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A-level Chemistry Oxford AQA > 3.1.9 Electrode Potentials and Electrochemical Cells > Flashcards

3.1.9 Electrode Potentials and Electrochemical Cells Flashcards

1
Q

What are electrochemical cells made out of?

A

Made from 2 different metals dipped in salt solutions of their own ions and connected by wire (external circuit)

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2
Q

What occur within electrochemical cell?

A

Redox reactions occur within it

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3
Q

What do electrochemical cells do?

A

Make electricity

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4
Q

Describe what happens to zinc in a zinc/copper electrochemical cell

A
  1. Zinc loses electrons more easily than copper
  2. Zinc (from zinc electrode) is oxidised to from Zn2+(aq) ions
  3. = releases electrons into external circuit
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5
Q

Describe what happens to copper in a zinc/copper electrochemical cell

A

Same no. of electrons (as zinc releases) are taken from external circuit, reducing Cu2+ ions to copper atoms

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6
Q

How are the 2 solutions connected in electrochemical cells

A

By a salt bridge

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7
Q

What is a salt bridge made out of?

A

Filter paper soaked in KNO3(aq)

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8
Q

What does the salt bridge enable?

A

Enables ions to flow through and balance out the charges

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9
Q

In an electrochemical cell, electrons flow through wire from ____ ____ ____ to ___ _____ ___

A

Electrons flow through wire from more reactive metal to less reactive one

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10
Q

In an electrochemical cell, what is the voltage that the voltmeter between the 2 half-cells measures known as?

A

Cell potential or EMF, known as Ecell

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11
Q

A half-cell can involve solutions of 2 aq ions of same element. Give an example of ions.

A

Fe2+ / Fe3+

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12
Q

Where does the conversion between these Fe2+ and Fe3+ occur?

A

On surface of platinum electrodes

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13
Q

Why do you make electrodes out of platinum?

A

∵ it’s inert

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14
Q

The reactions occurring at the electrodes are ______

A

reversible

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15
Q

Write the half equations for a zinc/copper electrochemical cell

A
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16
Q

In a cell (i.e. 2 half cells joined) which direction each reaction will go in depends on….

A

how easily each metal loses electrons (i.e. how easily it’s oxidised)

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17
Q

How easily metal is oxidised is measured using ____ ____

A

electrode potentials

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18
Q

Metal easy to oxidise = ______ electrode potential

A

very negative electrode potential

(On the LHS)

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19
Q

Metal harder to oxidise = ______ or ____ electrode potential

A

less negative or positive electrode potential

(On RHS)

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20
Q

Write the overall equation for the zinc/copper electrochemical cell

A
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21
Q

What are electrode potentials measured against?

A

Standard Hydrogen Electrodes

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22
Q

Why do we use standard conditions to measure electrode potentials?

A

∵ Cell potential is affected by…

  • Temperature
  • Pressure
  • Concentration
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23
Q

State the standard conditions

A
  • 1.00 mol dm-3
  • 298K
  • 100 kPa
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24
Q

Define standard electrode potenial (E⦵) of half-cells

A

Voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode

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25
Q

State the overall equation in a cell allowing you to find the standard electrode potential of Zn2+/Zn half-cell

A
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26
Q

What is the standard hydrogen electrode made from?

A

Platinum

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27
Q

What solution is used in the half-cell with the standard hydrogen electrode?

A

An acid 1.00 mol dm-3 of H+(aq)

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28
Q

More negative electrode potentials mean that:

  1. The right-hand substances are …
  2. The left-hand substances are …
A

More negative electrode potentials mean that:

  1. The right-hand substances are more easily oxidised
  2. The left-hand substances are more stable
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29
Q

More postive electrode potentials mean that:

  1. The right-hand substances are …
  2. The left-hand substances are …
A

More postive electrode potentials mean that:

  1. The right-hand substances are more stable
  2. The left-hand substances are more easily reduced
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30
Q

State the equation you can use to calculate the standard cell potential (from standard electrode potential values)

A
  • E*⦵cell = E⦵reduced — E⦵oxidised
    i. e. RHS - LHS
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31
Q

Calculate the standard cell potential of a Mg/Fe electrochemical cell

A
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32
Q

State the form for drawing a standard convention

A

Half-cell with more negative potential goes on left

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33
Q

Write the standard convection for a Zn/Cu cell

A
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34
Q

State the standard convection for the standard hydrogen electrode

A

Pt(s) | H2(g) | H+(aq)

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35
Q

State the standard convection for

Ni2+(aq) + 2e- ⇌ Ni(s)

2H+(aq) + 2e- ⇌ H2(g)

A

Ni(s) | Ni2+(aq) || H+(aq) | H2(g) | Pt(s)

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36
Q

State the standard convection for

Fe2+(aq) + 2e- ⇌ Fe(s)

MnO4-(aq) + 8H+(aq) + 5e- ⇌ Mn2+(aq) + 4H2O

A

Fe(s) | Fe2+(aq) || MnO4-(aq), H+(aq), Mn2+(aq) | Pt(s)

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37
Q

State the standard convection for

2H+(aq) + 2e- ⇌ H2(g)

Cr2O72-(aq) + 14H+(aq) + 6e- ⇌ 2Cr3+(aq) + 7H2O

A

Pt(s) | H2(g) | H+(aq) || Cr2O72-(aq), Cr3+(aq), H+(aq) | Pt(s)

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38
Q

Predict whether zinc metal reacts with aqueous copper(II) ions

39
Q

Calculate standard electrode potentials for the electrodes & write an equation to show the reaction that takes place in the cell.

Th4+(aq)/Th(s) Pt(s) | H2(g) | H+(aq) || Th4+(aq) | Th(s)

Ecell = -1.90 V

(standard electrode potential of H = 0)

A

H = 0 V & Th = -1.90 V

4H+ + Th → 2H2 + Th4+

40
Q

Write the balanced equation for

Br2(l) + 2e- ⇌ 2Br-(aq) E⦵= 1.09V

2H+(aq) + 2e- ⇌ H2(g) E⦵= 0V

A

H2 + Br2 → 2H+ + 2Br-

41
Q

What are batteries?

A

Types of electrochemical cell which provide electricity

42
Q

Give an example of non-rechargeable batteries

A

Zinc/carbon cells

43
Q

Describe zinc/carbon cells

A
  • Uses zinc anode & manganese dioxide cathode
  • Carbon is added to cathode to increase conductivity & retain moisture
  • Manganese dioxide takes part in reaction NOT the carbon
44
Q

What is the overall reaction occuring in zinc/carbon cells

A

Zn + 2MnO → ZnO + Mn2O3

45
Q

Name 3 rechargeable batteries

A
  • Lithium batteries
  • Lead-acid batteries
  • Nickel/cadmium batteries
46
Q

Where are lithium batteries found?

A

Found in lots of devices e.g. phones, laptops, cars

47
Q

Name what the electrodes are made out of in a lithium cell?

A

Lithium cobalt oxide (LiCoO2) electrode and a graphite electrode

48
Q

Name what the electrolyte is in a lithium cell?

A

Lithium salt in organic solvent

49
Q

State the half equations for lithium cells

50
Q

Lithium Batteries

State the equation occuring at the negative electrode

A

Li → Li+ + e-

51
Q

Lithium Batteries

State the equation occuring at the positive electrode

A

Li+ + CoO2 + e- → Li+[CoO2]-

52
Q

Lithium Batteries

Calculate the E⦵cell

53
Q

How are rechargeable batteries recharged?

A

A current is supplied to force electrons to flow in the opposite direction around the circuit and reverse the reaction

54
Q

Explain why non-rechargeable batteries cannot be recharged

A

Reactions that occur in non-rechargeable batteries are difficult/impossible to reverse

55
Q

What are lead-acid batteries used for?

A

Used to operate the starter motor of cars

56
Q

Lead-Acid Batteries

State the anode

A

Lead plate

57
Q

Lead-Acid Batteries

State the cathode

A

Lead oxide coated lead plate

58
Q

Lead-Acid Batteries

State the electrolyte

59
Q

Lead-acid Batteries

Write the half-equation occuring at the positive electrode

A

PbO2 + 3H+ + HSO4- + 2e- ⇌ PbSO4 + 2H2O

(V = +1.69)

60
Q

Lead-acid Batteries

Write the half-equation occuring at the negative electrode

A

PbSO4 + H+ + 2e- ⇌ Pb + HSO4-

(V = -0.36)

61
Q

Lead-acid Batteries

Write the overall equation that occurs

A

PbO2 + Pb + 2H2SO4 ⇌ 2PbSO4 + 2H2O

(Vcell = 2.05)

62
Q

Lithium Batteries

Write the overall equation that occurs

A

Li + CoO2 → Li+[CoO2]-

63
Q

What are nickel/cadmium batteries used for?

A

Used to replace zinc-carbon batteries

64
Q

Nickel/cadmium Batteries

Write the half-equation occuring at the anode

A

Cd(OH)2(s) + 2e- ⇌ Cd(s) + 2OH-(aq)

65
Q

Nickel/cadmium Batteries

Write the half-equation occuring at the cathode

A

NiO(OH)(s) + H2O(l) + e- ⇌ Ni(OH)2(s) + OH-(aq)

66
Q

Nickel/cadmium Batteries

Write the overall equation that occurs

A

2NiO(OH)(s) + Cd(s) + 2H2O(l) ⇌ 2Ni(OH)2(s) + Cd(OH)2(s)

(e- flow from Cd → Ni)

67
Q

What do fuel cells do?

A

Generate electricity from hydrogen and oxygen

68
Q

In most cells, where are the chemicals used to generate electricity contained?

A

In electrodes and electrolyte

69
Q

In fuels cells, where are the chemicals used to generate electricity stored?

A

Chemicals stored separately outside cell and fed in when electricity is required

70
Q

Give an example of a fuel cell

A

Alkaline hydrogen fuel cell

71
Q

What are fuel cells used for?

A

Used to power electric vehicles

72
Q

Alkaline hydrogen fuel cell

Hydrogen and oxygen gases are fed into…

A

2 separate platinum-containing electrodes

73
Q

Alkaline hydrogen fuel cell

What does platinum act as?

74
Q

Alkaline hydrogen fuel cell

How are the electrodes separated and what does this allow?

A

Separated by anion-exchange membrane that allows anions (OH-) and water to pass through but NOT hydrogen and oxygen gas

75
Q

Alkaline hydrogen fuel cell

State the electrolyte

A

Aqueous alkaline (KOH) solution

76
Q

Describe what occurs in an alkaline hydrogen fuel cell (i.e. in terms of electrons and ions)

A
  1. Electrons flow from -ve electrode through external circuit to +ve electrode
  2. OH- ions pass through anion-exchange membrane towards -ve electrode
  3. Overall effect: H2 & O2 react to make water
77
Q

Alkaline hydrogen fuel cell

State the reaction that occurs at the negative electrode

78
Q

Alkaline hydrogen fuel cell

State the reaction that occurs at the postive electrode

79
Q

Alkaline hydrogen fuel cell

State the overall reaction that occurs

80
Q

Name and describe 3 advantages of fuel cells

A
  1. In cars, more efficient than internal combustion engine
    • Convert more of their available energy into kinetic energy
    • Internal combustion engine = waste a lot of their energy producing heat
  2. Only waste product is water
    • No toxic chemicals or CO2 emissions
  3. Don’t need to be recharged like batteries
    • As long as H2 + O2 is supplied = continues to produce electricity
81
Q

Name and describe 2 disadvantages of fuel cells

A
  1. Produce a supply hydrogen and oxygen from electrolysis of water which requires electricity
    • Electricity normally generated by burning fossil fuels
    • Whole process ≠ carbon neutral
  2. Hydrogen = highly flammable
    • Storing pressurised H2 requires heavy gas cylinders
82
Q

Describe how people are planning to store hydrogen in the future in fuel cells & the benefit

A
  • Storing H2 absorbed into metals as metal hydrides mean if the cylinder is punctured, H2 escapes slowly
  • Reduced explosion risk
83
Q

Pt | H2SO3(aq), SO42–(aq), H+(aq) || Fe3+(aq), Fe2+(aq) | Pt

Explain why the e.m.f increases when the concentration of Fe3+(aq) ions is increased (2)

A
  • More Fe3+ ions to accept electrons
  • Fe3+/Fe2+ electrode becomes more postive
84
Q
A
  • B - Only Cu(I) undergoes disproportionation
85
Q

Describe a standard hydrogen electrode (4)

A
  • H2 gas
  • 1.0 mol dm–3 HCl/H+
  • At 298K and 100kPa
  • Pt (electrode)
86
Q

Suggest what reactions occur, if any, when hydrogen gas is bubbled into a solution containing a mixture of iron(II) and iron(III) ions. Explain your answer. (2)

A
  • Fe3+ ions reduced to Fe2+
  • Because E(Fe3+/Fe2+) > E(H+/H2)
87
Q

(6)

88
Q

The aluminium used as the electrode is rubbed with sandpaper prior to use. Suggest the reason for this. (1)

A

To remove oxide layer on aluminium

89
Q

Al(s) | Al3+(aq) || H+(aq) | H2(g) | Pt(s)

A simple salt bridge can be prepared by dipping a piece of filter paper into potassium carbonate solution. Explain why such a salt bridge would not be suitable for use in this cell. (2)

A
  • Carbonate ion react with acid (in the SHE)
  • H+ concentration change/cell e.m.f altered
90
Q

Draw labelled diagram of SHE

91
Q

Give one reason, rather than cost, why the platinum electrodes are made by coating a porous ceramic material with platinum rather than by using platinum rods (1)

A

Increases surface area

92
Q

Suggest why the emf of a hydrogen-oxygen fuel cell, operating in acidic conditions, is exactly the same as that of an alkaline fuel cell (1)

A

Overall reaction is the same

93
Q

Part 1: Describe how you would set up a cell

A
  1. Clean piece of copper and zinc using emery paper (fine grade sandpaper)
  2. Degrease metal using some cotton wool and propanone
  3. Place copper into a beaker with about 50 cm3 of CuSO4 solution
  4. Place zinc into a beaker with about 50 cm3 of ZnSO4 solution
  5. Lightly plug one end of the plastic tube with cotton wool and fill the tube with NaCl solution
  6. Plug the free end of the tube with cotton wool which has been soaked in NaCl
    • Join the 2 beakers with inverted U-tube
  7. Connect half-cells by connecting the metals (using the crocodile clips and leads provided) to the voltmeter and read off the voltage
94
Q

Part 2: Describe how you would compare electrode potentials of different metals

A
  1. Clean copper using emery paper/fine grade sandpaper
  2. Connect positive terminal of voltmeter to copper (using a crocodile clip and one of the leads)
  3. Cut piece of filter paper about same area as copper, moisten filter paper with NaCl solution and place on top of copper
  4. Connect second lead to voltmeter & to another metal
  5. Hold metal against filter paper and note voltage reading and sign
  6. Repeat steps 4 and 5 with different metals and record your results

A-level Chemistry Oxford AQA (34 decks)

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