Redox

Redox

• In redox reactions one element is reduced and oxidised.
• Reduction and oxidation must occur together because if one element loses an electron another element must gain them.

Definitions

• Oxidation is the loss of electrons or increase in oxidation number.
• Reduction is the gain of electrons or decrease in oxidation number.
• Disproportionation a reaction in which an element is simultaneously oxidised and reduced.

Oxidation States

• Oxidation states allow us to work out which elements have been oxidised and reduced in a reaction.
• Oxidation states are "charges" assigned to each element in a reaction .

Rules

• The rules for assigning oxidation states are:
  • Elements are 0.
  • In compounds, H = +1 and O = -2.
  • Group 1 = +1, Group 2 = +2, Group 6 = -2, Group 7 = -1.
  • In a neutral compound the oxidation states add up to 0.
  • In an ion the oxidation states add up to the charge on the ion.

Example

• NO₃^- = -1
• Oxidation state of O = -2 x 3 = -6.
• Oxidation of N must be +1

Redox Reactions

• In reactions, metals generally lose electrons = oxidation.
• In reactions, non - metals generally gain electrons = reduced.
• Metals undergo redox reactions with dilute hydrochloric and dilute sulphuric acids.

Examples

Zn + 2HCl --> ZnCl₂ + H₂

• Zn goes from 0 to +2 = oxidation.
• H goes from +1 to 0 = reduction. 

Mg + H₂SO₄ --> MgSO₄ + H₂

• Mg goes from 0 to +2 = oxidation
• H goes from +1 to 0 = reduction