AQA Chemistry Unit 2: 10 Redox Reactions

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Definitions

Redox: Reaction where both reduction and oxidation take place

Oxidation: The gain of oxygen, loss of hydrogen and loss of electrons

Reduction: The loss of oxygen, gain of hydrogen and gain of electrons

Oxidising Agent: Causes oxidation, causes gain of oxygen and loss of hydrogen and electrons. The agent itself is reduced in the reaction

Reducing Agent: Causes reduction, causes loss of oxygen and gain of hydrogen and electrons. The agent itself is oxidised in the reaction

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10.1 Oxidation and Reduction

  • Half equations are used to show the gain or loss of electrons
  • In a reaction if an ion takes no part in the actual reaction it is known as a spectator ion
  • In a chemical reaction if one species is oxidised another must be reduced

Oxidising and Reducing Agents

  • Reducing agents give away electrons so are electron donors
  • Oxidising agents accept electrons
  • Reducing agentscauses an increase in oxidation number thus their oxidation number decreases
  • Oxidising agents cause a decrease in oxidation number thus their oxidation number increases
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10.2 Oxidation States

  • Oxidation State: The number of electrons lost/gained by an atom in a compound compared to the uncombined atom
  • It is the charge on a simple ion
  • For an atom in a group which makes an ion it is the charge it would have if all other atoms were removed
  • The more electronegative an element the more negative its oxidation state

Rules

  • Every uncombined element has an oxidation state of zero
  • A positive number shows the element has lost electrons so has been oxidised
  • A negative number shows the element has gained electrons so has been reduced
  • The sum of all oxidation states in a compound is zero
  • The sum of oxidation states in an ion equals the charge on the ion
  • The most electronegative element has a negative oxidation state
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Oxidation States and Species

Uncombined elements: Zero

Oxygen in compounds: -2

Oxygen in peroxides: -1

Hydrogen in compounds: +1

Hydrogen in metal hydrides: -1

Group 1 metals: +1

Group 2 metals: +2

Group 6: -2

Group 7: -1

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Full Reactions to Half Equations

  • An equation for a redox reaction can be split into 2 equations
  • One shows oxidation and the other shows reduction

Rules:

  • Only one element in a half equation changes
  • It must balance for atoms
  • It must balance for charge

Additional Rules for Aqueous Solutions

  • Water provides a source of oxygen
  • Any surplus oxygen is converted to water by a reaction with acidic hydrogen ions
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Combining Half Equations to give Full Equations

  • Multiply whole half equations so the number of electrons given by the oxidised species equals the number needed by the species being reduced
  • No electrons therefore appear in any overall redox equation

Strategy:

  • Build the first half equation
  • Build the second half equation
  • Multiply (so no electrons are equal) and combine
  • Cancel species like H+ and H20 that appear on both sides of the redox equation
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