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Half Equations (DP IB Chemistry: SL): Revision Note

Philippa Platt

Written by: Philippa Platt

Reviewed by: Richard Boole

Updated on

DP Chemistry: SLDP Chemistry: HL

Half equations

How to write a half-equation

  • A half-equation shows the oxidation or the reduction on its own, with the electrons included

  • To write a half-equation in acidic solution:

    1. Balance the atoms other than O and H

    2. Balance O by adding H2O

    3. Balance H by adding H+

    4. Balance the charge by adding electrons, e-

  • To combine two half-equations, multiply one or both so the numbers of electrons are equal, then add them and cancel the electrons

Worked Example

Manganate(VII) ions, MnO4-, oxidise Fe2+ ions to Fe3+ in acidic solution and are reduced to Mn2+. Write the two half-equations and the overall equation.

Answer:

Reduction:

MnO4- → Mn2+

  • Mn is already balanced

  • Add 4H2O to the right to balance O:

MnO4- → Mn2+ + 4H2O

  • Add 8H+ to the left to balance H:

MnO4- + 8H+ → Mn2+ + 4H2O

  • The charge is 7+ on the left and 2+ on the right, so add 5e- to the left:

MnO4- (aq) + 8H+ (aq) + 5e- → Mn2+ (aq) + 4H2O (l)

Oxidation:

Fe2+ (aq) → Fe3+ (aq) + e-

Overall:

  • Multiply the oxidation half-equation by 5 so both have 5 electrons:

5Fe2+ (aq) → 5Fe3+ (aq) + 5e-

  • Combine the reduction and oxidation half equations

MnO4- (aq) + 8H+ (aq) + 5e- + 5Fe2+ (aq) → Mn2+ (aq) + 4H2O (l) + 5Fe3+ (aq) + 5e-

  • Cancel the electrons:

MnO4- (aq) + 8H+ (aq) + 5Fe2+ (aq) → Mn2+ (aq) + 4H2O (l) + 5Fe3+ (aq)

  • The oxidation number method below gives the same overall equation

How to balance a redox equation

  • Oxidation numbers can be used to balance chemical equations

  • Go through these steps to balance a redox equation:

    1. Write the unbalanced equation and identify the atoms which change in ox. no.

    2. Deduce the oxidation number changes

    3. Balance the oxidation number changes

    4. Balance the charges

    5. Balance the atoms

Worked Example

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 the reaction

Answer:

  1. Write the unbalanced equation and identify the atoms which change in oxidation number 

    Step 1: unbalanced equation MnO4- + Fe2+ + H+ -> Mn2+ + Fe3+ + H2O with oxidation numbers written under each species (Mn +7 to +2, Fe +2 to +3).

  2. Deduce the oxidation number changes

    Step 2: oxidation number change for Mn is -5 (+7 to +2) and for Fe is +1 (+2 to +3).

  3. Balance the oxidation number changes

    Step 3: 5 Fe2+ placed on each side so the oxidation number changes balance: 1 x (-5) = -5 and 5 x (+1) = +5.

  4. Balance the charges

    Step 4: charge balance ignoring H+: left 9+, right 17+, so 8 H+ are needed on the left.

  5. Balance the atoms

    Step 5: final balanced equation MnO4- + 5Fe2+ + 8H+ -> Mn2+ + 5Fe3+ + 4H2O.

Redox titrations

  • In titrations, the concentration of a solution is determined by titrating with a solution of known concentration.

  • In redox titrations, an oxidising agent is titrated against a reducing agent

    • Electrons are transferred from one species to the other

  • Indicators are sometimes used to show the endpoint of the titration

  • However, most transition metal ions naturally change colour when changing the oxidation state

  • There are two common redox titrations you should know about:

    • Manganate(VII) titrations

    • Iodine-thiosulfate titrations

Manganate(VII) titrations

  • A redox reaction occurs between acidified manganate(VII) ions and iron(II) ions:

MnO4– (aq) +  8H+ (aq) +  5Fe2+ (aq)  → Mn2+ (aq) +  5Fe3+ (aq) + 4H2O (l)

  • This reaction needs no indicator because manganate(VII) is a strong purple colour: it is decolourised by Fe2+ as it is added, and the first permanent pale pink colour marks the endpoint, so the titration is self-indicating

  • This reaction is often used for the analysis of iron for example in iron tablets (health supplements)

Iodine-thiosulfate titrations

  • A redox reaction occurs between iodine and thiosulfate ions:

2S2O32– (aq) + I2 (aq) → 2I–(aq) + S4O62– (aq)

  • The light brown/yellow colour of the iodine turns paler as it is converted to colourless iodide ions

  • When the solution is a straw colour, starch is added to clarify the endpoint

  • The solution turns blue/black until all the iodine reacts, at which point the colour disappears.

  • This titration can be used to determine the concentration of an oxidising agent, which oxidises iodide ions to iodine molecules

  • The amount of iodine is determined from titration against a known quantity of sodium thiosulfate solution

  • This reaction can be used for the analysis of chlorine in bleach

Worked Example

A health supplement tablet containing iron(II)sulfate was analysed by titration. A tablet weighing 2.25 g was dissolved in dilute sulfuric acid and titrated against 0.100 mol dm-3 KMnO4.The titration required 26.50 cm3 for a complete reaction. Calculate the percentage by mass of iron in the tablet.

Answer:

  1. Write the balanced equation for the reaction

    • oxidation: Fe2+ (aq)  →   Fe3+ (aq)  + e-

    • reduction: MnO4- (aq) + 8H+ (aq)  + 5e-   →   Mn2+ (aq)  + 4H2O (l)

    • overall: MnO4- (aq) + 8H+ (aq)  + 5Fe2+ (aq)   →   Mn2+ (aq)  + 4H2O (l) + 5Fe3+ (aq)
       

  2. Determine the amount of MnO4- used in the titration

    • moles of MnO4-  = 0.0265 dm3  x  0.100 mol dm-3 = 0.00265 mol
       

  3. Determine the amount of iron in the reaction

    • From the equation for the reaction, we know the reacting ratio  MnO4- : Fe2+ = 1: 5

    • ∴ moles of Fe2+ = 0.00265 mol MnO4- x 5 = 0.01325 mol
       

  4. Convert moles into the mass of iron

    • Mass of iron = 0.01325 mol x 55.85 gmol-1 = 0.740 g
       

  5. Find the percentage of iron in the tablet

    • ∴ % Fe in the tablet = (0.740/ 2.25)  x 100 = 32.9%

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Philippa Platt

Author: Philippa Platt

Expertise: Chemistry Content Creator

Philippa has worked as a GCSE and A level chemistry teacher and tutor for over thirteen years. She studied chemistry and sport science at Loughborough University graduating in 2007 having also completed her PGCE in science. Throughout her time as a teacher she was incharge of a boarding house for five years and coached many teams in a variety of sports. When not producing resources with the chemistry team, Philippa enjoys being active outside with her young family and is a very keen gardener

Richard Boole

Reviewer: Richard Boole

Expertise: Curriculum Expert

Richard has taught Chemistry for over 15 years as well as working as a science tutor, examiner, content creator and author. He wasn’t the greatest at exams and only discovered how to revise in his final year at university. That knowledge made him want to help students learn how to revise, challenge them to think about what they actually know and hopefully succeed; so here he is, happily, at SME.