5.6- Redox Titrations and Electrode Potentials

variable oxidation states

• when one element can have different oxidation states

• common in transition metals

redox reactions

a reaction where both oxidation and reduction take place

oxidising agents

a substance that gains electrons to oxidise another substance

reducing agent

a substance that donates electrons to reduce another substance

forming redox equations

1. Form two ionic half equations. Half equations can be balanced by adding:

• e^-

• H⁺

• OH^-

• H₂O

2. add the two half equation to make an overall equation. overall equation must have the same number of electrons on each side of the equation so when balancing you can only add:

• H⁺

• OH^-

• H₂O

redox titrations

• used to determine unknown concentration of a substance

• involves transfer of electrons from one species to another resulting in oxidation and reduction

• most don’t require an indicator since redox reactions are often self indicating→ colour change between two oxidation states

potassium manganate (VII) titrations

• manganate (VII) is oxidising agent→ reduced to Mn²⁺

• iron (II) is reducing agent→ oxidised to Fe³⁺

• reaction mixture must be acidified so xs acid is added to iron (II) ions before reaction begins

choice of acid in permanganate titrations

• must not react with manganate ions

• dilute sulfuric acid→ does not oxidise under these conditions and does not react with manganate ions

why other acids are not suitable for permanganate titration

• HCl→ oxidised to chlorine by manganate ions

• HNO₄→ is an oxidising agent so may oxidise substance being analysed

• CH₃COOH→ weak acid, conc. of H⁺ ions is insufficient

• conc. H₂SO₄→ may oxidise substance being analysed

end point in permanganate titration

• potassium permanganate acts as its own indicator→ reacts with Fe²⁺.

• potassium permanganate solution is purple

  • burette must have white numbering to ensure readability of titres.

• manganese ions have pale pink colour→ in low concentration so solution looks colourless

• when all Fe²⁺ have reacted, pale pink tinge appears due to excess manganate ions

iodine-thiosulfate titration equation

2S₂O₃^2– (aq) + I₂ (aq) → 2I^–(aq) + S₄O₆^2– (aq)

procedure for thiosulfate titrations

1. Record mass of alloy

2. dissolve alloy in conc HNO₃

3. dilute acidic Cu²⁺ solution in a volumetric flask

4. pipette 25cm³ Cu²⁺ into conical flask

5. add Na₂CO₃ to neutralise xs HNO₃:

   • effervescence

   • CuCO₃ ppt forms

6. add CH₃COOH dropwise until all CuCO₃ reacts to form Cu(CH₃COO)₂:

   • weak acid used to ensure Cu²⁺ solution is as close to neutral as possible

7. titrate liberated iodine with sodium thiosulfate

8. as colour becomes straw yellow add starch indicator→ clarifies endpoint

9. END POINT: when blue/black fades, leaving white ppt. in colourless solution

what is thiosulfate titration used for

• to determine conc. of oxidising agent→ oxidises iodide ions to iodine molecules

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