Oxidation-reduction reactions spontaneous

As discussed in Chapter 4, oxidation is the loss of electrons in a chemical reaction, and reduction is the gain of electrons, ooo (Section 4.4) Thus, oxidation-reduction (redox) reactions occur when electrons are transferred from an atom that is oxidized to an atom that is reduced. Redox reactions are involved not only in the operation of batteries but also in a wide variety of important natural processes, including the rusting of iron, the browning of foods, and the respiration of animals. Electrochemistry is the study of the relationships between electricity and chemical reactions. It includes the study of both spontaneous and nonspontaneous processes. 

In electroless deposition, the substrate, prepared in the same manner as in electroplating (qv), is immersed in a solution containing the desired film components (see Electroless plating). The solutions generally used contain soluble nickel salts, hypophosphite, and organic compounds, and plating occurs by a spontaneous reduction of the metal ions by the hypophosphite at the substrate surface, which is presumed to catalyze the oxidation—reduction reaction. 

N204 also forms expl mixts with incompletely halogenated hydrocarbons, NGu, carbon disulfide, etc (Ref 33). The effect of spontaneous decompn by oxidation-reduction reactions when N204 is mixed with a number of fuels (hydrazine, gasoline, liq paraffin, etc) has resulted in its extensive use in liq propint rocket engines (Refs 12, 22, 27 35)... 

We can use the electrochemical series to predict the thermodynamic tendency for a reaction to take place under standard conditions. A cell reaction that is spontaneous under standard conditions (that is, has K > 1) has AG° < 0 and therefore the corresponding cell has E° > 0. The standard emf is positive when ER° > Et that is, when the standard potential for the reduction half-reaction is more positive than that for the oxidation half-reaction. 

STRATEGY Find the standard potentials of the two reduction half-reactions in Appendix 2B. The couple with the more positive potential will act as an oxidizing agent (and be the site of reduction). That couple will be the right-hand electrode in the cell diagram corresponding to the spontaneous cell reaction. To calculate the standard emf of the cell, subtract the standard potential of the oxidation half-reaction (the one with the less-positive standard potential) from that of the reduction half-reaction. To write the cell reaction, follow the procedure in Toolbox 12.2. 

The last chapter in this introductory part covers the basic physical chemistry that is required for using the rest of the book. The main ideas of this chapter relate to basic thermodynamics and kinetics. The thermodynamic conditions determine whether a reaction will occur spontaneously, and if so whether the reaction releases energy and how much of the products are produced compared to the amount of reactants once the system reaches thermodynamic equilibrium. Kinetics, on the other hand, determine how fast a reaction occurs if it is thermodynamically favorable. In the natural environment, we have systems for which reactions would be thermodynamically favorable, but the kinetics are so slow that the system remains in a state of perpetual disequilibrium. A good example of one such system is our atmosphere, as is also covered later in Chapter 7. As part of the presentation of thermodynamics, a section on oxidation-reduction (redox) is included in this chapter. This is meant primarily as preparation for Chapter 16, but it is important to keep this material in mind for the rest of the book as well, since redox reactions are responsible for many of the elemental transitions in biogeochemical cycles. 

The spontaneous redox reaction shown in Figure 19-7 takes place at the surfaces of metal plates, where electrons are gained and lost by metal atoms and Ions. These metal plates are examples of electrodes. At an electrode, redox reactions transfer electrons between the aqueous phase and the external circuit. An oxidation half-reaction releases electrons to the external circuit at one electrode. A reduction half-reaction withdraws electrons from the external circuit at the other electrode. The electrode where oxidation occurs is the anode, and the electrode where reduction occurs is the cathode. 

In general, we expect that an external potential wiii drive the pair of reactions whose spontaneous reaction has the ieast positive ceii potentiai. In this case, combining the reduction of water to hydrogen and hydroxide with the oxidation of water to H3 O and oxygen gas gives the ieast positive E ° ... 

Examples of tunneling in physical phenomena occur in the spontaneous emission of an alpha particle by a nucleus, oxidation-reduction reactions, electrode reactions, and the umbrella inversion of the ammonia molecule. For these cases, the potential is not as simple as the one used here, but must be selected to approximate as closely as possible the actual potential. However, the basic qualitative results of the treatment here serve to explain the general concept of tunneling. 

The tarnish on silver, Ag2S, can be removed by boiling the silverware in slightly salty water (to improve the water s conductivity) in an aluminum pan. The reaction is an oxidation-reduction reaction that occurs spontaneously, similar to the redox reaction occurring in a voltaic cell. The Ag in Ag2S is reduced back to silver, while the A1 in the pan is oxidized to Al3+. 

One easily demonstrated electrical characteristic of moist soil is seen in the production of electricity when two different metals, namely, copper and zinc, are inserted into it. This is not unexpected because any salt-containing solution adsorbed in media, such as paper or cloth, and placed between these same two electrodes will cause a spontaneous reaction that produces electricity. The source of this flow of electrons is an oxidation-reduction reaction, represented as two half-reactions as shown in Figure 9.1 for copper and zinc. 

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. 

Tabulated E values can be used to calculate the for any reaction, as illustrated in Table 7.2 for the Zn/Cu galvanic cell. The redox reaction is spontaneous when the half-reaction (Cu /Cu) with the larger reduction (+0.34V) acts as the oxidizing agent. In this case, the other half-reaction (Zn /Zn) proceeds as an oxidation. The halfcell potential for this reduction is +0.76 V as it represents the reverse of the half-cell reduction potential as listed in Table 7.2. The sum of the oxidation and reduction half reactions is +0.34V + 0.76 V = +1.10 V. Thus for the galvanic Zn/Cu cell is +1.10V. 

The thermodynamic criterion for spontaneity (feasibility) of a chemical and electrochemical reaction is that the change in free energy, AG have a negative value. Free-energy change in an oxidation-reduction reaction can be calculated from knowledge of the cell voltage ... 

Electrons created in the oxidation reaction at the anode of a voltaic cell flow along an external circuit to the cathode, where they fuel the reduction reaction taking place there. We use the spontaneous reaction between zinc and copper as an example of a voltaic cell here, but you should realize that many powerful redox reactions power many types of batteries, so they re not limited to reactions between copper and zinc. 

The energy made available by this spontaneous electron flow (the free-energy change for the oxidation-reduction reaction) is proportional to AE ... 

Identify the spontaneous reactions among the following reactions and, for the spontaneous reactions, write balanced reduction and oxidation half-reactions. Show that the reaction is spontaneous by calculating the standard free energy of the reaction. 

The E° value for the overall reaction of Zn(s) with Ag+ aq) is the sum of the E° values for the two half-reactions 0.80 V + 0.76 V = 1.56 V. Because E° is positive (and AG° is negative), oxidation of zinc by Ag+ is a spontaneous reaction under standard-state conditions. Just as Ag+ can oxidize Zn, it s evident from Table 18.1 that Ag+ can oxidize any reducing agent that lies below it in the table (Fe2+, H202,1-, and so forth). The sum of E° for the Ag+/Ag reduction (0.80 V) and -E° for any half-reaction that lies below the Ag + /Ag half-reaction always gives a positive E° for the overall reaction. 

While such a device has yet to be constructed, Debreczeny and co-workers have synthesized and studied a linear D-A, -A2 triad suitable for implementation in such a device.11641 In this system, compound 6, a 4-aminonaphthalene monoimide (AN I) electron donor is excited selectively with 400 nm laser pulses. Electron transfer from the excited state of ANI to Ai, naphthalene-1,8 4,5-diimide (NI), occurs across a 2,5-dimethylphenyl bridge with x = 420 ps and a quantum yield of 0.95. The dynamics of charge separation and recombination in these systems have been well characterized.11651 Spontaneous charge shift to A2, pyromellitimide (PI), is thermodynamically uphill and does not occur. The mechanism for switching makes use of the large absorption cross-section of the NI- anion radical at 480 nm, (e = 28,300). A second laser pulse at 480 nm can selectively excite this chromophore and provide the necessary energy to move the electron from NI- to PI. These systems do not rely on electrochemical oxidation-reduction reactions at an electrode. Thus, switching occurs on a subpicosecond time scale. 

Many oxidation-reduction reactions may be carried out in such a way as to generate electricity. These cells are known as voltaic (older term galvanic) cells. In principle, any spontaneous, oxidation-reduction reaction (aqueous) can be set up to generate electricity by the following requirements ... 

An oxidation-reduction reaction that is not spontaneous, for which the calculated cell potential is negative, may be induced by electrolysis. This reaction can be due to an external electrical potential to force electrons into the couple undergoing reduction and to extract electrons from the couple undergoing oxidation. The minimum external potential required for electrolysis is the value of the calculated cell potential for the reaction. 

In (7), the (T1+ T1) couple functions as the oxidizing agent in (2), the (T13+ T1+) couples functions as the reducing agent. The reaction would occur at standard concentrations if E° for the reducing couple were to be less than E° for the oxidizing couple. Since the reduction potential of 1.25 V is greater than —0.336 V, the reaction cannot occur as written. We can now conclude that Tl+ does not spontaneously decompose to T1 and Tl3+. Actually, the reverse reaction is the spontaneous reaction 2T1 + Tl3+ 3T1+. 

For each of these overall cell reactions, write the oxidation and reduction half-reactions, calculate the standard cell potential, and determine if the reaction is spontaneous or not. 

The cell reaction is written as if oxidation occurs spontaneously at the left electrode and reduction at the right electrode. 

These processes change the initial spontaneous reaction so, that at the hydrogen electrode where formerly oxidation occurred reduction takes place, while at the chloride electrode the reduction process is replaced by oxidation. The total reaction in the electrolyzer which proceeds in an opposite direction to that in the galvanic cell can be expressed by the equation ... 

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