C3 - Redox I

Redox

• Oxidation is loss of electrons

• Reduction is gain of electrons

• When this happens at once, this is redox (OIL RIG)

• Oxidising agent

  • Accepts electrons and undergoes reduction

• Reducing agent

  • Donates electrons and undergoes oxidation

Oxidation numbers

• Oxidation

  • Increase of oxidation numbers

• Reduction

  • Decrease of oxidation number

• E.G. Na + 1/2 Cl2 ---> NaCl2

  • Sodium is oxidised because oxidation number increases from 0 to +1

  • Chlorine is reduced because oxidation number decreases from 0 to -1

Redox reactions

• Metal acid reactions are redox

  • Metal atoms lose electrons and are oxidised

  • Hydrogen ions gain electrons and are reduced

• Any species whose oxidation number does not change is a spectator ion

  • Not part of ionic equation

Rules of oxidation numbers

• Un-combined elements have an oxidation number of 0

• Atoms in molecules bonded to identical atoms (diatomic molecules) have oxidation number of 0

• Monatomic ion's ON is same as its charge (Fe3+ has ON of +3)

• In molecular ion, sum of ON is equal to overall charge

• Neutral compounds have ON sum of 0

• Combined oxygen has ON of -2 except in peroxides where its -1

  • In OF2 its +2

• Combined hydrogen usually has an ON of +1 except in metal hydrides where its -1

• Oxidation number of compound containing 2 non-metals is given to more electronegative element

Roman numerals

• Roman numerals are used after element's name to specify its oxidation number

  • Iron (II) has ON of +2

  • Iron (III) has ON of +3

-ate compunds

• Ions ending in -ate are made up of oxygen and another element

• Element bonding to oxygen has variable oxidation number

  • Roman numerals are used after compound's name to differentiate between these numbers

Reduction in acidic conditions

• Check oxi/red element is balanced on both sides

• Add H20 to balance out any oxygen in compound

• Add H+ to balance out any hydrogen in H20

• Use electrons to balance the charge

Reduction in alkaline conditions

• Check element is balanced

• Add 2x OH- needed to balance oxygens in compound

• Add H2O to balance OH-

• Use e- to balance the charge

Combining half equations

• Scale up and multiply half equations to ensure e- are equal

• Cancel out anything that appears on both sides (water, H+, OH- and e-)

Subtracting half equations

• Scale up half equation so moles of element is equal

• Subtract half equation from redox

• Place same number of e- from half equation on the opposite side

• Simplify if needed

Disproportionation

• When single element is both oxidised and reduced

• e.g. in Cl2 + 2OH- ----> ClO- + Cl- + H2O

  • Cl is both oxidised (increase in ON)

  • Cl is both reduced (decrease in ON)