Oxidation/Reduction Reactions

In these reactions we see a transfer of electrons from one atom or molecule to another. First let's look at some terms.

• Oxidation refers to taking electrons away from a substance. So to oxidize means to behave like oxygen normally does and "steal" electrons.
• Reduction refers to receiving electrons - it is the opposite of oxidation.
• Note that oxidation and reduction always go together. When oxygen oxidizes it is itself reduced (it gains electrons).
• Consider the example of burning natural gas:

CH₄ + 2 O₂ CO₂ + 2 H₂O

Notice that the methane is oxidized by the oxygen. We say that the carbon and hydrogen are both oxidized to give the new covalent products, water and carbon dioxide.

Examples:

• Consider the reaction of calcium metal with hydrochloric acid. This is similar to the reaction of sodium and water, and hydrogen gas is given off.
  • Write a net ionic equation for this reaction:

H⁺_(aq) + Cl^-_(aq) + Ca⁰_(s) H_{2 (g)} + Ca²⁺_(aq) + Cl^-_(aq)

Balancing: 2 H⁺ + Ca⁰_(s) H_{2 (g)} + Ca²⁺

• Notice that this net ionic equation involves electron transfer, so it is a redox equation.
• Notice that charge is also balanced in the final equation.
• A piece of clean copper is placed in a solution of silver nitrate. Assuming the copper goes to Cu(II) write a balanced equation. Notice that in this case you must balance the charge in order to get the final equation.

Ag⁺ + NO₃^- + Cu⁰ Ag⁰ + NO₃^- + Cu²⁺

Balancing: 2Ag⁺ + Cu⁰ 2Ag⁰ + Cu²⁺

• A piece of clean iron is placed in a solution of copper(II) sulfate. Assuming iron goes to Fe(III) write a balanced equation. Again the problem is to balance the charge.

Cu²⁺ + SO₄^2- + Fe⁰ Cu⁰ + Fe³⁺ + SO₄^2-

Balancing: 3 Cu²⁺ + 2 Fe⁰ 3 Cu⁰ + 2 Fe³⁺

As another example we can look at a key oxidation reaction in glycolysis, the central pathway of metabolism. Don't worry about these reactions - they will not be on an exam. They are presented for your interest. In this case we use a biological oxidizing agent, NAD+ to take electrons (in the form of a hydride ion, H:-) Frequently organisms need to operate this reaction under anaerobic (oxygen free) conditions. For example the maximum power (energy/sec) you can get from your muscles is anaerobically. For this to occur you need to regenerate the oxidizer (NAD+) without the presence of oxygen. For this to occur we use another reaction: The mechanism for this electron transfer is shown below:

Balancing Redox Equations

There are two common methods for balancing redox reactions: the oxidation number method and the half-reaction method. The half-reaction method works very well for ionic reactions, it is relatively easy to give partial credit, and it is the only method I will use in this class. If you know how to do the other method you are welcome to do so, but be careful to make sure you show your work or I won't be able to give partial credit!

The Half-Reaction Method

In the half-reaction method what we do is first break an equation into two parts and then balance the parts individually. Presented stepwise:

• Separate the reaction into two half-reactions.
• Balance each half-reaction separately:
  1. Balance atoms other than O & H by inspection.
  2. Balance O by adding H₂O to the opposite side.
  3. Balance H:
     • In acid solution add H⁺ as appropriate.
     • In base solution add H₂O to the side needing H, balance O by adding equal numbers of OH^- to the opposite side.
  4. Balance the charge by adding electrons (e^-) - add to same side as excess of positive charge, or opposite side if excess negative charge.
• Balance the charges of the two half-reactions by multiplying appropriately, then add to two equations together and cancel items appearing on both sides.

Example. Balance the following equation as it occurs in acid solution:

MnO₄^- + Cl^- Mn²⁺ + Cl₂

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© R A Paselk

Last modified 10 December 2009