Redox

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  • Created by: Sophie
  • Created on: 27-01-14 14:10

Redox - Introduction

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.
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Oxidation States

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
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Redox Reactions

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
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