Oxidation–Reduction Reactions for AP Chemistry (page 2)
Practice problems for these concepts can be found at:
- Reactions and Periodicity: Review Questions for AP Chemistry
- Reactions and Periodicity: Free-Response Questions for AP Chemistry
Oxidation–reduction reactions, commonly called redox reactions, are an extremely important category of reaction. Redox reactions include combustion, corrosion, respiration, photosynthesis, and the reactions involved in electrochemical cells (batteries). The driving force involved in redox reactions is the exchange of electrons from a more active species to a less active one. You can predict the relative activities from a table of activities or a halfreaction table. Chapter 16 goes into depth about electrochemistry and redox reactions.
The AP free-response booklet includes a table of half-reactions, which you may use for help during this part of the exam. A similar table can be found in the back of this book. Alternatively, you may wish to memorize the common oxidizing and reducing agents.
Redox is a term that stands for reduction and oxidation. Reduction is the gain of electrons and oxidation is the loss of electrons. For example, suppose a piece of zinc metal is placed in a solution containing the blue Cu2+ cation. Very quickly a reddish solid forms on the surface of the zinc metal. That substance is copper metal. As the copper metal is deposited, the blue color of the solution begins to fade. At the molecular level, the more active zinc metal is losing electrons to form the Zn2+ cation, and the Cu2+ ion is gaining electrons to form the less active copper metal. These two processes can be shown as:
The electrons that are being lost by the zinc metal are the same electrons that are being gained by the copper(II) ion. The zinc metal is being oxidized and the copper(II) ion is being reduced. Further discussions on why reactions such as these occur can be found in the section on single-displacement reactions later in this chapter.
Something must cause the oxidation (taking the electrons) and that substance is called the oxidizing agent (the reactant being reduced). In the example above, the oxidizing agent is the Cu2+ ion. The reactant undergoing oxidation is called the reducing agent because it is furnishing the electrons that are being used in the reduction half-reaction. Zinc metal is the reducing agent above. The two half-reactions, oxidation and reduction, can be added together to give you the overall redox reaction. When doing this, the electrons must cancel—that is, there must be the same number of electrons lost as electrons gained:
On the AP exam, you might be asked to identify what is being oxidized and reduced or to identify the oxidizing and reducing agents. (Be careful.)
In these redox reactions there is a simultaneous loss and gain of electrons. In the oxidation reaction (commonly called a half-reaction) electrons are being lost, but in the reduction half-reaction those very same electrons are being gained. So, in redox reactions electrons are being exchanged as reactants are being converted into products. This electron exchange may be direct, as when copper metal plates out on a piece of zinc, or it may be indirect, as in an electrochemical cell (battery).
Another way to determine what is being oxidized and what is being reduced is by looking at the change in oxidation numbers of the reactant species. (See the Basics chapter for a discussion of oxidation numbers and how to calculate them.) On the AP exam you may be asked to assign oxidation numbers and/or identify changes in terms of oxidation numbers. Oxidation is indicated by an increase in oxidation number. In the example above, the Zn metal went from an oxidation state of zero to +2. Reduction is indicated by a decrease in oxidation number. Cu2+ went from an oxidation state of +2 to zero. In order to figure out whether a particular reaction is a redox reaction, write the net ionic equation. Then determine the oxidation numbers of each element in the reaction. If one or more elements have changed oxidation number, it is a redox reaction.
There are several types of redox reaction that are given specific names. In the next few pages we will examine some of these types of redox reaction.
Combination reactions are reactions in which two or more reactants (elements or compounds) combine to form one product. Although these reactions may be of a number of different types, some types are definitely redox reactions. These include reactions of metals with nonmetals to form ionic compounds, and the reaction of nonmetals with other nonmetals to form covalent compounds.
In the first reaction, we have the combination of an active metal with an active nonmetal to form a stable ionic compound. The very active oxygen reacts with hydrogen to form the stable compound water. The hydrogen and potassium are undergoing oxidation, while the oxygen and chlorine are undergoing reduction.
Decomposition reactions are reactions in which a compound breaks down into two or more simpler substances. Although not all decomposition reactions are redox reactions, many are. For example, the thermal decomposition reactions, such as the common laboratory experiment of generating oxygen by heating potassium chlorate, are decomposition reactions:
In this reaction the chlorine is going from the less stable +5 oxidation state to the more stable –1 oxidation state. While this is occurring, oxygen is being oxidized from –2 to 0.
Another example is electrolysis, in which an electrical current is used to decompose a compound into its elements:
The spontaneous reaction would be the opposite one; therefore, we must supply energy (in the form of electricity) in order to force the nonspontaneous reaction to occur.
Single Displacement Reactions
Single displacement (replacement) reactions are reactions in which atoms of an element replace the atoms of another element in a compound. All of these single replacement reactions are redox reactions, since the element (in a zero oxidation state) becomes an ion. Most single displacement reactions can be categorized into one of three types of reaction:
- A metal displacing a metal ion from solution
- A metal displacing hydrogen gas (H2) from an acid or from water
- One halogen replacing another halogen in a compound
Remember: It is an element displacing another atom from a compound. The displaced atom appears as an element on the product side of the equation.
Reactions will always occur in the free-response section of the AP Chemistry exam. This may not be true in the multiple-choice part.
For the first two types, a table of metals relating their ease of oxidation to each other is useful in being able to predict what displaces what. Table 6.1 shows the activity series for metals, which lists the metal and its oxidation in order of decreasing ease of oxidation. An alternative to the activity series is a table of half-cell potentials, as discussed in Chapter 16. In general, the more active the metal, the lower its potential.
Elements on this activity series can displace ions of metals lower than themselves on the list. If, for example, one placed a piece of tin metal into a solution containing Cu(NO3)2(aq), the Sn would replace the Cu2+ cation:
The second equation is the net ionic form that is often required on the AP exam.
If a piece of copper metal was placed in a solution of Sn(NO3)2(aq) there would be no reaction, since copper is lower than tin on the activity series. This table allows us to also predict that if sodium metal is placed in water, it will displace hydrogen, forming hydrogen gas:
The Group IA and IIA elements on the activity table will displace hydrogen from water, but not the other metals shown. All the metals above hydrogen will react with acidic solutions to produce hydrogen gas:
Halogen reactivity decreases as one goes from top to bottom in the periodic table, because of the decreasing electronegativity. Therefore, a separate activity series for the halogens can be developed:
The above series indicates that if chlorine gas were dissolved in a KI(aq) solution, the elemental chlorine would displace the iodide ion:
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