Half-reaction reduction

A summary of oxidation-reduction half-reactions arranged in order of decreasing oxidation strength and useful for selecting reagent systems. 

Standard-state potentials are generally not tabulated for chemical reactions, but are calculated using the standard-state potentials for the oxidation, E°o, and reduction half-reactions, fi°red- By convention, standard-state potentials are only listed for reduction half-reactions, and E° for a reaction is calculated as... 

The two strongest oxidizing titrants are Mn04 and Ce +, for which the reduction half-reactions are... 

Note that the Pt cathode is an inert electrode that carries electrons to the reduction half-reaction. The electrode itself does not undergo oxidation or reduction. 

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. 

We begin by writing unbalanced equations for the oxidation and reduction half-reactions in part (a). 

Standard Reduction Potentials for Several Biological Reduction Half-Reactions ... 

Any redox reaction can be split into two half-reactions, an oxidation and a reduction. It is possible to associate standard voltages x (standard oxidation voltage) and (standard reduction voltage) with the oxidation and reduction half-reactions. The standard voltage for the overall reaction, °, is the sum of these two quantities... 

Standard half-cell voltages are ordinarily obtained from a list of standard potentials such as those in Table 18.1 (page 487). The potentials listed are the standard voltages for reduction half-reactions, that is,... 

Chromium metal can be electroplated from an aqueous solution of potassium dichromate. The reduction half-reaction is... 

As is always the case, a reduction half-reaction occurs at the cathode of an electrolytic cell. This half-reaction may be—... 

A voltaic cell consists of two half-cells. One of the half-cells contains a platinum electrode surrounded by chromium(III) and dichromate ions. The other half-cell contains a platinum electrode surrounded by bromate ions and liquid bromine. Assume that the cell reaction, which produces a positive voltage, involves both chromium(III) and bromate ions. The cell is at 25°C. Information for the bromate reduction half reaction is as follows ... 

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. 

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 1 Write the equation for the electrode on the right of the cell diagram as a reduction half-reaction (remember Right for Reduction). 

Standard potentials are also called standard electrode potentials. Because they are always written for reduction half-reactions, they are also sometimes called standard reduction potentials. 

The more positive the potential, the greater the electron-pulling power of the reduction half-reaction and, therefore, the more strongly oxidizing the redox couple (the stronger the tendency for the half-reaction to occur as a reduction). 

TABLE 12.1 Standard Potentials Species i at 25°C Reduction half-reaction E° (V)... 

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. 

One of the most useful applications of standard potentials is in the calculation of equilibrium constants from electrochemical data. The techniques that we develop here can be applied to any kind of reaction, including neutralization and precipitation reactions as well as redox reactions, provided that they can be expressed as the difference of two reduction half-reactions. 

Then find two reduction half-reactions in Appendix 2B that combine to give that equation. Reverse one of the half-reactions and add them together. 

Add this equation to the reduction half-reaction and cancel species that appear on both sides of the equation. 

One combination of electrodes that could be used to determine pH is a hydrogen electrode connected through a salt bridge to a calomel electrode. The reduction half-reaction for the calomel electrode is... 

The number of electrons required to reduce a species is related to the stoichiometric coefficients in the reduction half-reaction. The same is true of oxidation. Therefore, we can set up a stoichiometric relation between the reduced or oxidized species and the amount of electrons supplied. The amount of electrons required is calculated from the current and the length of time for which the current flows. 

Balance each of the following skeletal equations by using oxidation and reduction half-reactions. All the reactions take place in acidic solution. Identify the oxidizing agent and reducing agent in each reaction. 

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