23) Redox and electrode potentials

Question 1

What two things will there always be in a redox reaction?

Answer 1

oxidising agent

reducing agent

Question 2

Define oxidising agent

Answer 2

a reagent that oxidises (takes e- from) another species

Question 3

Define reducing agent

Answer 3

a reagent that reduces (adds e- to) another species

Question 4

What should you do when writing an overall equation for redox reactions?

Answer 4

balance the changes in oxidation number and then balance any remaining atoms afterwards

Question 5

If one species in an overall redox equation has an oxidation number change of +6, and a second species has an oxidation number change of -1, what should be done?

Answer 5

multiply second species by 6 to give -6

Question 6

Describe manganate (VII) titrations

Answer 6

a standard solution of KMnO4 is added to burette with excess of dilute sulfuric acid
as it reacts it is decolourised from deep purple colour until end point = permanent pink colour - indicating an excess of MnO4 - ions (self-indicating)

Question 7

What is the purpose of excess of dilute sulfuric acid in a manganate (VII) titration?

Answer 7

provides H+ for the reduction of MnO4 - ions

Question 8

Why are burette readings taken from the top of the meniscus in a manganate (VII) titration?

Answer 8

manganate (VII) solution is a deep purple colour

Question 9

What can manganate (VII) titrations be used for?

Answer 9

• analysis of reducing agents e.g. Fe 2+ (analysing the purity of an iron (II) compound) or ethanedioic acid (COOH)2
• reduce MnO4 - to Mn 2+

Question 10

What can iodine/thiosulfate titrations be used to determine?

Answer 10

• chlorate ions, ClO- content in household bleach (HCl added to acidify solution)
• Cu 2+ content in copper (II) compounds
• Cu content in copper alloys (dissolved in conc. nitric acid + neutralisation to produce Cu 2+)

Question 11

Describe iodine/thiosulfate titrations

Answer 11

• a solution of oxidising agent to be analysed is added to a conical flask with an excess of potassium iodide - iodine produced resulting in a yellow-brown colour
• titrated with sodium thiosulfate solution (S2O3 2- oxidised and I2 reduced)
• when end point is approached, sol. has faded to a pale straw colour, a small amount of starch indicator is added -> blue-black solution which turns colourless at end point

Question 12

Define voltaic cell

Answer 12

a type of electrochemical cell which converts chemical energy into electrical energy

Question 13

What does a half cell contain?

Answer 13

the chemical species present in a redox half-equation

Question 14

How is a voltaic cell made?

Answer 14

by connecting together 2 different half-cells allowing e- flow

Question 15

Describe metal/metal ion half-cells

Answer 15

a metal rod dipped into a solution of its aqueous metal ion
at the phase boundary where the metal is in contact with its ions, an equilibrium is set up e.g. Zn 2+(aq) + 2e- <=> Zn(s)

Question 16

Describe ion/ion half-cells

Answer 16

contains ions of the same element in different oxidation states
there is no metal to transport e- into/out of the half-cell so an inert metal electrode made out of platinum is used

Question 17

What happens in an operating cell with two metal/metal ion half-cells connected?

Answer 17

the electrode with more reactive metal loses e- and is oxidised
the electrode with the less reactive metal gains e- and is reduced

Question 18

Define standard electrode potential E⦵

Answer 18

the emf of a half-cell compared with a standard hydrogen half-cell under standard conditions

Question 19

What is the standard electrode potential of a standard hydrogen half-cell under standard conditions?

Answer 19

0V

Question 20

How is standard electrode potential measured?

Answer 20

the half-cell is connected to a standard hydrogen electrode by a wire to allow controlled flow of e-
recorded by voltmeter

Question 21

Describe a standard hydrogen half-cell

Answer 21

platinum electrode
H2(g) pumped in
acid sol. 1moldm-3 H+(aq)

Question 22

What does a salt bridge typically contain?

Answer 22

a concentrated solution of electrolyte that does not react with either solution e.g. a strip of filter paper soaked in KNO3(aq)

Question 23

How is the equilibrium for standard electrode potentials shown?

Answer 23

with the forward reaction being the reduction

Question 24

How are electrode potentials sorted in a data reference table?

Answer 24

in order of E⦵ - most negative at top

Question 25

The more negative E⦵ value, the greater the tendency to _ e- and undergo _?

Answer 25

lose

oxidation

Question 26

The more positive E⦵ value, the greater the tendency to _ e- and undergo _?

Answer 26

gain

reduction

Question 27

E⦵cell = ?

Answer 27

E⦵(positive electrode) - E⦵(negative electrode)

Question 28

What should you remember when predicting redox reactions from a data reference table?

Answer 28

top-right, bottom-left rule
the strongest reducing agent is at the top right
the strongest oxidising agent is at the bottom left

Question 29

What should you remember when writing an overall equation from half-equations in a data reference table?

Answer 29

the oxidation half-equation is obtained by reversing the equilibrium

Question 30

What are the limitations of predicting redox reactions from a data reference table? (4)

Answer 30

• gives no indication of rate of a reaction
• if concentration is not equal to 1moldm-3 electrode potential will be different to standard value
• may not be under standard conditions
• many reactions are not aqueous (aqueous equilibria)

Question 31

Describe primary cells

Answer 31

• non-rechargeable
• electrical energy is produced by reduction and oxidation at electrodes (irreversible reactions)
• eventually chemicals will be used up, voltage decreases and battery goes flat
• alkaline (Zn/MnO4 and potassium hydroxide alkaline electrolyte, Zn oxidised -> ZnO and Mn in MnO2 reduced -> Mn2O3)

Question 32

Describe secondary cells

Answer 32

• rechargeable - cell reaction can be reversed during recharging so chemicals are regenerated
• lithium-ion and lithium ion polymer cells (when in use e- flow from negative -> positive electrode and Li+ from anode to cathode to keep electrical charge balance)

Question 33

Describe fuel cells

Answer 33

• uses the energy from the reaction of a fuel with oxygen to create voltage
• fuel and oxygen flow into cell and products flow out, electrolyte remains
• operate continually provided fuel and O2 supplied
• do not have to be recharged

Question 34

Give the overall equation for hydrogen fuel cells and the E⦵cell

Answer 34

H2(g) + 1/2 O2(g) -> H2O(l)

E⦵cell = 1.23V

Question 35

Suggest two advantages of using methanoic acid as the fuel in a fuel cell rather than hydrogen

Answer 35

liquid - easier to store and transport
greater cell potential / voltage
hydrogen is flammable