Edexcel Chemistry - Topic 3: Redox I

• All uncombined elements have an oxidation number of ZERO
• The oxidation numbers of the elements in a neutral compound add up to ZERO
• The oxidation number of a monatomic ion is equal to the IONIC CHARGE
• In a polyatomic ion, the sum of the individual oxidation numbers of the elements equals the IONIC CHARGE
• Hydrogen is +1 except in metal hydrides (e.g. NaH) where it is -1
• Fluroine is -1
• Chlorine, Bromine and Iodine are -1 except in compounds with oxygen and fluorine
• Oxygen is -2 except in peroxides (e.g. H2O2) where it is -1 and with fluorine
• Group 1 metals have a charge of +1
• Group 2 metals have a charge of +2
• Transition metals have their charge in ROMAN NUMERALS (e.g. Iron(III) = +3)
• Elements or ions that have variable oxidation states must have their oxidation number stated as roman numerals in their name e.g. NaClO = Sodium Chlorate(I) / NaClO3 = Sodium Chlorate (II)(

• Redox reactions occur when a substance is simultaneously reduced and oxidised in a reaction
• Reduction is the gain of electrons
• Oxidation is the loss of electrons
• Oxidation Is Loss | Reduction Is Gain
• Half equations show parts of a chemical equation involved in either reduction or oxidation
• Reduction half equations have the electrons on the left
  Br2(aq) + 2e- --> 2Br-(aq)
• Oxidation half equations have the electrons on the right
  2I-(aq) --> I2(aq) + 2e-
• The overall equation would be:
  Br2(aq) + 2I-(aq) --> I2(aq) + 2Br-(aq) 
• A reducing agent is an electron donor
• An oxidising agent is an electron acceptor
• In the equation the oxidising agent is the bromine and the reducing agent the iodide ions

• Metals generally oxidise to form ions by losing electrons causing an increase in oxidation number e.g. Zn --> Zn2+ + 2e-
• Non-metals generally reduce to cause a decrease in oxidation number
  e.g. Cl2 + 2e- --> 2Cl-
• Oxygen gas often reduces as its oxidation number decreases e.g. 4Li + O2 --> 2Li2O (Li = 0 -> +1) (O = 0 -> -2)
• Hydrogen when reacting with an oxide or oxygen often oxidises because its oxidation number increases from 0 to +1
• Nitrogen upon decomposing of a Nitrate(V) reduces from +5 to +4 in NO2
• When reacting ammonia with Sodium Chlorate(I) the Chlorine reduces from +1 to -1 in NaCl and the Nitrogen oxidses from -3 to -2 in N2H4

ACID + METAL --> SALT + HYDROGEN

2HCl + Mg --> MgCl2 + H2

• Hydrogen reduces as its oxidation number decreases from +1 to 0
• Magnesium oxidses because its oxidation number increases from 0 to +2
• The reaction will cause effervescence as hydrogen gas is evolved and the metal will dissolve

• Disproportionation is the name of a reaction where an element in a single species simultaneously oxidises and reduces
• Cl2 + H2O --> HClO + HCl
• Chlorine is both simultaneously reduced and oxidised changing its oxidation number from 0 to -1 and +1
• 2Cu+ --> Cu + Cu2+
• Copper(I) ions (+1) when reacting with sulphuric acid will disproportionate to Cu2+ (+2) and Cu metal (0)

• Work out the oxidation numbers for the element being oxidised/reduced
• Add electrons to the change in oxidation number
• Add H2O to balance out any Oxygens
• Add H+ to balance out the Hydrogens in H2O
• Ensure that the sum of the charges on the reactant side equals the sum of the charges on the product side
• Multiply equations to ensure that the number of atoms on each side are balanced