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Electronics

M - FROM MINERALS TO ELEMENTS

Introduction

Oxidation and Reduction

Group 7 Elements: The Halogens

Theory Of Solutions

Extracting Copper from its Ore

Acid-Base Theory

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Introduction to Oxidation and Reduction

Oxidation and Reduction in Terms of Electron Transfer

Using Half-equations to Write Full Redox Equations

Oxidation State

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Introduction to Oxidation and Reduction

In an elementary chemistry course, oxidation might be defined as a chemical reaction where oxygen is gained or hydrogen is lost.  Reduction is the opposite process to oxidation, a reaction where oxygen is lost or hydrogen is gained.

E.g.     CuO(s) + H₂(g)    Cu(s) + H₂O(g)

The copper (II) oxide loses oxygen and is reduced while the hydrogen gains oxygen and is oxidised.  The hydrogen is called the reducing agent (reductant) and is oxidised.  The copper (II) oxide is the oxidising agent (oxidant) and is reduced.

Oxidation and reduction occur together and the term redox reaction means a reaction where oxidation and reduction take place.  This elementary treatment is easy to understand but it does not include all possible reactions.

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Oxidation and Reduction in Terms of Electron Transfer

Oxidation is a process where electrons are lost.

Reduction is a process where electrons are gained.

Consider the reaction where magnesium burns in chlorine to produce magnesium chloride.

Mg(s) + Cl₂(g) MgCl₂(s)

There are two processes occurring:

Mg Mg²⁺ + 2e^-     and     Cl₂ + 2e^- 2Cl^-

Magnesium loses electrons and is oxidised; chlorine gains electrons and is reduced.  These two equations are called half-equations.

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Using Half-equations to Write Full Redox Equations

The concept of electron transfer is useful in balancing redox equations.

The following simple rules can be used to obtain a balanced equation:

1. Write down the oxidant and the reductant and determine their products.
2. Write separate half-equations for oxidation and reduction processes and balance these with respect to atoms and charge.
3. Combine the half-equations to obtain the overall equation.

Consider the reaction between permanganate ions and iodide ions in acid solution.

	Reduction	Oxidation
1.	MnO4- Mn2+ + 4H2O	2I- I2
2.	Balance the hydrogen atoms with H+.
	MnO4- + 8H+ Mn2+ + 4H2O	2I- I2
	Balance half-equations electrically by adding electrons.
	MnO4- + 8H+ + 5e- Mn2+ + 4H2O	2I- I2 + 2e-
3.	Multiply each half-equation by the number of electrons appearing in the other half-equation, and add the equations to eliminate the electrons.
4.	2MnO4- + 16H+ + 10e- 2Mn2+ + 8H2O	10I- 5I2 + 10e-
Adding:    2MnO4- + 16H+ + 10I- 2Mn2+ + 5I2 + 8H2O

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Oxidation State

The concept of oxidation state (or number) has been devised to give a guide to the extent of oxidation or reduction in a species.  The concept is without direct chemical foundation, but is very useful, being appropriate to both ionic and covalently bonded species.

Oxidation states are assigned to atoms in molecules or ions to shew how much they are oxidised or reduced according to some simple rules:

1. Atoms in elements are in oxidation state zero.
2. In simple ions, the oxidation state is the same as the charge on the ion.
   Magnesium has oxidation state 0 in Mg, but oxidation state +2 in Mg²⁺.
   Chlorine has oxidation state 0 in Cl₂, but oxidation state -1 in Cl^-.
3. The oxidation states of all the constituent elements in a compound add up to zero.
4. The oxidation states of all the constituent elements in an ion add up to the charge on the ion.
5. In compounds and ions, oxidation states are assigned as follows:
   
   

   F	-1
   O	-2 (except in peroxides, -1 and combined with fluorine)
   H	+1 (except in metal hydrides, -1)
   Cl	-1 (except when combined with O and F)

6. Some elements have different oxidation states.

An element is oxidised when its oxidation state increases.

An element is reduced when its oxidation state decreases.

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