REDOX
I) OXIDATION NUMBERS AND OXIDATION STATES
Oxidation state is a numerical value associated with atoms of each element in a compound or ion. An atom can be said to be in an oxidation state of +2 but with an oxidation number of +2.
Oxidation states are usually calculated as the number of electrons that atoms lose, gain or share when they form ionic or covalent compounds.
The oxidation state of uncombined elements (not in compounds) is always zero.
For a monoatomic ion, the oxidation state of the element is the charge on the ion.
In a chemical species (compound or ion), the most electronegative element is given a negative oxidation state. Other elements are given positive oxidation states.
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The oxidation state of hydrogen in compounds is +1 except in metal hydrides, when it is -1.
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The oxidation state of oxygen in compounds is -2 except in peroxides, when it is -1, or in OF2, when it is +2.
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The sum of all oxidation states in a neutral atom is zero. In an ion, the oxidation states total the charge on the ion.
Redox is the term used for the simultaneous processes of oxidation and reduction.
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Oxidation is defined as the loss of electrons from an element in a chemical species (increase in oxidation number).
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Reduction is defined as the gain of electrons from an element in a chemical species (decrease in oxidation number).
An element, compound or ion which gains electrons is called an oxidising agent, as it causes other species to be oxidised.
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Examples are: KMnO4 and K2Cr2O7
An element, compound or ion which loses electrons is called a reducing agent, as it causes other species to be reduced.
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Examples are: Zn, NaBH4, LiAlH4, NaH
Half equations are equations which show the oxidation and reduction of a reaction separately. They show the action of electrons with oxidation states.
Examples:
Ox. state of sulphur S in:
SO2 S + (-2 – 2) = 0 S = +4
H2SO4 S + (1 + 1) + (-2 – 2 – 2 – 2) = 0 S = +6
II) OXIDATION NUMBERS IN COMPOUND NAMES
The names of some compounds include the oxidation state of one of the elements involved.
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The oxidation state is shown in brackets as a Roman numeral.
An example would be iron (III) chloride, in which iron has an oxidation state of +3.
There is no need to state the oxidation numbers in all compounds.
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Sodium chloride NaCl always has sodium in an oxidation state of +1 and chloride always refers to chlorine in the -1 oxidation state.
However there is a compound called sodium chlorate, where chlorine can exist in four oxidation states.
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Sodium (I) chlorate has chlorine in the +1 oxidation state: Na+ ClO-
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Sodium (III) chlorate has chlorine in the +3 oxidation state: Na+ ClO2-
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Sodium (V) chlorate has chlorine in the +5 oxidation state: Na+ ClO3-
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Sodium (VII) chlorate has chlorine in the +7 oxidation state: Na+ ClO4-
III) METALS AND NON-METALS
Metals and non-metals often react together to form and ionic compound.
In such a reaction the metals become positively charged ions by losing electrons and each non-metal atom becomes a negatively charged ion by gaining these electrons.
Mg + Cl2 à Mg2+ + 2Cl-
Each magnesium atom loses 2 electrons, so its oxidation number increases to +2. Magnesium is oxidised in this reaction.
Each chlorine atom gains an electron. Its oxidation number decreases to -1. Chlorine is reduced in this reaction.
Metals reacting with acids are examples of redox reactions. Take magnesium and hydrochloric acid for example:
Magnesium is oxidised, as initially its oxidation number was zero, but during the reaction it increased to +2. The hydrogen ions in the acid are reduced as they become hydrogen atoms.
The chloride ions are neither oxidised nor reduced in this reaction.
Other metals react with dilute acids in this way too – only metals less reactive than hydrogen (e.g. copper, silver and gold) will fail to do so.
GLOSSARY
Oxidation number (oxidation state): a number (with a positive or negative sign) assigned to the atoms of each element in an ion or compound. Oxidation states are determined using a set of rules devised by chemists.
Oxidation: the loss of electrons.
Redox: reactions which involve reduction and oxidation processes.
Reduction: the gain of electrons
