Topic 3 - Redox I

• Oxidation involves the loss of electrons. Reduction involves the gain of electrons. This redox rule can be remembered using OILRIG.
• The oxidation number gives the oxidation state of an element or ionic substance.
• The oxidation number of an element is zero.
• Oxidation numbers in a neutral compound add up to zero.
• Oxidation numbers in a charged compound add up to total the charge.
• Hydrogen has an oxidation number of 1+.
• Oxygen has an oxidation number of 2-.
• Halogens have an oxidation number of 1-.
• Group I metals have an oxidation number of 1+.
• Group II metals have an oxidation number of 2+.
• Oxygen has an oxidation number of 1- in peroxides.
• Hydrogen has an oxidation number of 1- in metal hydrides.
• Roman numerals can be used to give the oxidation number of an element that has a variable oxidation state, depending on the compound it is in.
• In copper (II) sulfate, copper has an oxidation number of 2+.
• Elements are arranged in the periodic table by proton number and electron orbitals. Each orbital corresponds to a block on the periodic table, and each element in the block has their outer shell electrons in that orbital.
• Elements in the same block react in similar ways since their outermost electron is in the same type of orbital.
• S block elements (groups 1 and 2 metals) generally lose electrons, so are oxidised and form species with positive oxidation numbers.
• P block non-metals generally gain electrons, so are reduced and form species with negative oxidation states.
• An oxidising agent accepts electrons from the species that is being oxidised. Therefore, it gains electrons and is reduced. This is seen as an increase in oxidation number (gets more positive).
• A reducing agent donates electrons to the species being reduced. Therefore, it loses electrons and is oxidised. This is seen as a reduction in oxidation number (gets more negative).
• Redox reactions are reactions in which oxidation and reduction occur simultaneously take place when one species loses electrons, which are then donated and gained by the other species.
• Being able to work out if the oxidation number of atoms in a reaction enables you to work out if a redox reaction is a disproportionation reaction too.
• In a disproportionation reaction, a species is both oxidised and reduced, seen as both an increase and decrease in oxidation number for that species.
• Half equations are used to show the separate oxidation and reduction reactions that occur in a redox reaction. They must be balanced in terms of the species present and the charges of the species on both sides of the equation.
• To balance half equations, firstly balance all the species apart from oxygen and hydrogen. Next, balance oxygen using H2O, and balance hydrogen using H+ ions. Finally, balance the charges using e- (electrons).
• Half equations can be combined in order to determine the overall redox reaction. In order to do this, the number of electrons must be the same for both half equations. This can be done by scaling up the number of moles. Once the half equations are combined, the electrons should be cancelled out on each side of the equation.