Oxidation-Reduction Reactions

Oxidation is a chemical change in which electrons are lost by an atom or group of atoms, and reduction is a chemical change in which electrons are gained by an atom or group of atoms.

As a simple example of an oxidation-reduction consider what happens when you dip a strip of metallic zinc into a blue solution of copper(II) sulfate.  The strip of zinc becomes coated with a reddish-brown layer of metallic copper.  The molecular equation for this reaction is         

In this simple example we can say that zinc has been oxidized and copper has been reduced.  For convenience, we can think of this reaction as two separate parts or half reactions one involving the loss of two electrons by an atom of zinc and the other being the gain of two electrons by the copper(II) ion.  The two half reactions are

Two terms used in conection with oxidation-reduction reactions are oxidizing agent, a species that oxidizes another species and is itself reduced, and reducing agent, a species that reduces another species and is itself oxidized.  In our example above, the copper(II) ion is the oxidizing agent, and zinc metal is the reducing agent.  It is not always immediately apparent from the ionic charge alone whether a compound is undergoing oxidation or reduction.  For example, MnO₂ reacts with hydrochloric acid to produce, chlorine gas and the Mn²⁺ ion.  It is obvious that neutral chlorine is produced by oxidation of Cl^-.   We assume, therefore, that MnO₂ is undergoing reduction, in spite of the fact that a neutral substance, MnO₂ is converted to a cationic substance, Mn²⁺.  The concept of oxidation numbers was developed as a simple way of keeping track of electrons in a reaction.  By using oxidation numbers, you can determine whether electrons have been transferred and thereby whether an oxidation-reduction reaction has occurred.

Oxidation Numbers    

The oxidation number (sometimes called the oxidation state) of an atom in a substance is defined as the charge of the atom if it existed as a monoatomic ion, or a hypothetical charge assigned to the atom in the substance by a set of rules.  We then define an oxidation-reduction reaction as one in which one or more atoms change oxidation numbers, signaling that there was a transfer of electrons.

Rules for Assigning Oxidation Numbers         

Balancing Oxidation-Reduction Equations

Oxidation-reduction equations are often too dificult to balance by the inspection method(trial and error) we have used up to this point.  There are two systematic methods for balancing redox equations. One is the half-reaction or ion-electron method and the second is the oxidation-state method.  In this summary, I am only going to cover the half-reaction method.  First we will list the rules, and then we will go through a couple of examples in detail.

Rules for balancing oxidation-reductions by the half-reaction method

As our first example lets balance the following oxidation-reduction equation:

First we balance the half-reaction for the oxidizing agent as follows. Note that we have an acidic solution because of HCl.

Next we balance the half-reaction for the reducing agent.

Now we multiply each half-reaction by numbers chosen to make the electrons cancel when added.

The 14 H⁺ was added as 14 HCl, and 6 of the 14 chloride ions were oxidized.  To each side of the equation 8 more Cl^- are added to represent those Cl^- which were not oxidized.  Similarly, 2 K⁺ can be added to each side to show that Cr₂O₇^2- came from K₂Cr₂O₇. So our balanced molecular equation is

For our second example we will balance the following equation, which is in an alkaline solution.

The balancing of the half-reaction for the oxidizing agent goes as follows.

The balancing of the half-reaction for the reducing agent is as follows.

Now we multiply each half-reaction by numbers chosen to make the electrons cancel when added.

Canceling terms that are common to both sides we have.

Convert the net ionic equation to a molecular equation by adding ions.