Reduction half-reaction, example

This half-reaction, too, is conceptual the electrons are not actually free. In the equation for a reduction half-reaction, the electrons gained always appear on the left of the arrow. In this example, the redox couple is Agf/Ag. 

Table 16-4 Examples of redox reactions, consisting of an oxidation and reduction half reaction...

A starting material loses electrons in an oxidation, so electrons appear among the products of the oxidation half-reaction. A starting material gains electrons in a reduction, so electrons appear among the reactants of the reduction half-reaction. The reaction of magnesium metal with hydronium ions to produce hydrogen gas provides an example Mg(.y) -I- 2H3 0 ((2 q) q) H2(g) + 2H2 0(/) Here are the half-reactions for this... 

When reactions do not divide clearly into an oxidation and a reduction, changes in oxidation numbers reveal how to divide the starting materials into half-reactions. Example illustrates the separation process. 

After oxidation and reduction half-reactions are balanced, they can be combined to give the balanced chemical equation for the overall redox process. Although electrons are reactants in reduction half-reactions and products in oxidation half-reactions, they must cancel in the overall redox equation. To accomplish this, multiply each half-reaction by an appropriate integer that makes the number of electrons in the reduction half-reaction equal to the number of electrons in the oxidation half-reaction. The entire half-reaction must be multiplied by the integer to maintain charge balance. Example illustrates this procedure. 

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 the example above, the electron transfer was direct, that is, the electrons were exchanged directly from the zinc metal to the cupric ions. But such a direct electron transfer doesn t allow for any useful work to be done by the electrons. Therefore, in order to use these electrons, indirect electron transfer must be done. The two half-reactions are physically separated and connected by a wire. The electrons that are lost in the oxidation half-reaction are allowed to flow through the wire to get to the reduction half-reaction. While those electrons... 

Recall that, if you consider a redox reaction as two half-reactions, electrons are lost in the oxidation half-reaction, and electrons are gained in the reduction half-reaction. For example, you know the reaction of zinc with aqueous copper(II) sulfate. 

Define a half-reaction. Give an example of an oxidation half-reaction and a reduction half-reaction. 

For example, in Figure 1.1 (as well as Figures 1.2 and 1.3), the voltage of the horizontal line comes from the following reduction half-reaction at all pH values ... 

Note that not all the reactions in Table 19-1 show the reduction of solid metals, as in our examples so far. We ve thrown in liquids and gases as well. Not every voltaic cell is fueled by a reaction taking place between the metals of the electrodes. Although the cathode itself must be made of a metal to allow for the flow of electrons, those electrons can be passed into a gas or a liquid to complete the reduction half-reaction. Examine Figure 19-2 for an example of such a cell, which includes a gaseous electrode. 

A large positive value of F implies that the oxidized form of the couple is a good oxidizing agent. For example, the reduction potential for the reduction of dichlorine to aqueous chloride ion is + 1.36 V, and the reduction half-reaction ... 

Identify the ion being converted to metal in each of the examples in the Electrolysis activity in eChapter 18.11. Write the reduction half reaction for the conversion of each ion to metal. Which do you expect to deposit the greatest number of moles per second Why Does the activity confirm your conclusion What determines the number of moles of metal deposited ... 

An organic molecule can be used as the sacrificial donor in a reduction half reaction. Generally there is no net energy storage, but (depending on the reaction) hydrogen may be evolved at the same time as a surplus reaction product. For example, the reaction... 

It follows that log Kr for any reduction half-reaction may be considered equally as log K for a complete redox reaction in which electrons are transferred from H2(g) to the reduced species in the reduction reaction. In the case of Eq. 2.18, for example, log KR for the reduction half-reaction is the same as log K for the redox reaction ... 

The firm thermodynamic status of log KR for reduction half-reactions permits the use of these parameters in the normal way (see Section 1.2 and Special Topic 1) to evaluate equilibrium activities of oxidized and reduced species and to compare the stabilities of reactants and products in redox reactions. As an example of a stability comparison, consider the possible reduction of N(V) to N(0) through the oxidation of C(0) to C(IV) in a soil solution.13 The reduction half-reaction for denitrification is implicit in Eq. 2.20 that for C oxidation is... 

You have probably worked with tables of standard reduction potentials before. These tables provide the reduction potentials of various substances. It describes an oxidized species s ability to gain electrons in a reduction half-reaction (like copper in the voltaic cell example). According to this definition, we can use a value from the table to represent the E°red in the expression above, but how do you find the E°ox ... 

Step 7 The oxidation half-reaction for this example has two electrons involved, and the reduction half-reaction has 8 electrons involved. To make the numbers of electrons equal, we multiply the former by 4. 

The two reduction half-reactions in this example represent the halfcells of a voltaic cell. 

For the example presented above, one electron is gained in the reduction half-reaction (Fe3+ +e - Fe2+), but two are lost in the oxidation half-reaction (Sn2+ -> Sn4+ + 2e ). A charge balance is obtained by multiplying the reduction half-reaction by two to obtain the half-reactions ... 

The tendency of an oxidation or reduction half-reaction to proceed can tell us what products to expect in an electrolysis reaction. For example, the reduction potential of sodium is far less (more negative) than that of water, so when we electrolyze a solution of sodium chloride, for example, water is reduced and not sodium ion. Electrolysis of a concentrated solution of sodium chloride produces chlorine along with the hydrogen, whereas electrolysis of a dilute solution of sodium chloride produces oxygen and hydrogen. To get sodium metal, we need to electrolyze molten (melted) sodium chloride in the absence of water altogether. Do not forget that aqueous solutions contain water as well as any solutes. 

By combining many pairs of half-cells into voltaic cells, we can create a list of reduction half-reactions and arrange them in decreasing order of standard electrode potential (from most positive to most negative). Such a list, called an emf series or a table of standard electrode potentials, appears in Appendix D, with a few examples in Table 21.2 on the next page. 

From this example we can see how a redox reaction is composed of the reduction half-reaction and the oxidation half-reaction. We note that ferrous iron is oxidized (it loses electrons) and permanganate is reduced... 

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