Oxidation half-reactions

The electrochemical potential for the reaction is the difference between the reduction potentials for the reduction and oxidation half-reactions thus,... 

Balance the net charge in each half-reaction by adding electrons the electrons should be a reactant for the reduction half-reaction and a product for the oxidation half-reaction. 

All of the reactions considered in this chapter are of the oxidation-reduction type. You will recall from Chapter 4 that such a reaction can be split into two half-reactions. In one halfreaction, referred to as reduction, electrons are consumed in the other, called oxidation, electrons are produced. There can be no net change in the number of electrons the number of electrons consumed in reduction must be exactly equal to the number produced in the oxidation half-reaction. 

At the zinc anode (connected to the red wire shown in Figure 18.2), electrons are produced by the oxidation half-reaction... 

To obtain the standard voltage for an oxidation half-reaction, all you have to do is change the sign of the standard potential listed in Table 18.1. For example, knowing that... 

The oxidation number of cadmium increases from 0 (Cd) to +2 (Cd2+), so the oxidation half-reaction is... 

We consider oxidation first. To show the removal of electrons from a species that is being oxidized in a redox reaction, we write the chemical equation for an oxidation half-reaction. A half-reaction is the oxidation or reduction part of the reaction considered alone. For example, one battery that Volta built used silver and zinc plates to carry out the reaction... 

An oxidation half-reaction is a conceptual way of reporting an oxidation the electrons are never actually free. In an equation for an oxidation ha If-reaction, the electrons released always appear on the right of the arrow. Their state is not given, because they are in transit and do not have a definite physical state. The reduced and oxidized species in a half-reaction jointly form a redox couple. In this example, the redox couple consists of Zn2+ and Zn, and is denoted Zn2+/Zn. A redox couple has the form Ox/Red, where Ox is the oxidized form of the species and Red is the reduced form. 

Balancing the chemical equation for a redox reaction by inspection can be a real challenge, especially for one taking place in aqueous solution, when water may participate and we must include HzO and either H+ or OH. In such cases, it is easier to simplify the equation by separating it into its reduction and oxidation half-reactions, balance the half-reactions separately, and then add them together to obtain the balanced equation for the overall reaction. When adding the equations for half-reactions, we match the number of electrons released by oxidation with the number used in reduction, because electrons are neither created nor destroyed in chemical reactions. The procedure is outlined in Toolbox 12.1 and illustrated in Examples 12.1 and 12.2. 

Multiply the reduction half-reaction hy 2 and multiply the oxidation half-reaction by 5. 

The chemical equation for a reduction half-reaction is added to the equation for an oxidation half-reaction to form the balanced chemical equation for the overall redox reaction. 

The electrode at which oxidation takes place is called the anode. The electrode at which reduction takes place is called the cathode. Electrons are released by the oxidation half-reaction at the anode, travel through the external circuit, and reenter the cell at the cathode, where they are used in the reduction half-reaction. A commercial galvanic cell has its cathode marked with a + sign and its anode with a — sign. 

Step 2 Write the equation for the electrode on the left of the cell diagram as an oxidation half-reaction. 

The more negative the potential, the greater the electron-donating power of the oxidation half-reaction and therefore the more strongly reducing the redox couple (that is, the stronger the tendency for the half-reaction to occur as an oxidation). 

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. 

Identify the reactions with K > 1 among the following reactions and, for each such reaction, write balanced reduction and oxidation half-reactions. For those reactions, show that K >... 

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... 

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. 

The problem states that the nickel electrode is the cathode. Because reduction takes place at the cathode, we know that the nickel half-reaction occurs as reduction. In the oxidation half-reaction, cadmium is oxidized at the anode. 

They are the basis of many products and processes, from batteries to photosynthesis and respiration. You know redox reactions involve an oxidation half-reaction in which electrons are lost and a reduction half-reaction in which electrons are gained. In order to use the chemistry of redox reactions, we need to know about the tendency of the ions involved in the half-reactions to gain electrons. This tendency is called the reduction potential. Tables of standard reduction potentials exist that provide quantitative information on electron movement in redox half-reactions. In this lab, you will use reduction potentials combined with gravimetric analysis to determine oxidation numbers of the involved substances. 

In the ion-electron method of balancing redox equations, an equation for the oxidation half-reaction and one for the reduction half-reaction are written and balanced separately. Only when each of these is complete and balanced are the two combined into one complete equation for the reaction as a whole. It is worthwhile to balance the half-reactions separately since the two half-reactions can be carried out in separate vessels if they are suitably connected electrically. (See Chap. 14.) In general, net ionic equations are used in this process certainly some ions are required in each half-reaction. In the equations for the two half-reactions, electrons appear explicitly in the equation for the complete reaction—the combination of the two half-reactions—no electrons are included. 

Balance the change in oxidation number by adding electrons to the side with the higher total of oxidation numbers. That is, add electrons on the left for a reduction half-reaction and on the right for an oxidation half-reaction. One way to remember on which side to add the electrons is the following mnemonic ... 

It is helpful to think of the photosynthesis reaction as the sum of an oxidation half reaction and a reduction half reaction as shown in Figure 1. In fact, nature does separate these half reactions, in that the reduction of C02 to carbohydrates occurs in the stroma of the chloroplast, the organelle in the leaf where the photosynthesis reaction occurs, - whereas, the light-driven oxidation half reaction takes place on the thylakoid membranes which make up the grana stacks within the chloroplast. Reduced nicotinamide adenine dinucleotide phosphate (NADPH) carries the reducing power and most of the energy to the stroma to drive the fixation of C02 with the help of some additional energy provided... 

In such devices the light-absorbing semiconductor electrode immersed in an electrolyte solution comprises a photosensitive interface where thermodynamically uphill redox processes can be driven with optical energy. Depending on the nature of the photoelectrode, either a reduction or an oxidation half-reaction can be light-driven with the counterelectrode being the site of the accompanying half-reaction. N-type semiconductors are photoanodes, p-type semiconductors are photocathodes, and... 

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