Oxidation-reduction reactions redox spontaneous

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+. 

Redox reactions are more conveniently described in terms of relative electrical potentials instead of the equivalent changes in Gibbs free energy. The electrons in Equation 6.8 come from or go to some other redox couple, and whether or not the reaction proceeds in the forward direction depends on the relative electrical potentials of these two couples. Therefore, a specific electrical potential is assigned to a couple accepting or donating electrons, a value known as its oxidation-reduction or redox potential. This redox potential is compared with that of another couple to predict the direction for spontaneous electron flow when the two couples interact—electrons spontaneously move toward higher redox potentials. The redox potential of species /, ), is defined as... 

Electrochemistry is the branch of chemistry that deals with the interconversion of electrical energy and chemical energy. Electrochemical processes are redox (oxidation-reduction) reactions in which the energy released by a spontaneous reaction is converted to electricity or in which electrical energy is nsed to cause a nonspontaneous reaction to occur. Although redox reactions were discnssed in Chapter 4, it is helpful to review some of the basic concepts that will come np again in this chapter. 

Voltaic cells are based on spontaneous oxidation-reduction reactions. Conversely, it is possible to use electrical energy to cause nonspontaneous redox reactions to occur. For example, electricity can be used to decompose molten sodium chloride into its component elements ... 

Voltaic, or galvanic, cells are electrochemical cells in which spontaneous (product-favored) oxidation-reduction reactions produce electrical energy. The two halves of the redox reaction are separated, requiring electron transfer to occur through an external circuit. In... 

Some oxidation-reduction reactions do not occur spontaneously but can be driven by electrical energy. If electrical energy is required to produce a redox reaction and bring about a chemical change in an electrochemical cell, it is an electrolytic cell. Most commercial uses of redox reactions make use of electrolytic cells. 

Highly protective layers can also fonn in gaseous environments at ambient temperatures by a redox reaction similar to that in an aqueous electrolyte, i.e. by oxygen reduction combined with metal oxidation. The thickness of spontaneously fonned oxide films is typically in the range of 1-3 nm, i.e., of similar thickness to electrochemical passive films. Substantially thicker anodic films can be fonned on so-called valve metals (Ti, Ta, Zr,. ..), which allow the application of anodizing potentials (high electric fields) without dielectric breakdown. 

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. 

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. 

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. 

To predict whether a redox reaction is spontaneous, remember that an oxidizing agent can oxidize any reducing agent that lies below it in the table but can t oxidize one that lies above it. To calculate E° for a redox reaction, sum the E° values for the reduction and oxidation half-reactions. 

In the lemon battery shown in Figure 17.9, a different chemical reaction occurs at each of the metal-strip electrodes. The electrode made of the metal that is more easily oxidized becomes the anode—the electrode at which the oxidation reaction occurs. The second electrode becomes the cathode, and a reduction reaction proceeds at its surface. The substance in a lemon that is most easily reduced is the abundant hydrogen ion of the electrolyte. When these two reactions occur together, in the same cell, they combine to produce a spontaneous redox reaction. This type of reaction is represented by the equation below, where M is the metal that is oxidized. 

In a voltaic cell, a spontaneous redox reaction (AG < 0) is separated into an oxidation half-reaction (anode half-cell] and a reduction halfreaction (cathode half-celf). Electrons flow from anode to cathode through an external circuit, releasing electrical energy, and ions flow through a salt bridge to complete the circuit and balance the charge within the celi. 

Plan To determine whether a redox reaction is spontaneous under standard conditions, we first need to write its reduction and oxidation hah-reactions. We can then use the standard reduction potentials and... 

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