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Edexcel Chemistry Unit 5 > 5.3.1 Application of Redox Equilibria > Flashcards

5.3.1 Application of Redox Equilibria Flashcards

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1
Q

What is a redox reaction?

A

A reaction where electrons are transferred.

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2
Q

What is oxidation?

A

A loss of electrons.

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3
Q

What is reduction?

A

Gaining electrons.

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4
Q

What is the term for oxidation and reduction happening simultaneously?

A

Redox.

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5
Q

What is an oxidising agent?

A

Something that accepts electrons and is itself reduced.

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6
Q

What is a reducing agent?

A

Something that donates electrons and itself is oxidised.

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7
Q

How are atoms treated when calculating oxidation numbers?

A

As ions, even if they are covalently bonded.

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8
Q

What is the oxidation number of uncombined elements?

A

0

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9
Q

What is the oxidation number of elements bonded to identical atoms?

A

0

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10
Q

What is the oxidation number of simple monatomic ions?

A

The same as their charge.

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11
Q

What is the overall oxidation number of compound ions?

A

The ion charge.

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12
Q

What is the sum oxidation numbers for a neutral compound?

A

0

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13
Q

What is the oxidation number of combined oxygen?

A

-2 except in peroxides where it is -1 and in fluorides.

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14
Q

What is the oxidation number of combined hydrogen?

A

+1 except in metal hydrides where it’s -1.

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15
Q

What do the roman numerals in a chemical name represent?

A

Oxidation numbers.

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16
Q

What happens to the oxidation state of an atom when an electron is lost?

A

It increases by 1.

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17
Q

What happens to the oxidation state of an atom when an electron is gained?

A

It decreases by 1.

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18
Q

What is the name for when an element is oxidised and reduced at the same time?

A

Disproportionation.

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19
Q

What does an ionic half equation show?

A

Oxidation or reduction.

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20
Q

How can you make a full equation from ionic half equations?

A

Combine an oxidation half equation with a reduction half equation.

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21
Q

When combining ionic half equations, what must balance?

A

Electrons lost/gain and therefore oxidation number.

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22
Q

Give the ions half equation for zinc loses two electrons.

A

Zn(s) → Zn2+(aq) + 2e-

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23
Q

What happens to zinc in this equation?

Zn(s) → Zn2+(aq) + 2e-

A

Zinc is oxidised.

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24
Q

What happens to the oxidation number of zinc in this equation?
Zn(s) → Zn2+(aq) + 2e-

A

It increases by 2.

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25
Q

Balance these two ionic half equations:
Zn(s) → Zn2+(aq) + 2e-
Ag+(aq) + e- → Aq(s)

A

Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)

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26
Q

How can you make a simple electrochemical cell?

A

Dip two different metals in salt solutions of their own ions and connect them with a wire and salt bridge.

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27
Q

What happens in a zinc/copper cell?

A

Zinc loses electrons (is oxidised) in the left cell, realising electrons into the external circuit. Electrons are taken from the external circuit by copper ions, reducing them to copper atoms.

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28
Q

Give an example of two ions of the same element that could be used to create an electrochemical cell.

A

Fe2+ and Fe3+

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29
Q

What are electrode potentials measured against?

A

A standard hydrogen electrode.

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30
Q

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

A

The voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode.

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31
Q

What is a standard hydrogen electrode?

A

Platinum foil submerged in 1mol dm^-3 of H+ ions (acid) with H2 gas bubbled through at 101kPa.

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32
Q

Why is the platinum electrode platinized?

A

To increase its surface area.

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33
Q

Give the equation for the equilibrium setup in the standard hydrogen electrode.

A

H2(g) ⇋ 2H+(aq) + 2e-

34
Q

What is the electrode potential of the standard hydrogen electrode defined as?

A

Zero.

35
Q

Why is the standard hydrogen electrode always shown on the left?

A

So that the measured cell voltage is the standard electrode potential for the right-hand half-cell.

36
Q

What is the equation for Eº cell?

A

Eº cell = Eº RHS - Eº LHS

37
Q

Which half cell goes on the left and why?

A

The half cell with the more negative standard electrode potential. This always gives a positive value.

38
Q

The more reactive a metal is…

A

The more it wants to lose electrons to form a positive ion.

39
Q

Describe what happens to Eº as reactivity of the metals increases?

A

It becomes more negative.

40
Q

The more reactive a non-metal is…

A

The more it wants to gain electrons to form a negative ion.

41
Q

What happens to Eº as the reactivity of non metals increases?

A

It becomes more positive.

42
Q

What can Eº cell be used for?

A

For predicting whether or not a reaction will go.

43
Q

What happens if cell potential is greater than +0.4V?

A

The reaction goes to completion.

44
Q

What happens if cell potential is between 0.0V and +0.4V?

A

The reaction is likely to be reversible.

45
Q

Why might cell potential predictions be wrong?

A

The conditions such as temperature or concentrations may not be standard. A high activation energy may stop it from happening.

46
Q

Why might call potential predictions appear to be wrong?

A

The rate of reaction may be so slow that the reaction doesn’t appear to happen.

47
Q

How is cell potential related to entropy?

A

Eº ∝ ∆S total

48
Q

How is cell potential related to equilibrium constant?

A

Eº ∝ ln K

49
Q

What can redox titrations be used for?

A

To find out exactly how much oxidising agent is needed to exactly react with a quantity of reducing agent.

50
Q

What does a sharp colour change tell you in a redox titration?

A

The reaction has just been completed.

51
Q

Give the two main oxidising agents used in redox titrations.

A

Mangante (VII) ions (MnO4-) in aqueous KMnO4-. These are purple. Dichromate (VI) ions (Cr2O7’2-) in aqueous potassium dichromate (VI), K2CrO7. These are orange.

52
Q

Briefly outline how you can calculate the concentration of a reagent from titration results.

A

Work out the moles of oxidising agent added. Find the ratio between oxidising agent and reagent and use this to work out the moles of reagent per dm^3 of solution.

53
Q

What are Iodine-Sodium Thiosulfate titrations used for?

A

Finding the concentration of an oxidising agent.

54
Q

What happens when you add an oxidising agent to excess potassium iodide?

A

The iodide ions are oxidised to iodine.

55
Q

How would you find out how many moles of iodine have been produced by an oxidising agent?

A

Titrate it with sodium thiosulfate in the burette.

56
Q

How can you see the end point of a thiosulfate titration?

A

When the iodine colour fades to pale yellow, add starch to detect the presence of iodine.

57
Q

Give four sources of error in a thiosulfate titration.

A

The starch needs be be added at the right time and freshly made so it behaves as expected. Copper iodide can make seeing the end point hard. Iodine produced can evaporate from the solution.

58
Q

How does a fuel cell generate electricity?

A

Reacting a fuel with an oxidant.

59
Q

Name a common fuel for a fuel cell.

A

Hydrogen.

60
Q

Name a common oxidant for a fuel cell.

A

Oxygen.

61
Q

What happens at the anode of a fuel cell?

A

Hydrogen is oxidised to H+ ions.

62
Q

What happens at the cathode of a fuel cell?

A

O2 combines with H+ from the anode and e- from the circuit to make H2O.

63
Q

How does a fuel cell force electrons around an external circuit?

A

The polymer electrolyte membrane, (PEM) only allows H+ ions across it, force e- to travel around the circuit to reach the cathode.

64
Q

Roughly what is the voltage produced by a hydrogen fuel cell?

A

0.6V

65
Q

What is the only waste product of a hydrogen fuel cell?

A

Water.

66
Q

Why might not hydrogen fuel cells be as green as they appear?

A

Manufacturing hydrogen fuel is energy intensive or may use non renewable sources of energy.

67
Q

Give two ways of creating hydrogen fuel that aren’t very green and explain why that is so.

A

Reaction steam with natural gas uses fossil fuels and produces CO2. Electrolysis of water often uses fossil fuels as their source of energy since it is energy intensive.

68
Q

How can H2 gas be produced sustainably?

A

Electrolysis powered by a renewable source such as hydroelectric, solar or wind power.

69
Q

Name two fuels other than hydrogen that can be used in a fuel cell.

A

Methanol and ethanol.

70
Q

Give the equation for methanol being oxidised at the anode of a fuel cell.

A

CH3OH + H2O → CO2 + 6e- + 6H+

71
Q

Explain four advantages of using alcohols instead of hydrogen in fuel cells.

A

They have a higher hydrogen density than liquid hydrogen. There is already infrastructure for renewable; alcohols. They don’t need refrigerated storage. Methanol can be made from CO2 so can reduce levels in the atmosphere.

72
Q

Give a disadvantage of using alcohol fuel cells.

A

CO2 is produced.

73
Q

Give two advantages of fuel cell vehicles over petrol or diesel powered vehicles.

A

They produce far less pollution when in use. They are twice as efficient at converting fuel to power.

74
Q

Give three disadvantages of making fuel cells for vehicles.

A

They’re expensive, made using toxic chemicals and have a limited life-span.

75
Q

Outline two uses of fuel cells in addition to powering cars.

A

Power plants in the space shuttle where they also produce drinking water. Breathalysers.

76
Q

Explain how old breathalysers worked.

A

Alcohol was oxidised by potassium dichromate (VII) changing its colour from orange to green. The colour change was measured by a photocell system.

77
Q

How do breathalysers at police stations work and why aren’t they used on the streets.

A

Infrared spectrometry detects the presence and quantity of ethanol. They are very accurate but not portable.

78
Q

How do ethanol fuel cell breathalysers work?

A

The amount of alcohol in the breath is proportional to current produced when breath is fed into the anode.

79
Q

Give two advantages of using fuel cell breathalysers.

A

They are more accurate because they’re less susceptible to give false readings from other substances and they’re portable.

80
Q

How much breath contains the same amount of alcohol as 1cm^3 of blood?

A

2100cm^3

Edexcel Chemistry Unit 5 (5 decks)

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