Oxidation-reduction ionic equations

A more general and fundamental view is obtained by a consideration of (a) the number of electrons involved in the partial ionic equation representing the reaction, and (b) the change in the oxidation number of a significant element in the oxidant or reductant. Both methods will be considered in some detail. 

We can now apply our knowledge of partial ionic equations to the subject of equivalents. The standard oxidation-reduction process is H H+ + e, where e represents an electron per atom, or the Avogadro number of electrons per mole. If we know the change in the number of electrons per ion in any oxidation-reduction reaction, the equivalent may be calculated. The equivalent of an oxidant or a reductant is the mole divided by the number of electrons which 1 mole of the substance gains or loses in the reaction, e.g. ... 

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. 

Determining the net ionic equation by balancing the oxidation-reduction reaction skeletal equation Cu(s) + HN03(ag) - Cu2+(aq) + NO(g)... 

Molarity Solubility rules Acids and bases Oxidation and reduction Net ionic equations Titrations... 

Try the following practice problems to review your understanding of net ionic equations, and to work with the new concepts of oxidation and reduction. 

You already know that some metals are more reactive than others. You may also have carried out an investigation on the metal activity series in a previous course. In Investigation 10-A, located on page 470, you will discover how this series is related to oxidation and reduction. You will write chemical equations, ionic equations, and half-reactions for the single displacement reactions of several metals. 

Write an oxidation half-reaction and a reduction half-reaction for each net ionic equation you wrote in question 1. Use the smallest possible whole-number coefficients in each half-reaction. 

Step 2 Divide the unbalanced net ionic equation into an oxidation half-reaction and a reduction half-reaction. To do this, you may need to assign oxidation numbers to all the elements in the net ionic equation to determine what is oxidized and what is reduced. 

Determine which of the following balanced net ionic equations represent redox reactions. For each redox reaction, identify the reactant that undergoes oxidation and the reactant that undergoes reduction. 

It should be evident that with a little practice you can very quickly, efficiently, and infallibly balance the most complicated electron-transfer equations. It is a straightforward mechanical process. This statement is true IF you know what the products of oxidation and reduction are. The most difficult situation that exists for balancing equations is the one characterized by the following request "Write a balanced ionic equation for the reaction, if any, that occurs when you mix A and B. You know the potential reactants because they are given, but that is all. 

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... 

An alternative to the oxidation-number method for balancing redox reactions is the half-reaction method. The key to this method is to realize that the overall reaction can be broken into two parts, or half-reactions. One half-reaction describes the oxidation part of the process, and the other half-reaction describes the reduction part. Each half is balanced separately, and the two halves are then added to obtain the final equation. Let s look at the reaction of aqueous potassium dichromate (K2Cr2C>7) with aqueous NaCl to see how the method works. The reaction occurs in acidic solution according to the unbalanced net ionic equation... 

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

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. 

In Chapter 5, we learned to write formulas for ionic compounds from the charges on the ions and to recognize the ions from the formulas of the compounds. For example, we know that aluminum chloride is AICI3 and that VCI2 contains ions. We cannot make comparable deductions for covalent compounds because they have no ions there are no charges to balance. To make similar predictions for species with covalent bonds, we need to use the concept of oxidation number, also called oxidation state. A system with some arbitrary rules allows us to predict formulas for covalent compounds from the positions of the elements in the periodic table and also to balance equations for complicated oxidation-reduction reactions. 

The expression molecular weightlionic charge is actually mass of the substance per unit of the reference species, where ionic charge is the reference species. The expression molecular weight n is also actually mass of the substance per unit of the reference species. In the case of the method based on the oxidation-reduction reaction, no equation was developed however, the ratios used in the example are ratios of the masses of the respective substances to the reference species, where 4, the number of electrons, is the number of reference species. 

For most oxidation-reduction systems z — zl is relatively high, e.g., 7 for the Fe(CN)o, Fe(CN)F system, and so the last terra in equation (7), which represents the activity coefficient factor, may be quite considerable further, the terms in the ionic strength involve the square of the valence and hence t will be large even for relatively dilute solutions. In any case, the presence of neutral salts, w hich were frequently added to the solution in the earlier studies of oxidation-reduction potentials, increases the ionic strength they will consequently have an appreciable influence on the potential, although the ratio of the amounts of oxidized to reduced forms remains constant. 

Write the oxidation and reduction half-reactions for the net ionic equation. 

Step 2. Write the oxidation and reduction half-reactions for the net ionic equation, including oxidation numbers. Recall the rules for assigning oxidation numbers from page 641. 

Most analytical oxidation/reduction reactions are carried out in solutions that have such high ionic strengths that activity coefficients cannot be obtained via the Debye-Hiickel equation (see Equation 10-1, Section lOB-2). Significant errors may result, however, if concentrations are used in the Nernst equation rather than activities. For example, the standard potential for the half-reaction... 

To write ionic equations, we must recognize compounds that are (1) soluble in water and (2) ionized or dissociated in aqueous solutions. To determine which are oxidation-reduction reactions, we should assign an oxidation number to each element. 

When balancing redox equations, we often find it convenient to omit the spectator ions (Section 4-3) so that we can focus on the oxidation and reduction processes. We use the methods presented in this chapter to balance the net ionic equation. If necessary we add the spectator ions and combine species to write the balanced formula unit equation. Examples 11-15 and 11-16 illustrate this approach. 

The oxidation half-reaction involves one electron, and the reduction half-reaction involves five electrons. Now we balance the electron transfer and then add the two equations term by term. This gives the balanced net ionic equation. 

For each of the following unbalanced equations, (i) write the half-reactions for oxidation and reduction, (ii) identify the species that lose and the species that gain electrons, and (iii) write the balanced net ionic equation for the overall reaction. 

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