Chapter 23

Question 1

What does an oxidising agent do, and what other feature does it have in redox reactions?

Answer 1

• It takes electrons from the species being oxidised
• It contains the species being reduced

Question 2

What does an reducing agent do, and what other feature does it have in redox reactions?

Answer 2

• It gives electrons to the species being reduced
• It contains the species being oxidised

Question 3

What needs to be remembered when naming the oxidising and reducing agent?

Answer 3

• Oxidising/ reducing agents are not just the element that is reduced/ oxidised, but the whole molecule/ ion

Question 4

How can redox reactions be constructed?

Answer 4

• By using half equations
• By using oxidation numbers

Question 5

What needs to be done when using half equations to construct redox reactions?

Answer 5

• The number of electrons in both equations needs to be the same so the equations can be combined
• This means one or both need to be scaled up

Question 6

How can oxidation numbers be used to write redox equations?

Answer 6

• You find which species had a change in oxidation number, and by how much
• You scale up the other species (wherever it appears in the reaction- watch out for any that are diatomic) so that the overall change in oxidation number sums to 0

Question 7

What can be done to balance redox reactions that take place in acidic, aqueous conditions?

Answer 7

• Start by balancing elements other than hydrogen and oxygen
• Balance oxygen by adding water (take note of any oxygens already present)
• Balance hydrogens by adding protons
• Add electrons to balance out charges if necessary (equations also have to be balanced in terms of charge)

Question 8

If a redox reaction is missing a product, how can this product be found?

Answer 8

• It is usually either water, a proton or a hydroxide ion
• If the reducing and oxidising agents are already present, make sure you do not add anything that has a change in oxidation number

Question 9

How can the concentrations of reactants in a redox reaction be calculated?

Answer 9

• Through redox titrations

Question 10

Give 2 examples of redox titrations.

Answer 10

• Fe2+/MnO4–
• I2/S2O32−

Question 11

Describe how to carry out a titration between manganate (VII) ions and iron (II) ions.

Answer 11

1. Standard solution of potassium manganate (VII) is added to the burette.
2. A pipette is used to add a measured volume of the reducing agent to the conical flask with an excess of dilute sulfuric acid (as it is needed for the reduction of MnO4-).
3. No indicator is needed as the reaction is self-indicating; manganate (VII) is purple while manganese (II) is colourless.
4. The end point is judged by the first permanent pale pink colour, as it shows an excess of manganate (VII) is present.
5. Burette readings are taken from the top of the meniscus, as potassium manganate (VII) has a deep purple colour that makes it hard to take readings from the bottom of the meniscus.
6. Repeat until you obtain concordant results.

Question 12

Give the equation for the reaction of manganate (VII) and iron (II). How can it be derived, and why is it important for their titration?

Answer 12

• Manganate (VII) is reduced to manganese (II), and iron (II) is therefore oxidised to iron (III)
• The reactions for both equations can be written, balanced and combined (iron’s would need to be scaled up by a factor of 5)
• This tells you they react in a 1:5 ratio
• MnO4- + 5Fe2+ + 8H+ -> Mn2+ + 5Fe3+ + 4H2O