Chapter 23 - Redox and Electrode Potentials

Question 1

What is an oxidising agent?

Answer 1

Takes electrons from species being oxidised. Contains the species that is being reduced

Question 2

What is a reducing agent?

Answer 2

Adds electrons to the species being reduced. Contains the species that is being oxidised

Question 3

What are some of the important oxidation numbers?

Answer 3

Uncombined element = 0
Combined hydrogen = +1
Combined oxygen = -2
Ion of element = ionic charge

Question 4

What are the steps for constructing a redox equation from half equations?

Answer 4

1. balance the electrons
2. add and cancel electrons
3. cancel any species that are on both sides of the equation

Question 5

What are the steps for constructing a redox equation from oxidation numbers?

Answer 5

1. assign oxidation numbers to identify atoms that change their oxidation number
2. balance only the species that contain elements that have changed oxidation number
3. balance any remaining atoms

Question 6

Where do electrons appear in half equations for reduction?

Answer 6

Left hand side
e.g. Cl2 + 2e- → 2Cl-

Question 7

Where do electrons appear in half equations for oxidation?

Answer 7

Right hand side
e.g. Cu → Cu2+ + 2e-

Question 8

How can oxidation numbers be used to write a half equation?

Answer 8

1. assign oxidation numbers and the change in oxidation number
2. balance e- (e.g. decrease in oxidation number of 5 requires 5e- on left hand side)
3. balance any remaining atoms and predict any further species.

Question 9

What are the two common redox titrations?

Answer 9

• Potassium manganate (VII) under acidic conditions
• Sodium thiosulfate for determination of iodine

Question 10

What are manganate titrations used for the analysis of?

Answer 10

Reducing agents e.g. iron ions or ethanedioic acid

Question 11

How does manganate change in redox titrations?

Answer 11

Reduced from MnO4 - to Mn2+

Question 12

What are iodine-thiosulfate titrations used for the analysis of?

Answer 12

Oxidising agents e.g. Cu ions or chlorate ions

Question 13

What is oxidised and what is reduced in iodine-thiosulfate reactions?

Answer 13

Thiosulfate ions are oxidised
Iodine is reduced

Question 14

What are the 2 half equations for iodine-thiosulfate reactions?

Answer 14

Oxidation = 2S2O3 2- → S4O6 2- + 2e-
Reduction = I2 +2e- → 2I-

Question 15

What is the overall equation for iodine-thiosulfate reactions?

Answer 15

2S2O3 2- + I2 → 2I- + S4O6 2-

Question 16

What is the method for iodine-thiosulfate titrations?

Answer 16

1. excess of acidified iodide ions and known volume of oxidising agent react in a conical flask
2. products of reaction 1 left in a conical flask, titrated with known conc of sodium thiosulfate
3. add starch when the solution has turned straw yellow, the solution will go blue black and the endpoint is when the conical flask goes colourless

Question 17

What is an electrochemical cell?

Answer 17

A voltaic cell which converts chemical energy into electrical energy via the transfer of electrons

Question 18

What is a half cell?

Answer 18

Contains the chemical species present in a redox half equation

Question 19

What is a metal/metal ion half cell?

Answer 19

A simple half cell made up of a metal rod dipped into a solution of its aqueous metal ion. This sets up an equilibrium
M 2+ + 2e ⇌ M (s)
One metal half cell on its own will have no net electron transfer.

Question 20

What is an ion/ion half cell?

Answer 20

A half cell containing ions of the same element in different oxidation states. This sets up a redox equilibrium:
Fe3+ + e- ⇌ Fe2+
Uses a platinum electrode to transport electrons in and out of the half cell

Question 21

When two metal/metal ion half cells are connected, how do you know where oxidation and reduction occur?

Answer 21

Oxidation happens at the more reactive electrode (e- are more easily lost)
Reduction happens at the less reactive electrode (e- more easily gained)
This creates an electrode potential difference

Question 22

What is the standard electrode potential?

Answer 22

The e.m.f. of a half cell connected to a standard hydrogen half cell under standard conditions (state standard conditions!!! 1mol dm-3 1 atm/100kPa and 298K)

Question 23

How is a standard hydrogen electrode set up?

Answer 23

Acid solution containing 1.0 mol dm-3 H+ with platinum electrode inserted. Glass tube with holes in it to allow bubbles of H2 gas to escape. H2 at 298K and 100kPa

Question 24

How do you connect 2 half cells?

Answer 24

• electrodes connected by a wire to allow a controlled flow of electrons
• two solutions connected with a salt bridge to allow ions to flow

Question 25

What is a salt bridge?

Answer 25

Typically a concentrated solution of an electrolyte which doesn't react with either solution e.g. filter paper soaked in aq KNO3

Question 26

What does the E o value tell you?

Answer 26

More negative = greater tendency to lose electrons and undergo oxidation More positive = greater tendency to gain electrons and undergo reduction

Question 27

How do you calculate standard electrode potential?

Answer 27

E cell = E (positive electrode) - E (negative electrode)

Question 28

How can you predict if a redox reaction is feasible?

Answer 28

Using the electrochemical series (standard electrode potentials)

Question 29

What does it tell you if a redox system has a more negative standard E cell value?

Answer 29

More negative = greater tendency to be oxidised and LOSE electrons Therefore better REDUCING AGENTS

Question 30

What does it tell you if a redox system has a more positive standard E cell value?

Answer 30

More positive = greater tendency to be reduced and GAIN electrons Therefore better OXIDISING AGENTS

Question 31

How do you connect two redox systems from the electrochemical series?

Answer 31

Connect the two with an anti clockwise circuit (more negative redox system on the top) - keep more positive redox system the SAME - CHANGE the more negative redox system

Question 32

How does changing concentration affect standard electrode potentials?

Answer 32

If concentration of solution is not 1 mol dm-3, then value of the electrode potential will be different from standard value Shift to the side with less electrons makes electrode potential less negative Shift to side with more electrons makes electrode potential more negative

Question 33

What are the 3 types of cell?

Answer 33

Primary, secondary and fuel

Question 34

What are primary cells?

Answer 34

Cells which are non renewable and designed to be used once only. When in use, electrical energy is produced by oxidation and reduction at the electrodes. Most modern primary cells = Zn/MnO2 and KOH alkaline electrolyte

Question 35

What are some examples where primary cells are used?

Answer 35

Low current, long storage devices like wall clocks or smoke detectors

Question 36

What are secondary cells?

Answer 36

Rechargeable cells, where the cell reaction produced electrical energy can be reversed during recharging. Common examples = nickel-metal hydride or lead-acid

Question 37

What are some examples where secondary cells are used?

Answer 37

Car batteries, radios and torches, laptop and phone batteries

Question 38

What are fuel cells?

Answer 38

A fuel cell uses the energy from the reaction of a fuel with oxygen to create a voltage. Fuel and oxygen flow into fuel cell, products flow out. Can operate continuously provided fuel and oxygen are supplied, do not have to be recharged Common examples = hydrogen fuel cells or methanol

Question 39

What are the advantages and disadvantages of fuel cells over fossil fuels?

Answer 39

Advantages: only H2O produced/non polluting, greater efficiency Disadvantages: hydrogen gas is difficult to store, and difficult to manufacture initially

Question 40

What is the flow of electrons in 2 half cells?

Answer 40

Flows from more negative to half cell with more positive standard electrode potential