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Redox Main
Half-Reactions Balancing Redox Reactions in Neutral, Acidic, and Basic Solutions |
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Half-Reactions
All redox reactions involve electron transfer. To show this, we can use half-reactions, in which a species either gains or loses electrons. An example of zinc losing electrons is shown below:
Zn(s)
Zn2+(aq) + 2e-
Or, a species may gain electrons, as in the following half-reaction:
Cl2(g) + 2e-
Half-reactions illustrate some important terminology in redox reactions. First, a species that gains electrons is said to be reduced. Conversely, a substance that loses some of its electrons is oxidized. The individual reactants also have names: a reduced substance is called the oxidizing agent, because it allows another species to be oxidized. The oxidized reactant is, not surprisingly, called the reducing agent, because it allows another reactant to be reduced. So, in recap:
- A species that loses electrons is oxidized, and is called the reducing agent
- A species that gains electrons is reduced, and is called the oxidizing agent
These terms are applicable to all redox reactions, but half-reactions make them clearer. For example, in the following half-reaction,
Na(s)
solid sodium metal is oxidized, since it loses an electron. It is also the reducing agent, but this term is not really appropriate without some other reactant to reduce. In the following reaction, iron is reduced (gains electrons), and can be called an oxidizing agent, even though nothing is oxidized:
Fe3+(aq) + e-
Half-reactions are subject to the same balance requirements as all other reactions; in this type of reaction, the most important item is charge. Notice in the above half-reaction that the charge is +2 on both sides (since the electron has a -1 charge and the iron ion has a +3 charge, the net charge on the left side is +2). Using charge, balance the following half-reaction, and identify what is being reduced or oxidized:
Al(s)
Obviously, three electrons will be produced. In this case, aluminum is being oxidized to a +3 ion; it can also be called the reducing agent.
Adding two half-reactions will yield a complete reaction; for instance, the first two examples of the page add up to the following reaction:
Zn(s) + Cl2(g)
Adding half-reactions is integral to understanding the concepts of the next page. Balancing redox reactions is both different from and much trickier than good old stoichiometric balancing, since electron transfer must be accounted for. In addition, the effects of acids and bases must be accounted for as well.
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