Aqueous solution equilibrium potential

The electronic spectrum of a compound arises from its 7r-electron system which, to a first approximation, is unaffected by substitution of an alkyl group for a hydrogen atom. Thus, comparison of the ultraviolet spectrum of a potentially tautomeric compound with the spectra of both alkylated forms often indicates which tautomer predominates. For example, Fig. 1 shows that 4-mercaptopyridine exists predominantly as pyrid-4-thione. In favorable cases, i.e., when the spectra of the two alkylated forms are very different and/or there are appreciable amounts of both forms present at equilibrium, the tautomeric constant can be evaluated. By using this method, it was shown, for example, that 6-hydroxyquinoline exists essentially as such in ethanol but that it is in equilibrium with about 1% of the zwitterion form in aqueous solution (Fig. 2). 

The standard electrode potentials , or the standard chemical potentials /X , may be used to calculate the free energy decrease —AG and the equilibrium constant /T of a corrosion reaction (see Appendix 20.2). Any corrosion reaction in aqueous solution must involve oxidation of the metal and reduction of a species in solution (an electron acceptor) with consequent electron transfer between the two reactants. Thus the corrosion of zinc ( In +zzn = —0-76 V) in a reducing acid of pH = 4 (a = 10 ) may be represented by the reaction ... 

Table 1.7 shows typical half reactions for the oxidation of a metal M in aqueous solutions with the formation of aquo cations, solid hydroxides or aquo anions. The equilibrium potential for each half reaction can be evaluated from the chemical potentials of the species involved see Appendix 20.2) and it should be noted that there is no difference thermodynamically between equations 2(a) and 2(b) nor between 3(a) and 3(b) when account is taken of the chemical potentials of the different species involved. 

Immunity the state of a metal whose corrosion rate is low or negligible because its potential is below (less positive than) that of equilibrium with a very small concentration (or activity of its dissolved ions. The metal is thus regarded as thermodynamically stable. Pourbaix has suggested that the small metal ion concentration be 10 mol dm (Atlas of Electrochemical Equilibria in Aqueous Solutions, p. 71, Pergamon/ CEBELCOR, Oxford (1966)). 

Potential-pH Equilibrium Diagram (Pourbaix Diagram) diagram of the equilibrium potentials of electrochemical reactions as a function of the pH of the solution. The diagram shows the phases that are thermodynamically stable when a metal reacts with water or an aqueous solution of specified ions. 

Some metals are thermodynamically unstable in aqueous solutions because their equilibrium potential is more negative than the potential of the reversible hydrogen electrode in the same solution. At such electrodes, anodic metal dissolution and cathodic hydrogen evolution can occur as coupled reactions, and their open-circuit potential (OCP) will be more positive than the equilibrium potential (see Section 13.7). 

Their potentials in 0.1 N, lmolal, IN and saturated KC1 solutions are 0.3337, 0.2800, 0.2897 and 0.2415 V, respectively. The dilute types reach their equilibrium potentials more quickly and these potentials are less dependent on temperature the SCE has the advantage of being less sensitive to current flow (electrolysis). The AgCl-Ag electrodes are more compact, do not need a liquid function, which makes them exceedingly attractive for analysis in non-aqueous media, and support high temperatures. 

The activity coefficient of the solvent remains close to unity up to quite high electrolyte concentrations e.g. the activity coefficient for water in an aqueous solution of 2 m KC1 at 25°C equals y0x = 1.004, while the value for potassium chloride in this solution is y tX = 0.614, indicating a quite large deviation from the ideal behaviour. Thus, the activity coefficient of the solvent is not a suitable characteristic of the real behaviour of solutions of electrolytes. If the deviation from ideal behaviour is to be expressed in terms of quantities connected with the solvent, then the osmotic coefficient is employed. The osmotic pressure of the system is denoted as jz and the hypothetical osmotic pressure of a solution with the same composition that would behave ideally as jt. The equations for the osmotic pressures jt and jt are obtained from the equilibrium condition of the pure solvent and of the solution. Under equilibrium conditions the chemical potential of the pure solvent, which is equal to the standard chemical potential at the pressure p, is equal to the chemical potential of the solvent in the solution under the osmotic pressure jt,... 

Because of the adsorption equilibrium for H+ and OFT ions between the surface of semiconductors and an aqueous (aq) solution, the semiconductor surface attains the point of zero charge (PZC). The flat-band potential U[h of most semiconductors including all oxides and also other compounds such as n- and p-type GaAs, p-type GaP, and n- and p-type InP in an aqueous solution is determined solely by pH and shifts proportionately with pH with a slope of -59 mV/decade, that is, pH, for example,... 

Reaction (1) is a reversible process, and it can be a source of superoxide if only its equilibrium is shifted to the right. The estimation of the equilibrium constant for this reaction in aqueous solution is impossible because the reduction potential of water-insoluble ubiquinone in water is of course undetectable. However, Reaction (1) occurs in the mitochondrial membrane and therefore, the data for the aqueous solutions are irrelevant for the measurement of its equilibrium. Some time back we studied Reaction (1) in aprotic media and found out that Ki is about 0.4 [23]. As the ubiquinone concentration in mitochondria is very high (it is about... 

The simultaneous transfer of four electrons is unlikely, and the overall reaction must contain several steps. An important intermediate is hydrogen peroxide H2O2, and its occurrence makes it difficult to establish even the equilibrium potential experimentally. The reaction is further complicated by the fact that in aqueous solutions almost all metals are covered by an oxide film in the potential range over which the reduction occurs. 

Pyridinecarboxaldehyde, 3. Possible hydration of the aldehyde group makes the aqueous solution chemistry of 3 potentially more complex and interesting than the other compounds. Hydration is less extensive with 3 than 4-pyridinecarboxaldehyde but upon protonation, about 80% will exist as the hydrate (gem-diol). The calculated distribution of species as a function of pH is given in Figure 4 based on the equilibrium constants determined by Laviron (9). 

Equilibrium between simple salts and aqueous solutions is often relatively easily demonstrated in the laboratory when the composition of the solid is invariant, such as occurs in the KCI-H2O system. However, when an additional component which coprecipitates is added to the system, the solid composition is no longer invariant. Very long times may be required to reach equilibrium when the reaction path requires shifts in the composition of both the solution and solid. Equilibrium is not established until the solid composition is homogeneous and the chemical potentials of all components between solid and aqueous phases are equivalent. As a result, equilibrium is rarely demonstrated with a solid solution series. 

Stoichiometric saturation defines equilibrium between an aqueous solution and homogeneous multi-component solid of fixed composition (10). At stoichiometric saturation the composition of the solid remains fixed even though the mineral is part of a continuous compositional series. Since, in this case, the composition of the solid is invariant, the solid may be treated as a one-component phase and Equation 6 is the only equilibrium criteria applicable. Equations 1 and 2 no longer apply at stoichiometric saturation because, owing to kinetic restrictions, the solid and saturated solution compositions are not free to change in establishing an equivalence of individual component chemical potentials between solid and aqueous solution. The equilibrium constant, K(x), is defined identically for both equilibrium and stoichiometric saturation. 

Oxazolidines are five-membered cyclic ft-Mannich bases, some of which have, indeed, been examined as potential produgs of /l-amino alcohols of medicinal relevance such as ephedrines and /3-blockers. For example, 3,4-dime-thyl-5-phenyloxazolidine (11.106), the oxazolidine of ephedrine (11.107) undergoes hydrolysis to ephedrine and formaldehyde slowly at pH 1 and 12, but very rapidly in the neutral pH range (tm < 1 min at 37°) [135], Interestingly, the equilibrium reached between the reactants and products of hydrolysis was markedly pH- and concentration-dependent. However, despite its poor stability in aqueous solution, the oxazolidine was delivered through human skin significantly faster than ephedrine when applied as 1% aqueous solutions of pH 7 - 11. The lower basicity of the oxazolidine (pKa 5.5) compared to that of ephedrine (pKa 9.6) may explain the efficient skin permeation. 

In aqueous solutions, the peak potentials of the oxidation of thiols vary with pH (Aiip/ApH = 60 mV), reflecting the position of the acid-base equilibrium affecting the SH group. In basic solutions. 

Next, we consider the interface M/S of a nonpolarizable electrode where electron or ion transfer is in equilibrium between a solid metal M and an aqueous solution S. Here, the interfadal potential is determined by the charge transfer equilibrium. As shown in Fig. 4-9, the electron transfer equilibrium equates the Fermi level, Enn) (= P (M)), of electrons in the metal with the Fermi level, erredox) (= P s)), of redox electrons in hydrated redox particles in the solution this gives rise to the inner and the outer potential differences, and respectively, as shown in Eqn. 4-10 ... 

For ion adsoiption in equilibrium on the electrode interface, the electrochemical potential Pi.., of hydrated adsorbate ions in aqueous solution equals the electrochemical potential Pi..d of adsorbed ions as shown in Eqn. 5-20 ... 

We consider dehydration-adsorption of hydrated protons (cathodic proton transfer) and desorption-hydration of adsorbed protons (anodic proton transfer) on the interface of semiconductor electrodes. Since these adsorption and desorption of protons are ion transfer processes across the compact layer at the interface of semiconductor electrodes, the adsorption-desorption equilibrium is expressed as a function of the potential of the compact layer in the same way as Eqns. 9-60 and 9-61. In contrast to metal electrodes where changes with the electrode potential, semiconductor electrodes in the state of band edge level pinning maintain the potential d(hi of the compact layer constant and independent of the electrode potential. The concentration of adsorbed protons, Ch , is then determined not by the electrode potential but by the concentration of h3 ated protons in aqueous solutions. 

---