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Redox Equations (Oxford AQA International A Level Chemistry)

Revision Note

Alexandra Brennan

Written by: Alexandra Brennan

Reviewed by: Stewart Hird

Oxidation States

  • Oxidation states are used to:

    • Tell if oxidation or reduction has taken place

    • Work out what has been oxidised and/or reduced

    • Construct half equations and balance redox equations

  • Oxidation states are also referred to as oxidation numbers

  • A positive oxidation states = loss of electrons

    • The more positive the number, the more the element has been oxidised

  • A negative oxidation state = gain of electrons

    • The more negative the number, the more the element has been reduced

Oxidation State Rules

  • We can use the following rules to determine the oxidation state of an element

Rule

Example

Uncombined elements have an oxidation state of 0

  • H2

  • Cl2

  • Na

Some elements have the same oxidation state in their compounds with some exceptions

  • Group 1 elements are always +1

  • Group 2 element are always +2

  • Aluminium is always +3

  • Hydrogen is always +1 except in metal hydrides where it is -1

  • Fluorine is always -1

  • Oxygen is always -2 except in peroxides where is it -1 and OF2, where it is +2

  • Chlorine is always -1 except in compounds with F and O where it has a positive value

The sum of all of the oxidation states in a compound is equal to 0

  • In NaCl:

    • Na= +1

    • Cl = -1

  • Sum of oxidation states= 0

The oxidation state of an element in a monoatomic ion is always equal to its charge

  • Zn2+= +2

  • Fe3+ = +3

  • Cl- = -1

The sum of the oxidation states in an ion is equal to the charge on the ion

  • In SO42-

    • S= +6

    • O= (-2 x 4) = -8

  • Sum of oxidation states= -2

In a compound, the more electronegative element is given the negative oxidation state

  • In F2O

    • F is more electronegative so is (-1 x 2) = -2

    • O = +2

  • Sum of oxidation states= 0

Worked Example

Deducing oxidation states

Give the oxidation state of the bold atoms in these compounds or ions.

  1. P2O5

  2. SO42– 

  3. H2S

  4. Al2Cl6

  5. NH3 

  6. ClO2– 

  7. CaCO3 

Answers:

  1. P2O5 

    • 5 O atoms = 5 x (–2) = –10

    • The overall charge of the compound = 0

    • 2 P atoms = +10

    • Oxidation state of 1 P atom = (+10) / 2 = +5

  2. SO42– 

    • 4 O atoms = 4 x (–2) = –8

    • The overall charge of the compound = –2

    • The oxidation state of 1 S atom = +6

  3. H2S

    • 2 H atoms = 2 x (+1) = +2

    • The overall charge of the compound = 0

    • The oxidation state of 1 S atom = –2

  4. Al2Cl6 

    • 6 Cl atoms = 6 x (–1) = –6

    • The overall charge of the compound = 0

    • 2 Al atoms = +6

    • The oxidation state of 1 Al atom = (+6) / 2 = +3

  5. NH3 

    • 3 H atoms = 3 x (+1) = +3

    • The overall charge of the compound = 0

    • The oxidation state of 1 N atom = –3

  6. ClO2– 

    • 2 O atoms = 2 x (–2) = –4

    • The overall charge of the compound = –1

    • The oxidation state of 1 Cl atom = +3

  7. CaCO3 

    • 3 O atoms = 3 x (–2) = –6

    • 1 Ca atom = +2

    • The overall charge of the compound = 0

    • The oxidation state of 1 C atom = +4

Balancing Redox Reactions

  • Redox reactions are reactions in which oxidation and reduction take place together

    • While one species is being oxidised, another is being reduced in the same reaction

  • For example:

Cu2++ Zn → Zn2+ + Cu

  • Cu has been reduced from +2 to 0

  • Zn has been oxidised from 0 to +2

  • You should be able to split full or ionic equations into half equations and be able to identifying the species oxidised and reduced

Worked Example

In each of the following equations, write the half equations and state which reactant has been oxidised and which has been reduced.

  1. 2Na + Cl2 → 2NaCl

  2. Mg + Fe2+ → Mg2+ + Fe

  3. CO + Ag2O → 2Ag + CO2 

Answers:

  • Answer 1:

    • Na → Na+ + e-

    • Cl2 + 2e- → 2Cl-

    • Oxidised: Na as the oxidation state has increased from 0 to +1

    • Reduced: Cl2 as the oxidation state has decreased from 0 to –1

  • Answer 2:

    • Mg + 2e- → Mg2+ + 2e-

    • Fe2+ + 2e- → Fe

    • Oxidised: Mg as the oxidation state has increased by 2

    • Reduced: Fe2+ as the oxidation state has decreased by 2

  • Answer 3:

    • 2Ag++ 2e- → 2Ag

    • C2+→ C4+ + 2e-

    • Oxidised: C as it has lost electrons

    • Reduced: Ag as it has accepted electrons

Balancing redox equations

  • Oxidation states can be used to balance chemical equations

  • Roman numerals between brackets are used to show the ox. state of an atom that can have multiple oxidation states, e.g.:

    • Fe(II) = iron with an oxidation state of +2

    • Fe(III) = iron with an oxidation state of +3

  • You should be able to combine half equations and produce an overall redox equation

  • Using simple half equations involves adjusting the number of electrons and other coefficients to produce the overall equation:

Example:

Al → Al3+ + 3e-

O2 + 4e- → 2O2-

4Al + 3O2 → 2Al2O3

  • In this example it's easy to see that the lowest common factor is 12, so the first half equation is multiplied by 4 and the second one by 3, and then they are combined

  • More complicated examples involve being given an unbalanced redox equation and working through the redox changes, as well as having to add extra species such as H+ and H20 to balance the overall equation

  • Go through these steps to balance a redox equation:

    • Write the unbalanced equation and identify the atoms which change in oxidation state

    • Deduce the oxidation state changes

    • Balance the oxidation state changes

    • Balance the charges

    • Balance the remaining atoms

Worked Example

Writing overall redox reactions

Manganate(VII) ions, MnO4–, react with Fe2+ ions in the presence of acid, H+, to form Mn2+ ions, Fe3+ ions and water.

Write the overall redox equation for this reaction.

Answer:

  • Step 1: Write the unbalanced equation and identify the atoms which change in oxidation number:

Electrochemistry Step 1 Writing overall redox reactions, downloadable AS & A Level Chemistry revision notes

  • Step 2: Deduce the oxidation number changes:

Electrochemistry Step 2 Writing overall redox reactions, downloadable AS & A Level Chemistry revision notes

  • Step 3: Balance the oxidation number changes:

Electrochemistry Step 3 Writing overall redox reactions, downloadable AS & A Level Chemistry revision notes

  • Step 4: Balance the charges:

Electrochemistry Step 4 Writing overall redox reactions, downloadable AS & A Level Chemistry revision notes

  • Step 5: Balance the atoms:

Electrochemistry Step 5 Writing overall redox reactions, downloadable AS & A Level Chemistry revision notes

Examiner Tips and Tricks

It's really important that you learn the rules for oxidation states as these are not given to you in an exam!

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    Author
    Reviewer
    Alexandra Brennan

    Author: Alexandra Brennan

    Expertise: Chemistry

    Alex studied Biochemistry at Newcastle University before embarking upon a career in teaching. With nearly 10 years of teaching experience, Alex has had several roles including Chemistry/Science Teacher, Head of Science and Examiner for AQA and Edexcel. Alex’s passion for creating engaging content that enables students to succeed in exams drove her to pursue a career outside of the classroom at SME.

    Stewart Hird

    Author: Stewart Hird

    Expertise: Chemistry Lead

    Stewart has been an enthusiastic GCSE, IGCSE, A Level and IB teacher for more than 30 years in the UK as well as overseas, and has also been an examiner for IB and A Level. As a long-standing Head of Science, Stewart brings a wealth of experience to creating Topic Questions and revision materials for Save My Exams. Stewart specialises in Chemistry, but has also taught Physics and Environmental Systems and Societies.