Writing and balancing redox equations

In order to write a redox equation, it is first necessary to write the two half-reactions that identify the oxidation and reduction processes taking place. The overall redox equation is then obtained by adding these two half-reactions together, so that the electrons in each half reaction cancel out. 

To write the half reactions, follow these simple rules  

Identify the atoms that are oxidized and reduced, using the oxidation number method. 

Most of the reactions that you will come across at this stage, will occur in neutral or acid solution, and step 2 applies. 

Note that the rules are slightly different if the reaction occurs in basic solution hydrogen atoms are balanced using H2O molecules and then the same number of OH ions are added to the opposite side of the equation to balance the oxygens. Carry on as before, adding electrons to balance the charges. 

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Strategy Follow the general procedure for writing and balancing redox equations in Chapter 4 and reviewed in Chapter 18. Note that because nitric acid is strong, it should be represented as H+ and N03 ions. 

For each of these pairs of half-reactions, write the balanced equation for the overall cell reaction and calculate the standard cell potential. Express the reaction using cell notation. You may wish to refer to Chapter 20 to review writing and balancing redox equations. 

Write a balanced redox equation for the reaction of mercury with aqua regia, assuming the products include HgCl/ and N02(g). 

J 3 Write and balance chemical equations for simple redox reactions (Self-Test K.4). 

An effective way to balance a redox equation is to break down the reaction into separate half-reactions, write and balance half-equations for these half-reactions, and recombine the balanced half-equations into an overall balanced equation (Table 5.5). A slight variation of this method is used for a reaction that occurs in a basic aqueous solution (Table 5.6). A redox reaction in which the same substance is both oxidized and reduced is called a disproportionation reaction. 

In earlier sections of this chapter, we showed how to write and balance equations for precipitation reactions (Section 4.2) and acid-base reactions (Section 4.3). In this section we will concentrate on balancing redox equations, given the identity of reactants and products. To do that, it is convenient to introduce a new concept, oxidation number. 

The key to writing and balancing equations for redox reactions is to think of the reduction and oxidation processes individually. We saw in Section K that oxidation is the loss of electrons and reduction the gain of electrons. 

Write the balanced chemical equation for (a) the thermal decomposition of potassium chlorate without a catalyst (b) the reaction of bromine with water (c) the reaction between sodium chloride and concentrated sulfuric acid, (d) Identify each reaction as a Bronsted acid—base, Lewis acid—base, or redox reaction. 

Predict what will happen when the following pairs of substances are allowed to react. Write a balanced chemical equation for each reaction. When the reaction involves ions, write a net ionic equation. Identify each reaction as precipitation, as acid-base, or as redox, (a) AgN03(a q) and NaCl(a q) (b)... 

We conventionally cite the oxidized form first within each symbol, which is why the general form is o,r> so pb4+ Pb + is correct, but 2+ 4+ is not. Some people experience difficulty in deciding which redox state is oxidized and which is the reduced. A simple way to differentiate between them is to write the balanced redox reaction as a reduction. For example, consider the oxidation reaction in Equation (7.1). On rewriting this as a reduction, i.e. Al3+(aq) + 3e = A Em, the oxidized redox form will automatically precede the reduced form as we read the equation from left to right, i.e. are written in the correct order. For example, o,r for the couple in Equation (7.1) is Ai3+,ai-... 

Appiying Concepts The processes used to remove metals from their ores usually involve a reduction process. For example, mercury may be obtained by roasting the ore mercury(II) sulfide with calcium oxide to produce mercury metal, calcium sulfide, and calcium sulfate. Write and balance the redox equation for this process. 

Diluted hydrochloric acid can be used to remove limestone (calcium carbonate) surrounding phosphate and silicate fossils. The reaction produces carbon dioxide, water, and aqueous calcium chloride. Write the balanced chemical equation. Is it a redox reaction Explain. 

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. 

Problem Classify each of the following redox reactions as a combination, decomposition, or displacement reaction, write a balanced molecular equation for each, as well as total and net ionic equations for part (c), and identify the oxidizing and reducing agents ... 

Write and balance the net ionic equation for the redox reaction between KMn04 and H2C2O4. 

For each of the five experiments described in Model 3, write the balanced chemical equation (no "e " appears in the balanced chemical equation) for the redox reaction that could occur between the metal bar and the ion in solution. Note that the same number of electrons must be lost and gained in the transfer process. In each case indicate the oxidizing agent and the reducing agent. 

On commencing study of this topic, students should be comfortable with writing and balancing equations for chemical reactions (covered in greater depth in section 3.2 of Chapter 3). Knowledge of ions is also central. Concepts of basic electricity, such as electric current and potential difierence (a concept many students struggle with), are essential. Redox reactions (covered in Chapter 7) enter the study of electrochemistry... 

Fe is reduced (its oxidation state goes down from +3 to +2), and V is oxidized (its oxidation state goes up from +2 to +3). Appendix D discusses oxidation numbers and how to balance redox equations. You should be able to write a balanced redox reaction as the sum of two half-reactions, one an oxidation and the other a reduction. 

In this section, each equation identifies an oxidizer and a reducer, as well as the oxidized and reduced products of the redox reaction. Write separate oxidation and reduction half-reaction equations, assuming that the reaction takes place in an acidic solution, and add them to produce a balanced redox equation. 

Oxides. Chlorine oxides in which Cl has an even oxidation number undergo self-redox. In basic solution, the products are oxyanions such as C102, CIOs", and CIO4". Write the balanced ionic equations for the self-redox of (a) CIO2, (b) CI2O3. 

Write the balanced half-reactions and a balanced redox equation for each of the following reactions in acidic solution (15.3)... 

When inspecting a cell diagram, we first need to identify the species involved in oxidation and in reduction. Then we can write balanced half-cell equations. Finally, we can combine the half-cell equations to give the overall cell reaction. The new part to this example is the Ce /Ce couple in the presence of the inert platinum electrode at which the reduction takes place. Again, when balancing redox equations, it is... 

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