Oxidation Number

• Oxidation number refers to the hypothetical charge that an atom would have if all bonds to it were fully ionic. This value is an important concept in redox chemistry.

• In pure elements, the oxidation number is always zero, regardless of whether the element is monatomic (single atom) or polyatomic (multiple atoms).

• For ions, the oxidation number equals the charge of the ion. For example, the oxidation number of Na+ is +1, and for S2-, it’s -2.

• Hydrogen generally has an oxidation number of +1 when bonded to a non-metal and -1 when bonded to metals.

• Oxygen usually has an oxidation number of -2. There are exceptions, such as in peroxides (e.g., H2O2) where it is -1, and in compounds with fluorine where it is positive.

• In a neutral compound, the sum of the oxidation numbers for each atom must equal zero. In ions, the sum of the oxidation numbers must equal the overall charge of the ion.

• Oxidation is an increase in oxidation number, signifying a loss of electrons. Reduction is a decrease in oxidation number, indicating a gain of electrons.

• The process of assigning oxidation numbers to atoms can be used to identify redox reactions.

• A substance which loses electrons (i.e., is oxidised) is called the reducing agent, while a substance which gains electrons (i.e., is reduced) is called the oxidising agent.

• Redox reactions can be balanced by adjusting the coefficients of the reactants and products, or by using the half-equation method. The half-equation method involves balancing the atoms and charges on either side of the equation.

• Oxidation numbers are valuable in determining the stoicheiometry of redox reactions; they can provide an overview of how electrons are transferred in the reaction.

• Understanding these concepts and skills is crucial for mastering redox chemistry, including succeeding in a chemistry examination. Regular practise and application will help solidify comprehension and application knowledge.