Oxidation–Reduction Reactions for AP Chemistry (page 2)

By — McGraw-Hill Professional
Updated on Feb 1, 2011

Combination Reactions

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

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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