Redox (2.1.5)

Oxidation numbers are a way of describing the location of electrons in a chemical species. They indicate the degree of oxidation (loss of electrons) or reduction (gain of electrons) of an atom in a compound compared to in its elemental form.

There are rules for assigning oxidation numbers.

1. The oxidation number of an atom in its elemental form is (e.g., ).
2. For monatomic ions, the oxidation number is equal to the charge of the ion (e.g., is , is ).
3. Oxygen in compounds usually has an oxidation number of (except in peroxides where it is ).
4. Hydrogen in compounds usually has an oxidation number of (except when bonded to metals in hydrides, where it is ).
5. The sum of the oxidation numbers in a neutral compound must be , and in a polyatomic ion, it must equal the overall charge of the ion.

Roman numerals (without a sign) are used to indicate the magnitude of the oxidation number of an element which commonly holds different oxidation numbers in different compounds.

Iron (II) represents where the oxidation number of iron is +2.

Iron (III) denotes where the oxidation number of iron is +3.

Nitrate (V) represents where the oxidation number of nitrogen is +5.

Nitrate (III) represents where the oxidation number of nitrogen is +3. The common name for is the nitrite ion.

The name nitrate written without an oxidation number is always assumed as .

The rules for assigning oxidation numbers can be used to construct the chemical formulae of a compound from its name.

There are many varieties of manganese oxide. The chemical formula of each compound can be deduced using the oxidation states given by roman numerals in the name.

Manganese (II) Oxide – The oxidation state of manganese is +2 and the oxidation state of oxygen is -2. A neutral compound must contain a 1:1 ratio of as the sum of the oxidation states must equal 0. The chemical formula is therefore .

Manganese (III) Oxide – The oxidation state of manganese is +3 and the oxidation state of oxygen is -2. A neutral compound must contain a 1:1.5 ratio of as the sum of the oxidation states must equal 0. This translates to a 2:3 whole number ratio. The chemical formula is therefore .

In an ionic compound the oxidation state of an ion is linked to its ionic charge.

Metals are generally less electronegative than non-metals; the metal has the positive oxidation number while the non-metal has the negative oxidation number.

• Group 1 metals always have the oxidation number +1.
• Group 2 metals always have the oxidation number +2.
• Transition metals can have variable oxidation states.
• Group 6 non-metals generally have the oxidation number -2.
• Group 7 non-metals always have the oxidation number of -1.

The sum of the oxidation numbers in an ionic compound equals zero.

Oxidation can be defined as the removal of electrons from an atom (taken by an oxidising agent).

Reduction can be defined as the gain of electrons by an atom (provided by a reducing agent).

Oxidation can be defined as the increase of oxidation number of an atom.

Reduction can be defined as the reduction of oxidation number of an atom.

Metals lose electrons to achieve a more stable electronic configuration, resembling the nearest noble gas configuration for main block metals.

This loss of electrons results in the formation of positively charged ions (cations). For example, sodium () loses one electron to form , and magnesium () loses two electrons to form .

When a metal atom loses electrons, its oxidation number increases. For example, sodium goes from an oxidation state of in to in .

Non-metals tend to gain electrons to achieve a stable electron configuration, often resembling the nearest noble gas configuration.

This gain of electrons results in negatively charged ions (anions). For example, chlorine () gains one electron to form , and oxygen () gains two electrons to form .

When a non-metal gains electrons, its oxidation number decreases. For example, chlorine goes from an oxidation state of in to in .

Whenever an acid reacts with a metal, the products are a salt and hydrogen.

This is a redox reaction.

The metal is oxidised whilst the acid is reduced.