Oxidation-reduction equations for

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 for this reaction ... 

Write an oxidation-reduction equation for the formation of an acid of each of the important oxidation states of sulfur. 

Write a balanced oxidation-reduction equation for the reaction of each of the metals below... 

Sulfuric acid, H2SO4, oxidizes many metallic elements. One of the effects of acid rain is that it produces sulfuric acid in the atmosphere, which then reacts with metals used in construction. Write balanced oxidation-reduction equations for the reaction of sulfuric acid with Fe, Zn, Mg, Co, and Ni. 

Thus, the overall equation for the formation of CuO involves an oxidation and a reduction that occur simultaneously. In every oxidation and reduction, the number of electrons lost must be equal to the number of electrons gained. Therefore, we multiply the oxidation reaction of Cu by 2. Canceling the 4 on each side, we obtain the overall oxidation-reduction equation for the formation of CuO. 

A solution of chlorine gas in water is used in chemical analysis to test for the presence of Br and I ions in solution. CI2 is able to oxidize easily these two ions to the elemental forms Br2 and I2, which may then be identified by their colors. Write a balanced oxidation-reduction equation for each of these processes. 

Write plausible half-equations and a balanced oxidation-reduction equation for the disproportionation of Xep4 to Xe and Xe03 in aqueous acidic solution. Xe and Xe03 are produced in a 2 1 mole ratio, and 02(g) is also produced. 

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. 

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... 

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

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. 

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.)... 

In using the oxidation states method to balance an oxidation-reduction equation, we find the coefficients for the reactants that will make the total increase in oxidation state balance the total decrease. The remainder of the equation is then balanced by inspection. 

The procedures for balancing an oxidation-reduction equation by the oxidation states method are summarized below. 

The general procedure is to balance the equations for the half-reactions separately and then to add them to obtain the overall balanced equation. The half-reaction method for balancing oxidation-reduction equations differs slightly depending on whether the reaction takes place in acidic or basic solution. 

We can split any oxidation/reduction equation into two half-reactions that show which species gains electrons and which loses them. For example. Equation 18-1 is... 

In writing oxidation-reduction equations, care should be taken to write formulas only for compounds or ions that have true chemical existence, like Mn02, H3ASO3, HAsO . Mn (-1-4), As (-1-3), and As ( + 5) are not true species and should not be written as such in equations. 

Chapter 4 Types of Chemical Reactions and Solution Stoichiometry," now gives a more qualitative and intuitive method for balancing oxidation-reduction equations, which is based... 

When this oxidation-reduction equation is balanced in acidic solution, using only whole number coefiTicients, what is the coefficient for S(s l... 

The Nemst equation, as described above, is only strictly valid for mineral oxidation-reduction systems for example ... 

In lesson C the student must know the definition of molarity and moles, the quantitative relationship of a chemical equation to determine quantities of reactants and products for reactions. Lesson D has the following objectives Assignment of oxidation numbers to elements according to a set of rules balancing oxidation-reduction equations by the half-reaction method and identification of oxidizing and reducing agents. 

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