Galvanic Cells

In order to keep track of the electron flow in redox reactions we will define the concept of oxidation state and provide some rules for determining the oxidation states of atoms within a compound. The oxidation state of an atom in a covalent compound is an imaginary charge assigned to that atom if all the electrons in its bonds were completely given to the more electronegative atom in the bond. Of course, if two atoms of the same element are bonded together, then the two electrons in its bond are shared equally. In ionic compounds, the charge on the ion is equal to its oxidation number. shows, using hydrogen peroxide as an example, how one calculates the oxidation state of an atom in a covalent compound.

The above procedure for assigning oxidation states leads to the following useful observations about oxidation states which you should verify given the above discussion:

1. Atoms in elemental form have oxidation states of zero.
2. The charge on a monoatomic ion is equivalent to its charge.
3. Hydrogen has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals.
4. F, because it only forms one bond and is the most electronegative element, has an oxidation state of -1.
5. O, unless bonded to F or itself, has an oxidation state of -2.
6. The sum of all oxidation states in a compound must equal the total charge on the species. For hydrogen peroxide, we have seen that both O's have -1 oxidation states and both H's have +1 oxidation states, the sum of which is zero--the charge on hydrogen peroxide.

Oxidation-reduction reactions, also called redox reactions, involve the transfer of electrons from one species to another. That electron transfer causes a change in oxidation state for both reactive partners. The reducing agent is oxidized, meaning that its oxidation number increases due to the loss of one or more electrons. The oxidizing agent is reduced, meaning that its oxidation number has decreased due to the gain of one or more electrons. For example, the between permanganate and iron metal in acidic aqueous solution involves the transfer of five electrons to each permanganate (the oxidizing agent) from iron (the reducing agent).

Balancing redox equations by inspection is quite difficult, as you must take into account not only the mass balance but also the charge balance in the equation. To aid in this task a set of rules, called the Half-Reaction Method, has been devised. The following rules work for reactions performed in acidic or in basic solution.

1. Separate oxidation and reduction half-reactions:

2. Balance all atoms except for hydrogen and oxygen in each half-reaction. In this example they are already balanced.

3. Balance oxygen by adding H₂O as needed:

4. To balance hydrogen, add H⁺ as needed: (Note: You still do this if you are in basic solution, later on, you will add OH^- to "neutralize" the acid.)

5. Balance the charge of each reaction by adding electrons to side with the greater charge:

6. Multiply each half-reaction by the least integer factor that equalizes the number of electrons in each half-reaction. Then, add the half-reactions to obtain the overall balanced reaction in acidic solution:

If your redox reaction is in acidic solution, the above reaction is properly balanced. However, if the reaction you wish to balance is in basic solution, you need to add these three steps:

7. If the redox reaction is one in basic solution, then add OH^- to both sides of the equation to "neutralize" each H⁺:

8. "React" H⁺ and OH^- to form H₂O and eliminate water molecules on both sides of the equation:

9. Make sure that all atoms and charges are indeed balanced in your overall balanced equation for the redox reaction in basic solution: