Oxidation and Reduction

Introduction

Oxidation-reduction (redox) reactions are one of several types of reactions commonly seen in nature and in the lab.  In redox reactions, there is a transfer of electrons from one substance to another.

The electron donor, which loses electrons, is oxidized and the electron acceptor, which gains electrons, is reduced.  One of the primary tasks in understanding redox reactions is determining the direction of electron transfer.  Therefore, we need some method of “bookkeeping” electrons. 

 

Oxidation Numbers

Oxidation numbers allow us to determine if electrons have been lost or gained by a substance during a chemical reaction.  The rules for assigning oxidation numbers are given below:

 

Rules for Assigning Oxidation Numbers

 

1.  The oxidation number for any element in its uncombined state is 0.

            Explanation:  any element in its uncombined state has exactly the number of electrons it should have – therefore it has an oxidation number of 0

            Example:  Fe, Mg, H₂, F₂ all have oxidation numbers of 0.

 

2.  In simple ionic compounds, the oxidation number is the charge of the ion.

            Explanation:  when a chlorine atom becomes an ion, it gains one electron and becomes –1.  Thus it would have an oxidation number of –1.

            Example:  in NaCl, Na would have an oxidation number of +1, Cl would have an oxidation number of –1.

 

3.  In other compounds, many elements have characteristic oxidation numbers:

            Oxygen:  typically has an oxidation number of –2.  The major exception is in peroxides, where oxygen has an oxidation number of –1. 

            Hydrogen:  typicall has an oxidation number of –1.  The major exception is in hydrides, where hydrogen is combined with a metal.  In that case, hydrogen has a –1 oxidation number.

 

4.  The sum of the oxidation numbers must be equal to the charge on the ion or molecule.

            For neutral molecules, the oxidation numbers must sum to 0.

            For polyatomic ions, the oxidation numbers must sum to the charge on the ion.

Examples of Oxidation Numbers

Assign oxidation numbers to all atoms in the following substances:

 

            1.  S₈                Sulfur is in its uncombined state so the oxidation number is 0.

 

            2.  CaO            This is a simple ionic compound, so oxidation states equal charges

                                    Calcium is a +2 ion so the oxidation number is 2+

                                    Oxygen is a –2 ion so the oxidation number is 2-

 

            3.  BaCl₂          This is a simple ionic compound, so oxidation numbers equal charges

                                    Barium = 2+

                                    Chlorine = 1- (notice that oxidation numbers are calculated for each atom)

 

            4.  CO₂                        This is a molecular compound, not ionic.

                                    Typical oxidation number for oxygen is –2

                                    The oxidation numbers must sum to 0 because molecule is neutral.

                                    Oxidation number for carbon must be 4+.

 

            5.  PO₄^3-           This is a polyatomic ion.

                                    Typical oxidation number for oxygen is –2

                                    The oxidation numbers must sum to –3, the charge on the ion.

                                    x + 4(-2) = -3              x=5+, which is oxidation number for P

 

1,  Assign oxidation numbers to all atoms in the following substances:

a.       FeCl₃

b.       SO₂

c.       O₂

d.      Na₂S

e.       CO₃^2-

f.        H₂SO₄

 

Use of Oxidation Numbers to Predict Direction of Electron Transfer

Now that we can assign oxidation numbers, let’s use them to predict the direction of electron transfer in a chemical reaction.  Remember that electrons have a negative charge.  So, if a substance gains electrons during a reaction, the oxidation number will become more negative.  If a substance loses electrons during a reaction, its oxidation number will become more positive.

Let’s look at an example:

            2Na + Cl₂ à 2NaCl

 

First, we assign oxidation numbers.  

 

            2Na⁰ + Cl₂⁰ à 2Na⁺¹Cl^-1

 

Then, we use the oxidation numbers to determine how electrons have been transferred.

 

            Na goes from an oxidation number of 0 to 1+.  It has lost an electron.

            Each Cl goes from an oxidation number of 0 to 1-.  It has gained an electron.

           

We can use other terms to indicate electron transfer.

 

            If a substance has lost an electron, it has been oxidized.  This substance is the reducing agent.

            If a substance has gained an electron, it has been reduced.  This substance is the oxidizing agent

 

So, another way of indicating the transfer of electrons would be to say:

            Na has been oxidized and is the reducing agent. 

Cl has been reduced and is the oxidizing agent..

2.  Use oxidation numbers to indicate which substances have been oxidized and which substances have been reduced in the following reactions.

            a.  Fe + Cl₂ à FeCl₂

 

b.      2C₂H₆ + 7O₂ à 4CO₂ + 6H₂O

 

c.       NaI + 3HOCl à NaIO₃ + 3HCl