5.1-5.3 Redox reactions and Oxidation and reduction

• Oxidation is the loss of electrons, gain of oxygen

• Reduction is gain of electrons, loss of oxygen

• ^^u can remember this by OILRIG^^

• ^^Oxidation is Loss of electrons Reduction is Gain of electrons^^

• sometimes oxidation and reduction happen together this reaction or process is called Redox Reactions.

• Redox reactions can also be defined in terms of electron transfer

• Oxidation is a reaction in which an element, ion or compound loses electrons

  • The oxidation state of the element is increased
  • This can be shown in a half equation, e.g. when silver reacts with chlorine, silver is oxidised to silver ions:
  • Ag → Ag+ + e-

• Reduction is a reaction in which an element, ion or compound gains electrons

  • The oxidation state of the element is decreased
  • This can be shown in a half equation, e.g. when oxygen reacts with magnesium, oxygen is reduced to oxide ions:

O2 + 4e- → 2O2-

• Oxidation and reduction in terms of electron transfer can be remembered by the mnemonic 'OIL RIG': Oxidation Is Loss of electrons, Reduction Is Gain of electrons

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Mnemonic to remember oxidation and reduction in terms of electron transfer

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Names using oxidation states

• Transition elements can bond in different ways by forming ions with different charges

• When naming, the charge on the ion is shown by using a Roman numeral after the element's name

  • e.g. iron can form ions with a 2+ charge, called iron(II) ions or a 3+ charge, called iron(III) ions

• The Roman numeral is the oxidation state of the element

• When iron reacts with oxygen to form iron oxide, the formula depends on the oxidation state of the iron ions

  • The compound where iron has a 2+ charge has the formula FeO and is called iron(II) oxide
  • The compound where iron has a 3+ charge has the formula Fe2O3 and is called iron(III) oxide

  Oxidation number

• The oxidation state (also called oxidation number) is a number assigned to an atom or ion in a compound which indicates the degree of oxidation (or reduction)

• The oxidation state helps you to keep track of the movement of electrons in a redox process

• It is written as a +/- sign followed by a number (not to be confused with charge which is written by a number followed by a +/- sign)

• E.g. aluminium in a compound usually has the oxidation state +3

Assigning the oxidation state

• Oxidation state refers to a single atom or ion only

• The oxidation number of a compound is 0 and of an element (for example Br in Br2) is also 0

• The oxidation number of oxygen in a compound is always -2 (except in peroxide R-O-O-R, where it is -1)

• For example in FeO, oxygen is -2 then Fe must have an oxidation number of +2 as the overall oxidation number for the compound must be 0

  • Table to show some common oxidation states of elements within compounds
  • Example redox equation: electron loss/gain and oxidation state
  • zinc + copper sulphate → zinc sulphate + copper
  • Zn + CuSO4 → ZnSO4 + Cu

• Writing all of the ions present and including state symbols we get:

  • Zn(s) + Cu2+(aq) + SO42-(aq) →Zn2+(aq) + SO42-(aq) + Cu(s)

• The spectator ions (those that do not change) are SO42-(aq), removing these we can write the ionic equation as:

  • Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

• By analysing the ionic equation, we can split the reaction into two half equations by adding in the electrons to show how the changes in charge have occurred.

• It then becomes clear that zinc has been oxidised as its oxidation state has increased from 0 in Zn to +2 in Zn2+and it has lost electrons:

  • Zn(s) → Zn2+(aq) + 2e-

• Copper ions have been reduced as the oxidation state has decreased from +2 in Cu2+ to 0 in Cu and they have gained electrons:

  • Cu2+(aq) +2e- → Cu(s)

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Oxidation and Reducing Agents

Oxidation Agent

• A substance that oxidises another substance, in so doing becoming itself reduced
• An oxidising agent gains electrons as another substance loses electrons
• Common examples include hydrogen peroxide, fluorine and chlorine

Reducing agent

• A substance that reduces another substance, in so doing becoming itself oxidised
• A reducing agent loses electrons as another substance gains electrons
• Common examples include carbon and hydrogen
• The process of reduction is very important in the chemical industry as a means of extracting metals from their ores

Example

• CuO + H2 → Cu + H2O

• In the above reaction, hydrogen is reducing the CuO and is itself oxidised as it has lost electrons, so the reducing agent is therefore hydrogen:

• H2 → 2H+ + 2e-

• The CuO is reduced to Cu by gaining electrons and has oxidised the hydrogen, so the oxidising agent is therefore copper oxide

• Cu2+ +2e- →  Cu

  

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