Oxidation-reduction equations

If you know the reactants and products of a chemical reaction, you should be able to write an equation for the reaction and balance it. In writing the equation, first write down the correct formulas for all reactants and products. After they are written down, only then start to balance the equation. Do not balance the equation by changing the formulas of the substances involved. For simple equations, you should balance the equation by inspection. (Balancing oxidation-reduction equations will be presented in Chap. 13.) The following rules will help you to balance simple equations. 

One of the most important uses of oxidation numbers is in balancing redox (oxidation-reduction) equations. These equations can get very complicated, and a systematic method of balancing them is essential. There are many such methods, however, and each textbook seems to use its own. There are many similarities among the methods, however, and the following discussion will help no matter what method your instructor and your textbook use. 

The ratio by weight of potassium nitrate and sulfur corresponding to a balanced - or stoichiometric - mixture will be 4(101. 1) 404.4 grams (4 moles) of KNO 3 and 5(32.1) = 160.5 grams (5 moles) of sulfur. This equals 72% KNO 3 and 28% S by weight. An ability to balance oxidation-reduction equations can be quite useful in working out weight ratios for optimum pyrotechnic performance. 

The more diflicult oxidation-reduction equations cun often be written more easily by use of tile Stock system ol oxidation numbers, which are positive or negative valences or charges. Consider the reaction of potassium diehromale. K Cr 0-. with potassium sultile. KjSOi. in acid solution to lorni chromiuntlllll sulfate. Cr SOj i. and potassium sullale. K S04. The unbalanced expression for the ionic reaction ts... 

Balance the oxidation-reduction equation for the oxidation of H2S(aq) by HN03(aq) to produce NO(g) and S(s) in aqueous acidic solution (thus H+ and H20 may be involved). 

Balancing oxidation-reduction equations for reactions occurring in aqueous acidic solutions is usually fairly straightforward since we can use H20 to balance O, and then H+ to balance H. In basic solution,... 

Balance the oxidation-reduction equation H+NO + H2S - NO + S + H20. Half-Reaction Method... 

Balance the following oxidation-reduction equation FeS2 + 02 Fe203 + S02. 

Oxidation-reduction reactions Oxidation number Oxidizing and reducing agents Ionic notation for equations Balancing oxidation-reduction equations... 

Balance the oxidation-reduction equation for this reaction ... 

In order to balance oxidation-reduction equations, we must therefore find out how many electrons are released by the reducing agent and taken up by the oxidizing agent. This can easily be done if the half-cell reaction equations of the redox systems involved are known. In the above example, if we write up the two half-cell equations ... 

In general, balancing of oxidation-reduction equations should be made by taking the following steps ... 

It is very important that users of this book should acquire a routine knowledge of balancing oxidation-reduction equations. Further practice may be obtained by combining various oxidizing and reducing agents mentioned in Examples 25 to 30, for which the half-cell reactions can easily be checked from the numbered equations. A careful study of Section 1.38 which follows might also help. 

Net ionic equations are used in discussions of limiting quantities problems (Chapter 10), molarities of ions (Chapter 11), balancing oxidation-reduction equations (Chapter 16), acid-base theory (Chapter 19), and many other areas beyond the scope of this book. They make possible writing equations for halfreactions at the electrodes in electrochemical experiments (Chapter 17), which have electrons included explicitly in them. They make understandable the heat effects of many reactions such as those of strong acids with strong bases. 

Net ionic equations are used extensively in chemistry. For example, equilibrium expressions for acid-base reactions, as well as for the ionization of water itself, are conventionally written in the form of net ionic equations. Many complex oxidation-reduction equations are balanced using net ionic equations. These topics are introduced in Chapters 16 and 19. 

We have already balanced a number of simple oxidation-reduction equations, starting in Chapter 8. Most combination and decomposition reactions and all single substitution and combustion reactions are oxidation-reduction reactions. However, many oxidation-reduction reactions are much more complicated than the ones we have already considered, and we must use a systematic method for balancing equations for them. Unfortunately, many different systematic methods are used, and each chemistry instructor seems to have his or her own favorite method. Most instructors will accept any valid method that a student understands, however. The method outlined here is a standard method that should be acceptable. 

All methods of balancing oxidation-reduction equations are based on the overall gain of oxidation numbers in a reaction being the same as the overall loss of oxidation numbers in the reaction (because the same number of electrons must be gained as lost). 

The first step in any method of balancing oxidation-reduction equations is to identify the element that is oxidized and the one that is reduced. Because the change in oxidation number is equal to a change in the number of electrons controlled, and the electrons must be controlled by some atom, the total gain in oxidation number is equal to the total loss in oxidation number. The oxidation half of a reaction may be written in one equation, and the reduction half in another. Neither half-reaction can be carried out without the other, but they can be done in different locations if they are connected in such a way that a complete electrical circuit is made (Chapter 17). The half-reaction method is illustrated by balancing the equation for the reaction of zinc metal with dilute nitric acid to produce ammonium ion, zinc ion, and water ... 

Oxidation is defined as a gain in oxidation number, caused by a loss of electrons or of control of electrons. Reduction is defined as a loss in oxidation number, caused by a gain of electrons or of control of electrons. Complicated oxidation-reduction equations must be balanced according to some systematic method because they are too complex to be balanced by inspection. Although neither can take place alone, the oxidation and the reduction can occur in different locations if suitable electrical connections are provided. (Chapter 17) In the halfreaction method, the equation for the half-reaction involving oxidation and that for the half-reaction involving reduction are balanced separately then the two are combined. Each may be multiplied by a small integer if necessary to balance the numbers of electrons involved. 

Do not confuse oxidation numbers with charges when balancing oxidation-reduction equations. Use Roman numerals for pxrsitive oxidation numbers and Arabic numbers for charges. (To denote negative oxidation numbers, use Arabic numerals below the formula and circle them do not get them mixed up with charges. The Romans did not have negative numbers.)... 

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