Redox (OCR AS Chemistry A): Revision Note

Oxidation number

What is oxidation number?

• The oxidation number (or oxidation state) of an atom is the theoretical charge it would have if bonding were completely ionic

• Oxidation numbers are used to:

  • Identify oxidation and reduction

  • Construct half-equations

  • Balance redox equations

How to assign oxidation numbers

• Follow these rules when working out oxidation numbers:

1. Uncombined elements have an oxidation number of 0

   • For example:

     • Hydrogen, H₂ = 0

     • Sulfur, S₈ = 0

     • Chlorine, Cl₂ = 0

     • Zinc, Zn = 0

2. The oxidation number of a simple ion equals its charge

   • For example:

     • Sodium ion, Na⁺ = +1

     • Iron(III) ion, Fe³⁺ = +3

     • Chloride ion, Cl^- = –1

     • Phosphide ion, P^3- = -3

3. In a neutral compound, the sum of oxidation numbers is 0

4. In a polyatomic ion, the sum of oxidation numbers equals the overall charge

   • For example:

     • Nitrate ion, NO₃^- = -1

     • Sulfate ion, SO₄^2- = –2

     • Phosphate ion, PO₄^3- = -3

     • Ammonium ion, NH₄⁺ = +1

5. In compounds:

   • Group 1 metals = +1

   • Group 2 metals = +2

   • Fluorine = –1

   • Hydrogen = +1

     • Hydrogen = –1 in metal hydrides

   • Oxygen = –2

     • Oxygen = –1 in peroxides

     • Oxygen = +2 in F₂O

6. The more electronegative element is assigned the negative oxidation number

Examples of oxidation numbers in compounds

• Sodium chloride, NaCl:

  • Na = +1

  • Cl = –1

  • So, the overall oxidation number is (+1) + (-1) = 0

• Carbon dioxide, CO₂:

  • O = –2 × 2 = –4

  • C = +4

  • So, the overall oxidation number is (-4) + (+4) = 0

• Hydrogen peroxide, H₂O₂:

  • H = +1 x 2 = +2

  • O = –1 x 2 = -2

    • Remember, oxygen is -1 in peroxides

  • So, the overall oxidation number is (+2) + (-2) = 0

Periodic Table patterns

• Metals usually form positive oxidation states, often matching their group number

• Transition metals can form multiple oxidation states

  • For example:

    • V = +2, +3, +4, +5

    • Cr = +2, +3, +6

    • Mn = +2, +7

• Non-metals typically have negative oxidation numbers

  • For example:

    • Cl = –1 in Cl⁻, or +1, +3, +5, +7 in oxoanions

    • Br^- = -1

    • S^2- = -2

    • N^3- = -3

Are oxidation numbers always whole numbers?

• Most are whole numbers, but some ions have an average oxidation number

• For example:

  • The tetrathionate ion, S₄O₆^2-

  • The tetrathionate ion contains six oxygen atoms

    • 6 x –2 = –12

  • The overall ion has an oxidation state of –2

    • This means that the oxidation numbers of the elements in the ion must add up to –2

  • So, the oxidation number of all four S atoms is +10

    • 4S + (6 x –2) = –2 4S = +10

  • Therefore, the oxidation number of S is +2.5

    • 4S = +10

    • S = 2.5

  • This specific question is a rare example showing a fractional oxidation number for the element in question / sulfur

  • This is because the sulfur atoms are in two different environments 
    

    

Using Roman numerals

• Roman numerals indicate oxidation states in compound names:

  • Iron(II) choride

    • Fe²⁺ = +2

  • Iron(III) oxide

    • Fe³⁺ = +3

  • Chlorate(I), ClO^-

    • Cl = +1

  • Chlorate(V), ClO₃^-

    • Cl = +5

Redox reactions & equations

• A redox reaction is a reaction where both oxidation and reduction occur

  • This involves a transfer of electrons and changes in oxidation number

Interpreting redox

• Redox reactions always involve both oxidation and reduction happening together

  • Oxidation is an increase in oxidation number (loss of electrons)

  • Reduction is a decrease in oxidation number (gain of electrons)

• A reducing agent donates electrons

  • So, a reducing agent is oxidised

• An oxidising agent accepts electrons

  • So, an oxidising agent is reduced

• For example:

Zn (s) + H₂SO₄ (aq) → ZnSO₄ (aq) + H₂ (g)            

• Analysing Zn:

  • Oxidation number increases from 0 to +2

  • Zn is oxidised

  • Therefore, zinc is the reducing agent

• Analysing H in H₂SO₄:

  • Oxidation number decreases from +1 to 0

  • H is reduced

  • Therefore, sulfuric acid is the oxidising agent