Redox and Electrode Potentails A2

Question 1

Why do you not need to add an indicator to a Manganate (VII) titration?

Answer 1

Manganate (VII) solution is purple. It changes to colourless when reduced to Mn2+.

Question 2

What colour change do you look for when adding manganate solution?

Answer 2

Colourless to permanent pink.

Question 3

What is different when reading the burette solution with a manganate titration?

Answer 3

Read from top of meniscus because dark purple colour means you can’t read the bottom of the miniscus.

Question 4

In a thiosulphate-iodine titration, what is the initial colour change?

Answer 4

Orange-brown fades to yellow as iodine is used up.

Question 5

What indicator is used at the end of a thiosulphate-iodine titration?

When do you add it?
Colour change?

Answer 5

Starch solution

Add only when faded to pale yellow colour.

Solution turns blue/black in I2 which disappears at end point in I-.

Question 6

What are the 3 electrochemical cells?

Answer 6

• Metal / metal ion half cells
• non-metal / non-metal ion half cells
• Metal ion / metal ion half cells

Question 7

What does a half cell compromise of?

Answer 7

An element in 2 different oxidation states.

Question 8

How do you set up a non-metal/non-metal ion half cell? (specifically hydrogen)

Answer 8

A glass tube with holes is used to allow bubbles of gas (non-metal - e.g. Hydrogen) to escape into the solution, either side of the platinum electrode.

Solution has 1 moldm^3 H+ from HCl

Question 9

Which electrode is used to measure standard electrode potentials of other half cells?

What are the conditions for standard electrodes?

Answer 9

Standard hydrogen electrode
(1 mol dm-3 H+ conc - HCl solution)

Hydrogen at 298K (25 C), 100kPa, Platinum electrode for inert electrodes, 1 mol dm-3 conc solution

Question 10

Which electrochemical cell do you not use a platinum electrode?

Answer 10

In the metal/metal ion half cells:
The metal is used as the rod and the solution contains the metal ions.

Question 11

How are two half cells connected to make a cell?

Answer 11

• Wires with a voltmeter between electrodes
• A salt bridge between solutions

Question 12

What is a salt bridge?

Answer 12

Concentrated solution of an electrolyte soaked into filter paper to connect two half cells.

Electrolyte = a susbstance that dissociates in water to form ions.

Question 13

What occurs between metal and metal ions/ non-metal and non-metal ions / metal ions in half cells?

Answer 13

Equilibrium

Question 14

If a standard electrode potential is very negative/positive, what does that mean for the likelihood of gaining/loosing electrons? Why is this?

Answer 14

Very positive = Likely to gain electrons / be reduced - equilibrium shifts right.
BECAUSE stronger oxidising agent.

Very negative = Likely to lose electrons / be oxidised - equilibrium shifts left.
BECAUSE stronger reducing agent.

Question 15

How do you work out the cell potential from electrode potentials of half cells? When is the reaction feasible?

Answer 15

Eθ of more positive electrode - Eθ of more negative electrode

Feasible if cell potential is positive (opposite to ΔG).

Question 16

Give 2 limitations of predictions of feasibility from electrode potentials.

Answer 16

1) Reaction may not take place because activation energy too high

2) Reaction may not take place with 1 mol dm-3 solutions

Question 17

What are the limitations to standard electrode potentials (Eθ)? (4 limitations)

Answer 17

• Non-standard conditions alter the value of Eθ.
• Half equations are equilibia so changes in conc will shift the position, which alters electron transfer.
• Reaction rates may be slow due to high activation energy.
• Not all reactions take place in aqueous solution.

Question 18

What is the difference between primary and secondary cells?

Answer 18

Primary cells = non-rechargable - chemical used up

Secondary cells = rechargable

Question 19

What does a fuel cell do? Describe them. What are the limitations?

Answer 19

Use energy from reactions with fuel and oxygen to produce a voltage.

• Reactants flow in and out - electrolyte remains in the cell.
• Can operate continuously as long as fuel and O2 provided.
• Do not have to be recharged.

LIMITATIONS
- Reactions will generally only take place if the difference between Eθ values is large (over 0.4v).

Question 20

How do you work out half equations from electrochemical series?

Answer 20

Write the equation with the more positive Eθ on the top and the equation with the more negative Eθ backwards, on the bottom.

(Reaction is clockwise)

Question 21

How do you construct a redox equation from half equations?

Answer 21

Identify which is reduction and which is oxidation.

Multiply equations so number of electrons are equal.

Add/merge them together.

Question 22

How do you construct a balanced redox equation from a reaction (reactants and products)?

Answer 22

e.g. 2HI + H2SO4 –> H2S + I2

Identify changing oxidation numbers and total changes.
I (2x-1 –> 2x0) total = +2 S (+6–>-2) total = -8

Balance the total increase and decrease in oxidation number.
I x 4 - total = +8
8HI + H2SO4 –> H2S + 4I2

Balance the H and O
8H and 4O = 4H2O
8HI + H2SO4 –> H2S + 4I2 + 4H2O

Question 23

What will many redox equations require the addition of to balance them and when?

Answer 23

Acidic conditions = H+
Alkaline conditions = OH-
If H and O don’t balance = H2O