Redox reaction balancing using halfreactions

Simple redox reactions can be balanced by the trial-and-error method described in Section 3.1, but other reactions are so complex that a more systematic approach is needed. There are two such systematic approaches often used for balancing redox reactions the oxidation-number method and the half-reaction method. Different people prefer different methods, so we ll discuss both. The oxidation-number method is useful because it makes you focus on the chemical changes involved the halfreaction method (discussed in the next section) is useful because it makes you focus on the transfer of electrons, a subject of particular interest when discussing batteries and other aspects of electrochemistry (Chapter 18). 

You will learn more about the importance of half-reactions when you study electrochemistry in Chapter 21. For now, however, you can learn to use halfreactions to balance a redox equation. First, look at an unbalanced equation taken from Table 20-3 to see how to separate a redox equation into half-reactions. For example, the following unbalanced equation represents the reaction that occurs when you put an iron nail into a solution of copper(II) sulfate, as shown in Figure 20-8. Iron atoms are oxidized as they lose electrons to the copper(ll) ions. 

As a result of this redox reaction between iron and copper sulfate solution, solid copper metal is deposited on the iron nail. To balance the equation given in the text for this reaction you could use the method of halfreactions. 

Balance equations for redox reactions in aqueous solution, using the halfreaction method, and calculate the concentrations of substances during redox titrations (Section 11.4, Problems 27-40). 

If a redox reaction occurs in basic solution, the equation must be balanced by using OH and H2O rather than and H2O. One approach is to first balance the halfreactions as if they occurred in acidic solution and then count the number of in each... 

When balancing redox equations, we consider the gain of electrons (reduction) separately from the loss of electrons (oxidation), express each of these processes as a halfreaction, and then balance both atoms and charge in each of the two half-reactions. When we combine the halfreactions, the number of electrons released in the oxidation must equal the number used in the reduction. 

Section 5.2 provides the thermodynamic basis for predicting whether or not a specific redox transformation can occur spontaneously in a given environment. The necessary redox halfreactions involving contaminants are usually well characterized because contaminants are the primary motivation for many studies of environmental systems. However, difficulties often arise in selecting the appropriate "environmental" half-reaction with which to balance the overall equation. When an environmental half-reaction cannot be identified, it is tempting to use traditional electrode potential measurements (127,128) as a generic measure of in situ redox conditions. These values (Emeas) then might be used as E°red or E°ox in Equation 14, to assess the thermodynamic potential of a particular contaminant transformation in a particular environment. However, a number of fundamental difficulties arise with this approach, so we do not recommend the procedure. 

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