Topic 9: Redox Processes

Question 1

LEO GER

Answer 1

Loss of electrons is oxidation (loss of hydrogen, gain of oxygen).

Gain of electrons is reduction (gain of hydrogen, loss of oxygen).

Question 2

Oxidation-reduction Reactions

Answer 2

The transfer of electrons from one species to another - one species loses electrons and the other species gains electrons.

Question 3

Define the oxidation number

Answer 3

The charge that an atom would have if all the covalent bonds in a compound were broken.

Question 4

How is the oxidation number written?

Answer 4

Number after the charge! eg. +2
(Opposite to when we write ion charges).

Question 5

What happens if the oxidation number reduces in size?

Answer 5

Reduction has occurred.

Question 6

What happens if the oxidation number increases in size?

Answer 6

Oxidation has occurred.

Question 7

What is the oxidation number for elements?

Answer 7

Zero. Eg. Cl2 is 0

Question 8

What is the oxidation number for simple ions?

Answer 8

The same as the charge. Eg. Cl- is -1

Question 9

What is the oxidation number for compounds?

Answer 9

Oxygen is -2 (except in peroxides when it is -1).
Hydrogen is +1 (except in metal hydrides when it is -1).
Chlorine is -1 (except with O and F).
Potassium is +1.
Work backwards. Compounds are neutral so = 0.
eg. MgO
x - 2 = 0
+2 - 2 = 0
So must be Mg+2O-2

Question 10

What is the oxidation number for polyatomic ions?

Answer 10

Adds up to the net charge on the ion.

Question 11

What ions do we leave out of redox equations?

Answer 11

Spectator ions. Eg. Na+, K+, SO4 2-

Question 12

What ions are present in redox equations?

Answer 12

Reactive ions (these always have a transition metal). Eg. Cu, MnO4 -, Cr2O7 2-, Fe2+

Question 13

How does reactivity affect redox reactions?

Answer 13

More reactive metals are stronger reducing agents as they lose their electrons more readily.

More reactive non-metals are stronger oxidising agents as they gain electrons more readily.

Question 14

Define the biological oxygen demand (BOD)

Answer 14

The amount of oxygen used to decompose the organic matter in a sample of water over a specified time period (usually five days) at a specified temperature.

Question 15

What are the two types of electrochemical cells?

Answer 15

Voltaic (galvanic) and electrolytic cells.

Question 16

Explain voltaic (galvanic) cells

Answer 16

In a voltaic cell, electricity is generated by the spontaneous redox reaction taking place.

Question 17

Explain electrolytic cells

Answer 17

In an electrolytic cell, electricity is used to drive non-spontaneous reactions.

Question 18

What is the standard hydrogen electrode used for?

Answer 18

The SHE is used as a reference to measure the electrode potential of other half cells.

Question 19

What are the standard conditions for the hydrogen electrode reaction?

Answer 19

Temperature = 298K
Pressure = 100kPa
Concentration = 1mol dm-3

Question 20

Define the standard electrode potential of a half-cell

Answer 20

The electrode potential relative to the standard hydrogen electrode (SHE), measured under standard conditions (100kPa, 298K, all solutions 1 mol dm-3).

Question 21

What way is the electron flow if the half cell contains a metal above Hydrogen in the Activity Series?

Answer 21

Electrons flow from the half-cell to the hydrogen electrode (electrode has a negative value).

Question 22

What way is the electron flow if the half cell contains a metal below Hydrogen in the Activity Series?

Answer 22

Electrons flow from the hydrogen electrode to the half cell (electrode has a positive value).

Question 23

(Spontaneity)
If the cell EMF is positive…

Answer 23

reaction is spontaneous.

Question 24

(Spontaneity)
If the cell EMF is negative…

Answer 24

reaction is non-spontaneous and the reverse reaction is spontaneous.

Question 25

Metals with low Eo values (most negative) will be . . .

Answer 25

the strongest reducing agents (themselves oxidised).

Question 26

Non-metals with high Eo values (most positive) will be . . .

Answer 26

the strongest oxidising agents (themselves reduced).

Question 27

Electrolysis of aqueous solutions: equations for the reduction of water at the cathode and oxidation of water at the anode

Answer 27

At the cathode:
2H20 + 2e —> H2 + 2OH- (-0.83V)

At the anode:
2H20 —> O2 + 4H+ + 4e

Overall equation:
2H20 —> 2H2 + O2

Question 28

What are the observed changes at the electrodes for the electrolysis of water?

Answer 28

• Colourless gas is produced at both electrodes. This is seen through bubbling (H2 at the cathode and O2 at the anode)
• Ratio of the volume of gases is 2H2 : O2 (i.e. two times the amount of H2 compared to O2)
• The pH at the cathode will decrease as H+ is produced. The pH at the anode will increase as OH- is produced.

Question 29

What is selective discharge determined by?

Answer 29

• At the cathode, it is determined by the E0 value of the ions
• At the anode, it is determined by the concentration of the ions in the electrolyte
• The materials from which the electrodes are made

Question 30

What species are unable to be oxidised further in electrolysis?

Answer 30

Sulfates and nitrates.

Question 31

Define electroplating

Answer 31

The process of using electrolysis to deposit a layer of a metal on top of another metal or other conductive object.

Question 32

Define Faraday

Answer 32

The charge carried by one mole of electrons (96500 Colulomb)

Question 33

Steps to work out the amount of product in electrolysis

Answer 33

Q=It
F=Q/96500
Mole ratios to find moles of product
m=nxM

Question 34

What factors affect the amount of product formed in electrolysis?

Answer 34

The charge on the ion. Cu+ produces twice as much than Cu2+.
i.e. the greater the charge the more electrons per reduction so less product formed.

The current (I).
Increase in current = more product

The duration of the electrolysis (t).
Increase in time = more product