented on this diagram. Reproduced with permission from Ref. 30...

The standard reduction potential of Cr " (Table 2) shows that this ion is a strong reducing agent, and Cr(II) compounds have been used as reagents in analytical chemistry procedures (26). The reduction potential also explains why Cr(II) compounds are unstable in aqueous solutions. In the presence of air, the oxidation to Cr(III) occurs by reaction with oxygen. However, Cr(II) also reacts with water in deoxygenated solutions, depending on acidity and the anion present, to produce H2 and Cr(III) (27,28). 

In addition to simple dissolution, ionic dissociation and solvolysis, two further classes of reaction are of pre-eminent importance in aqueous solution chemistry, namely acid-base reactions (p. 48) and oxidation-reduction reactions. In water, the oxygen atom is in its lowest oxidation state (—2). Standard reduction potentials (p. 435) of oxygen in acid and alkaline solution are listed in Table 14.10- and shown diagramatically in the scheme opposite. It is important to remember that if or OH appear in the electrode half-reaction, then the electrode potential will change markedly with the pH. Thus for the first reaction in Table 14.10 O2 -I-4H+ -I- 4e 2H2O, although E° = 1.229 V,... 

For aqueous solutions, ascorbate can be included in the hierarchy, while a-tocopherol has to be replaced by its water-soluble analogue trolox, which is often assumed to have the same standard reduction potential. The ordering of the antioxidants based on the two different determinations of E in water is rather similar, and it should be noted that ascorbate is the antioxidant which will regenerate the other antioxidants, with the ascorbate itself ending up being oxidised. In contrast to what was observed for DMF, the ordering in water predicts that quercetin could regenerate a-tocopherol from its oxidised form. 

Fig. 16.5 Synergistic regeneration of a-tocopherol by quercetin at a lipid-water interphase. a-tocopherol is reacting with a lipid peroxyl radical in a chain-breaking reaction. According to the standard reduction potential, the phenoxyl radical of quercetin can further be regenerated by ascorbate.

The IrIV anion, [Ir(H20)Br5], oxidises ascorbic acid at 20.0 °C.51 This reaction is first order with respect to ascorbic acid concentration and first order with respect to the Irlv anion. Comparison of hexabromo-, hexachloro, aquopentachloro-, and di-aquotetracholoiridium(IV) reactions with ascorbic acid shows that replacing a halide ion with a water molecule increases the standard reduction potential of the IrIV complex and increases the rate of reaction. 

FIGURE3.7 The potential window for the redox chemistry of life. Redox chemistry in living cells is approximately limited by the standard potentials for reduction and oxidation of the solvent water at neutral pH. Approximate standard reduction potentials are also indicated for the commonly used oxidant ferricyanide and reductants NADH and dithionite. 

It is not possible to prepare F2 by electrolysis of an aqueous NaF solution. In electrolysis, the most easily oxidized and reduced species are the ones involved. To prepare F2, the oxidation of F would have to occur. However, water is more easily oxidized than is F, as seen by its position in the standard reduction potential chart (Appendix J and below). By inspection, H20 is a stronger reducing agent than F because the reduction half-reaction has a less positive E°. So H20 s oxidation is preferable to F s oxidation. F2 can be prepared from molten NaF, but not aqueous NaF. 

When water is electrolyzed with copper electrodes or using other common metals, the amount of 02(g) is less than when Pt electrodes are used, but the amount of H2(g) produced is independent of electrode material. Why does this happen In electrolysis, the most easily oxidized species is oxidized and the most easily reduced species is reduced. If we compare Cu and H20 by looking on the standard reduction potentials chart (data given below), we see that Cu is a stronger reducing agent than H20, because 0.337 V is less than 0.828 V. This means that Cu is more easily oxidized than water. 

The standard reduction potential for Be2+ is the least negative of the elements in the group and by the same token beryllium is the least electropositive and has the greatest tendency to form covalent bonds. The bulk metal is relatively inert at room temperature and is not attacked by air or water at high temperatures. Beryllium powder is somewhat more reactive. The metal is passivated by cold concentrated nitric acid but dissolves in both dilute acid and alkaline solutions with the evolution of dihydrogen. The metal reacts with halogens at 600°C to form the corresponding dihalides. 

The standard reduction potentials used to calculate E ceii for the decomposition of water apply only to reactants and products in their standard states. However, in pure water at 25°C, the hydrogen ions and hydroxide ions each have concentrations of 1 x 10 mol/L. This is not the standard state value of 1 mol/L. The reduction potential values for the non-standard conditions in pure water are given below. The superscript zero is now omitted from the E symbol, because the values are no longer standard. 

This is an aqueous solution. You are given the formula and concentration of the electrolyte. You have a table of standard reduction potentials, and you know the non-standard reduction potentials for water. 

Step 1 The Lh and Br concentrations are 1 mol/L, so use the standard reduction potentials for the half-reactions that involve these ions. Use the non-standard values for water. 

Use the relevant standard reduction potentials from the table in Appendix E, and the non-standard reduction potentials you used previously for water, to predict the electrolysis products. Predict which product(s) are formed at the anode and which product(s) are formed at the cathode. 

It is possible to use the standard reduction potentials for the reduction of hydrogen ions and the reduction of water molecules to show that the dissociation of water molecules into hydrogen ions and hydroxide ions is non-spontaneous under standard conditions. Describe how you would do this. How is this result consistent with the observed concentrations of hydrogen ions and hydroxide ions in pure water ... 

Which one of the following metals does not react with water to produce hydrogen (Use Appendix D for standard reduction potentials)... 

Although reduction potentials may be estimated for half-reactions, there are limits for their values that correspond to both members of a couple having stability in an aqueous system with respect to reaction with water. For example, the Na+/Na couple has a standard reduction potential of -2.71 V, but metallic sodium reduces water to dihydrogen. The reduced form of the couple (Na) is not stable in water. The standard reduction potential for the Co3 + / Co2 + couple is +1.92 V, but a solution of Co3+ slowly oxidizes water to dioxygen. In this case the oxidized form of the couple is not stable in water. The standard reduction potential for the Fe3T/Fe2+ couple is +0.771 V, and neither oxidized form or reduced form react chemically with water. They are subject to hydrolysis, but are otherwise both stable in the aqueous system. The limits for the stability of both oxidized and reduced forms of a couple are pH dependent,... 

The standard reduction potentials for the main species formed by the Group 17 elements in aqueous solution are given in Tables 6.16 and 6.17, for pH values 0 and 14, respectively. Irrespective of the pH of the solution, the halogen elements range from the extremely powerful F2 (which has the potential to oxidize water to dioxygen), through the powerful oxidants Cl2 and Br2, to 12, which is a relatively weak oxidant. 

The nature of ions in solution is described in some detail and enthalpies and entropies of hydration of many ions are defined and recalculated from the best data available. These values are used to provide an understanding of the periodicities of standard reduction potentials. Standard reduction potential data for all of the elements, group-bygroup, covering the s-and p-, d- and/- blocks of the Periodic Table is also included. Major sections are devoted to the acid/base behaviour and the solubilities of inorganic compounds in water. 

The rate constant for the redox step, kr, is unlikely to reflect a simple electron transfer from the monodentate diimine ligand to the metal center because replacement of coordinated water with coordinated hydroxide would be expected to decrease the oxidizing power of the metal-(III) center. This is well documented by the standard reduction potentials of the aqua and hydroxo complexes in Table IV, and it would seem... 

---