Oxidation Numbers
- 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
- Redox reactions occur when a substance is simultaneously reduced and oxidisedin 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
Metal and Nonmetal Redox Reactions
- 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
Metal and Acids Redox Reactions
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 Reactions
- 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)
Balancing Redox Equations
- 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