Electron transfer spontaneous

Why do electrons transfer spontaneously from a Zn atom to a Cu ion, either directly as in the reaction of Figure 20.3 or through an external circuit as in the voltaic cell of Figure 20.5 In this section we will examine the "driving force" that pushes the electrons through an external circuit in a voltaic cell. 

One form of selective reaction between an unwanted ion and the reaction gas molecules is electron transfer. Spontaneous electron transfer from a reaction gas molecule to an ion occurs when the ionization energy of the former is lower than that of the neutral entity at the basis of the interfering ion (first ionization energy in the case of a singly charged ion, second in the case of a doubly charged ion). A few examples of this approach are presented below. 

Decomposition of diphenoylperoxide [6109-04-2] (40) in the presence of a fluorescer such as perylene in methylene chloride at 24°C produces chemiluminescence matching the fluorescence spectmm of the fluorescer with perylene was reported to be 10 5% (135). The reaction follows pseudo-first-order kinetics with the observed rate constant increasing with fluorescer concentration according to = k [flr]. Thus the fluorescer acts as a catalyst for peroxide decomposition, with catalytic decomposition competing with spontaneous thermal decomposition. An electron-transfer mechanism has been proposed (135). 

A voltaic cell produces electrical energy through spontaneous redox chemical reactions. When zinc metal is placed in a solution of copper sulfate, an electron transfer takes place between the zinc metal and copper ions. The driving force for the reaction is the greater attraction of the copper ions for electrons ... 

In this reaction, copper metal plates out on the surface of the zinc. The blue color of the aqueous Cu2+ ion fades as it is replaced by the colorless aqueous Zn2+ ion (Figure 18.1). Clearly, this redox reaction is spontaneous it involves electron transfer from a Zn atom to a Cu2+ ion. 

Other postulated mechanisms for spontaneous initiation include electron transfer followed by proton transfer to give two monoradicals, hydrogen atom transfer between a charge-transfer complex and solvent,110 and formation of a di radical from a charge-transfer complex, JJ[Pg.111]

Experiments and calculations both indicate that electron transfer from potassium to water is spontaneous and rapid, whereas electron transfer from silver to water does not occur. In redox terms, potassium oxidizes easily, but silver resists oxidation. Because oxidation involves the loss of electrons, these differences in reactivity of silver and potassium can be traced to how easily each metal loses electrons to become an aqueous cation. One obvious factor is their first ionization energies, which show that it takes much more energy to remove an electron from silver than from potassium 731 kJ/mol for Ag and 419 kJ/mol for K. The other alkali metals with low first ionization energies, Na, Rb, Cs, and Fr, all react violently with water. 

The reactivities of potassium and silver with water represent extremes in the spontaneity of electron-transfer reactions. The redox reaction between two other metals illustrates less drastic differences in reactivity. Figure 19-5 shows the reaction that occurs between zinc metal and an aqueous solution of copper(II) sulfate zinc slowly dissolves, and copper metal precipitates. This spontaneous reaction has a negative standard free energy change, as does the reaction of potassium with water ... 

Spontaneous redox reactions can aiso occur by indirect eiectron transfer. In an indirect electron transfer, species invoived in the redox chemistry are not aiiowed to come into direct contact with one another. Instead, the oxidation occurs at one end of a wire and transfers eiectrons to the wire. Reduction occurs at the other end of the wire and removes eiectrons from the wire. The wire conducts eiectrons between the oxidation site and the reduction site. 

In Figure 19-9. the spontaneous reaction is electron transfer between zinc and copper ... 

The units of A G are J/mol. On the right side of Equation, the Faraday constant has units of C/mol. Potential differences are in volts, and 1J=IVC, solV=lG/C and the product FE has units of J/mol. In this equation, n is dimensionless because it is a ratio, the number of electrons transferred per atom reacting. Equation has a negative sign because a spontaneous reaction has a negative value forzlG but a positive value for E. 

C19-0123. A cell is set up using two zinc wires and two solutions, one containing 0.250 M ZnCl2 solution and the other containing 1.25 M Zn (N03)2 solution, (a) What electrochemical reaction occurs at each electrode (b) Draw a molecular picture showing spontaneous electron transfer processes at the two zinc electrodes, (c) Compute the potential of this cell. 

Fig. 5.2 The n-Cd(Se,Te)/aqueous Cs2Sx/SnS solar cell. P, S, and L indicate the direction of electron flow through the photoelectrode, tin electrode, and external load, respectively (a) in an illuminated cell and (b) in the dark. For electrolytes, m represents molal. Electron transfer is driven both through an external load and also into electrochemical storage by reduction of SnS to metaUic tin. In the dark, the potential drop below that of tin sulfide reduction induces spontaneous oxidation of tin and electron flow through the external load. Independent of illumination conditions, electrons are driven through the load in the same direction, ensuring continuous power output. (Reproduced with permission from Macmillan Publishers Ltd [Nature] [60], Copyright 2009)...

MEMED has also been used to investigate the nature of coupled ion-transfer processes involved in spontaneous electron transfer at ITIES [80]. In this application, a key strength of MEMED is that all of the reactants and products involved in the reaction can be measured, as shown in Figs. 19 and 20. The redox reaction studied involved the oxidation of either ferrocene (Fc) or decamethylferrocene (DMFc) in a DCE phase (denoted by Fcdce) by either IrCle or Fe(CN)g in the aqueous phase (denoted by Ox ) ... 

The difference in magnitudes of the currents for ion transfers coupled with the electron transfer from DMFC in NB to FMN in W is responsible for that Na" " transfers from NB to W spontaneously in the system of Eq. (5), though K" " does not. 

Flavin-cyclobutane pyrimidine dimer and flavin-oxetane model compounds like 1-3 showed for the first time that a reduced and deprotonated flavin is a strong photo-reductant even outside a protein environment, able to transfer an extra electron to cyclobutane pyrimidine dimers and oxetanes. There then spontaneously perform either a [2n+2n cycloreversion or a retro-Paternd-Buchi reaction. In this sense, the model compounds mimic the electron transfer driven DNA repair process of CPD- and (6-4)-photolyases. 

If the EDA and CT pre-equilibria are fast relative to such a (follow-up) process, the overall second-order rate constant is k2 = eda c e In this kinetic situation, the ion-radical pair might not be experimentally observed in a thermally activated adiabatic process. However, photochemical (laser) activation via the deliberate irradiation of the charge-transfer absorption (hvct) will lead to the spontaneous generation of the ion-radical pair (equations 4, 5) that is experimentally observable if the time-resolution of the laser pulse exceeds that of the follow-up processes (kf and /tBet)- Indeed, charge-transfer activation provides the basis for the experimental demonstration of the viability of the electron-transfer paradigm in Scheme l.21... 

Electron-transfer activation. The observation of intense coloration upon mixing the solutions of hydroquinone ether MA and nitrogen dioxide at low temperature derives from the transient formation of MA+ cation radical, as confirmed by the spectral comparison with the authentic sample. The oxidation of MA to the corresponding cation radical is effected by the nitrosonium oxidant, which is spontaneously generated during the arene-induced disproportionation of nitrogen dioxide,239 i.e.,... 

EAa is limited because the molecules, as well as their cations or anions, must be stable in ambient air or solvent donors that are too powerful as reducers, and acceptors that are too powerful as oxidizers, will not persist until their rectification is measured. Further, the donor s HOMO must not be so high, or its LUMO so low, that spontaneous electron transfer occurs to convert the molecule into a zwitterion. 

SN P spontaneously releases N O both thermally and photochemically [61-65], but is quite stable in the dark and in aqueous in vitro physiological media [66]. This implies that absorption of heat and light energy induces electron transfer from the Fe2+ center to the N 0+ ligand, resulting in weakening of the Fe-N O bond and subsequent release of NO [65]. SNP also decomposes in an aqueous environment in the presence of biological reductants [65, 66] and some transition metal ions to produce nitric oxide. 

Let s look at the little strip cartoon in Figure 7.7, which shows the surface of a copper electrode. For clarity, we have drawn only one of the trillion or so atoms on its surface. When the cell of which it is a part is permitted to discharge spontaneously, the copper electrode acquires a negative charge in consequence of an oxidative electron-transfer reaction (the reverse of Equation (7.7)). During the oxidation, the surface-bound atom loses the two electrons needed to bond the atom to the electrode surface, becomes a cation and diffuses into the bulk of the solution. 

We now consider a slightly different cell in which the copper half-cell is the positive pole. Perhaps the negative electrode is zinc metal in contact with Zn2+ ions. If the cell discharges spontaneously, then the electron-transfer reaction is the reduction reaction in Equation (7.7) as depicted in the strip cartoon in Figure 7.8. A bond forms between the surface of the copper electrode and a Cu2+ cation in the solution The electrons needed to reduce the cation come from the electrode, imparting a net positive charge to its surface. 

A first turning point in the dichotomy between radical and ionic chemistry is located at the level of the primary radical, usually an ion radical, formed upon single electron transfer to the substrate. If, for a reduction, the reaction medium is not too acidic (or electrophilic), and for an oxidation, not too basic (or nucleophilic), radical reactions involving the primary radical, such as self-coupling, have a first opportunity to compete successfully with acid-base reactions. In this competition, the acidity (for a reduction) or basicity (for an oxidation) of the substrate should also be taken into account insofar as they may lead to father-son acid-base reactions. It should also be taken into consideration that the primary radical may undergo spontaneous acid-base reactions such as expelling a base (or a nucleophile) after a reduction, and an acid (or an electrophile) after an oxidation. 

If the provoked or spontaneous acid-base reactions overcome the radical reactions of the primary radical, the secondary radical is easier to reduce, or to oxidize, than the substrate in most cases. Exceptions to this rule are scarce, but exist. They involve substrates that are particularly easy to reduce thanks to the presence of a strongly electron-withdrawing substituent (for reductions, electron-donating for oxidation), which is expelled upon electron transfer, thus producing a radical that lacks the same activation. Alkyl iodides and aryl diazonium cations are typical examples of such systems. 

The mechanism of electrochemical reduction of nitrosobenzene to phenylhydroxylamine in aqueous medium has been examined in the pH range from 0.4 to 13, by polaro-graphic and cyclic voltametry. The two-electron process has been explained in terms of a nine-membered square scheme involving protonations and electron transfer steps565. This process is part of the overall reduction of nitrobenzene to phenylhydroxylamine, shown in reaction 37 (Section VI.B.2). Nitrosobenzene undergoes spontaneous reaction at pH > 13, yielding azoxybenzene471. 

Electron-transfer reactions, especially step (4), are too fast and the intermediates too fleeting to be properly characterized and discriminated they are, moreover, contaminated by recycling of the free ferric cytochrome (Fe " ) and by spontaneous nonhydroxylating decay of the oxyferrous compound. The reaction cycle lacks any consistent knowledge of the nature and reactivity of unstable intermediates occurring beyond the metastable oxyferrous compound (O2—Fe) RH. 

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