Redox

RULES

• In elements, the oxidation state is 0
• In ions, the charge on the ion is equal to the total oxidation states added together in the case of molecular ions. 
• Oxgyen is always -2
• Group 1 metals are always +1
• Group 2 metals are always +2
• Group 7 are always -1 except when bonded to a more electronegative atom.
• The more electronegative atom always has the -ve oxidation state.
• H is always +1 except in metal hydrides. 

Spectator ions are not included in half-equations.

For ionic half-equations, the oxidation numbers on each side of the equation do not balance.

RULES

• Balance all atoms except O + H
• Balance O with water
• Balance H with H+
• Balance charges with e-
• Construct the other half-equation
• Balance with e-
• Combine and cancel

Oxidation is the process where electrons are lost.

Reduction is the process where electrons are gained. 

OIL RIG - Oxidation is loss, reduction is gain. 

In a redox reaction, the reducing agent (reduces another substance) is oxidised and the oxidising agent (oxidises another substance) is reduced. 

METALS WITH OXYGEN

The metal loses electrons and is oxidised. Oxygen gains electrons and is reduced.

METALS WITH WATER

The metals loses electrons and is oxidised. The water gains electrons and is reduced. 

METALS WITH ACID

The metals loses electrons and is oxidised. The hydrogen ions gain electrons so are reduced. 

DISPLACEMENT

Displacement reactions involve a shift of electron density and so are redox reactions.

A disproportionation reaction takes place when a substance is both oxidised and reduced in the same reaction. 

E.g. 

• Breakdown of hydrogen peroxide into water and oxygen
• Copper(I) oxide with dilute sulfuric acid to form copper(0) and copper(II)
• Chlorine with water
• Chlorine with hot sodium hydroxide solution

These react to oxise something else, they themselves are reduced. The electrons are on the left of the equation: 

• Oxygen
• Chlorine
• Bromine
• Iodine
• Manganate(VII) in acid solution
• Dichromate(VI) in acid solution
• Iron(III) salts
• Hydrogen ions
• Hydrogen peroxide in absence of another oxidising agent
• Concentrated sulfuric acid

These react by reducing something else, they themselves are oxidised. The electrons are always on the right of the equation:

• Metals
• Iron(II) salts
• Acidified potassium iodide
• Thiosulfate
• Ethanedioic acid and ethanedioates
• Sulfuric(IV) acid
• Hydrogen peroxide in the presence of an acid and absence of a strong oxidising agent
• Hydrogen