Balancing Ionic Redox Equations

Balance a Net Ionic Redox Equation Balance the following redox equation. CI04 (aq) + Br (aq) —> Ch(aq) + Br2(g) (in acid solution)... 

C04-0100. Some of the following react when added together, but others do not. For those that do, write a balanced net ionic redox equation, (a) Cu + HCl(a g) (b) Cu + MgCl2(f2 g) (c)... 

Balance net Ionic redox equations by the oxidation-number method. 

Use the oxidation-number method to balance this net ionic redox equation for the reaction between the perchlorate ion and the bromide ion in acid solution. 

Use the oxidation-number method to balance the following net ionic redox equations. 

Write a balanced ionic redox equation using the following pairs of redox half-reactions. 

Balance these ionic redox equations by any method. 

The main difference between balancing ionic redox equations and molecular redox equations is in how we handle the ions. In the ionic redox equations, besides having the same number of atoms of each element on both sides of the final equation, we must also have equal net charges. In assigning oxidation numbers, we must therefore remember to consider the ionic charge. 

Several methods are used to balance ionic redox equations, including, with slight modification, the oxidation-number method just shown for molecular equations. But the most popular method is probably the ion-electron method. 

The ion-electron method uses ionic charges and electrons to balance ionic redox equations. Oxidation numbers are not formally used, but it is necessary to determine what is being oxidized and what is being reduced. 

Balance these ionic redox equations using the ion-electron method. These reactions occur in acidic solution. 

Balancing net ionic redox equations The simplest way to express a redox reaction is an equation that shows only the oxidation and reduction processes. In order to understand how this is done, consider the balanced equation for the reaction of iron(II) nitrate and... 

It is easier to balance redox equations in net ionic form than in overall form. 

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. 

Our goal in this chapter is to help you understand how to balance redox equations, know the different types of electrochemical cells, and how to solve electrolysis problems. Have your textbook handy—you may need to find some information in electrochemical tables. We will be using the mole concept, so if you need some review refer to Chapter 3, especially the mass/mole relationships. You might also need to review the section concerning net-ionic equations in Chapter 4. And don t forget to Practice, Practice, Practice. 

You could balance the chemical equation for the reaction of magnesium with aluminum nitrate by inspection, instead of writing half-reactions. However, many redox equations are difficult to balance by the inspection method. In general, you can balance the net ionic equation for a redox reaction by a process known as the half-reaction method. The preceding example of the reaction of magnesium with aluminum nitrate illustrates this method. Specific steps for following the half-reaction method are given below. 

The key to the oxidation-number method of balancing redox equations is to realize that the net change in the total of all oxidation numbers must be zero. That is, any increase in oxidation number for the oxidized atoms must be matched by a corresponding decrease in oxidation number for the reduced atoms. Take the reaction of potassium permanganate (KMn04) with sodium bromide in aqueous acid, for example. An aqueous acidic solution of the purple permanganate anion (Mn04 ) is reduced by Br- to yield the nearly colorless Mn2+ ion, while Br- is oxidized to Br2. The unbalanced net ionic equation for the process is... 

HCl + 2 K2Cr04 + 3 H2C2O4 — 6 CO2 + 2 CrClj + 4 KCl + 8 H2O It is often easier to balance redox equations in net ionic form than in overall form. 

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