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AQA A Level Physical Chemistry > 11. Electrode Potentials and Electrochemical Cells > Flashcards

11. Electrode Potentials and Electrochemical Cells Flashcards

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1
Q

What is a redox reaction?

A

In a reaction if an element is oxidised then another must be reduced

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2
Q

What happens when a piece of metal is dipped into a solution of its metal ions?

A

An equilibrium is set up and there is a tendency

  • for the metal to form positive ions and go into solution
  • as well as for the metal ions in solution to gain electrons and form metal
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3
Q

What is a half cell?

A

When a piece of metal is dipped into a solution of its metal ions and an equilibrium is set up

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4
Q

What two things can happen when a piece of metal is dipped into a solution of its metal ions and an equilibrium is set?

A
  • tendency for the metal to form positive ions and go into solution
  • tendency for the metal ions in solution to gain electrons and form metal
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5
Q

What is an electrode also known as?

A

A half cell

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6
Q

For a zinc rod immersed in Zn²⁺ ion solution, what could happen?

A
  • each Zn on rod could deposit two electrons on the rod and move into solution as Zn²⁺ ions (-ve charge on rod)
  • Zn²⁺ ions in solution could accept 2 electrons and move onto rod to become Zn atoms (+ve charge on rod)
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7
Q

What happens when a zinc rod is immersed in Zn²⁺ ion solution and electrons are transferred?

A

A potential difference is set up between the rod and the solution

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8
Q

What is an electrode potential?

A

When a metal is dipped in a solution of its metal ions and a potential difference is set up between the rod and the solution

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9
Q

When will the potential difference of a half cell be greater?

A

When there is a greater tendency to give up electrons (forming ions)

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10
Q

What does an electrode potential indicate?

A

How readily the metal gives up electrons

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11
Q

Does a chemical reaction take place in a lone half cell?

A

No - there is simply a potential difference between the rod and solution

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12
Q

What does the potential difference in a half cell depend on?

A
  • nature of the ions in solution
  • concentration of ions in solution
  • type of electrode used
  • temperature
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13
Q

What is required in a half cell where there is no solid metal involved in the half equation?

A

A metal electrode

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14
Q

Why is a metal electrode required in a half cell where there is no solid metal involved?

A

To allow the flow of electrons

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15
Q

When a half cell involves no solid metal, what metal is usually used as the electrode?

A

Platinum (Pt)

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16
Q

Why is platinum used for metal electrodes?

A

It is unreactive/inert

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17
Q

What are some types of electrodes?

A
  • gas electrodes

* redox electrodes

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18
Q

What is a gas electrode?

A

A gas and a solution of its ions

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19
Q

What is a redox electrode?

A

Two different ions of the same element present in solution

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20
Q

Can the potential of a half cell be measured directly?

A

No

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21
Q

What letter represents the potential of a cell?

A

E

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22
Q

How is the potential of a half cell measured?

A

Connected to another half-cell of known potential, and p.d. between the two half cells is measured

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23
Q

What is an electrochemical cell?

A

When two half cells are combined

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24
Q

What is formed when two half cells are combined?

A

An electrochemical cell

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25
Q

Are the two half reactions that make up a cell in equilibria?

A

Yes

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26
Q

What is the position of equilibria affected by?

A
  • temperature
  • concentration
  • pressure
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27
Q

When can cell potentials only be compared?

A

Under standard conditions

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28
Q

Why must cell potentials be compared under standard conditions?

A

Different conditions could change the position of equilibrium

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29
Q

What are the standard conditions that cell potentials are measured under?

A
  • 298K
  • 1 moldm⁻³ solution of ions
  • 100 kPa
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30
Q

How is the full potential difference (emf) of cells measured?

A

Under zero-current conditions

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31
Q

How are zero-current conditions achieved in order to measure the emf of a cell?

A

By using a high resistance voltmeter

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32
Q

What does S.H.E. stand for?

A

Standard hydrogen electrode

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33
Q

What had to happen before the potential of any half-cells could be measured?

A

A potential had to be assigned to one particular half cell (so the potential of all other electrodes could be measured against it)

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34
Q

What is the potential of the SHE?

A

0V by definition

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35
Q

What is the SHE known as?

A

The primary standard

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36
Q

Why is the SHE known as the primary standard?

A

It is the potential to which all others are compared

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37
Q

What does the standard hydrogen electrode consist of?

A
  • a platinum electrode with hydrogen gas of standard pressure (100 kPa) bubbling over its surface
  • a solution containing H⁺ ions of standard concentration (1 moldm⁻³)
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38
Q

What acids could be used in the standard hydrogen electrode?

A
  • 1 moldm⁻³ HCl/HNO₃

* 0.5 moldm⁻³ H₂SO₄

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39
Q

When a half cell is connected to the SHE, what happens?

A

The p.d. between the two electrodes causes a current to flow between them

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40
Q

When finding the potential of a half cell under test, which side is the standard electrode?

A

The left hand electrode

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41
Q

What can the SHE/primary standard also be known as?

A

The standard electrode

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42
Q

How is the circuit between the SHE and a half cell completed?

A

By using a salt bridge to allow ions to flow from one solution to another

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43
Q

What is a salt bridge usually made from?

A

A piece of filter paper saturated with a solution of an inert electrode, e.g. KNO₃ (aq)

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44
Q

Why can charge be transferred through the salt bridge?

A

The ions in it move

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45
Q

What does the transfer of charge through the salt bridge counter balance?

A

The electron flow

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46
Q

What must be true of the salt used in a salt bridge?

A

Must not react with either of the solutions in the half cells

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47
Q

What is the symbol for the standard electrode potential?

A

E^θ

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48
Q

What is the standard electrode potential?

A

The potential difference between the electrode and the hydrogen electrode under standard conditions

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49
Q

Why are electrode potentials sometimes known as reduction potentials?

A

The equilibria are written with the electrons on the left

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50
Q

What is the electrochemical series?

A

When the electrode potentials are recorded in order of increasing/decreasing potential

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51
Q

What is it called when the reduction potentials are recorded in order of increasing/decreasing potential?

A

The electrochemical series

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52
Q

What do very positive potentials mean?

A

These elements are good at attracting electrons (by taking them from something else which is oxidised)

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53
Q

Which elements in the electrochemical series are the best oxidising agents?

A

Those with very positive potentials

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54
Q

What do very negative potentials mean?

A

These elements are good at giving away electrons (by giving them to something else which is reduced)

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55
Q

What happens to reducing power as the value of potential decreases?

A

It increases

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56
Q

What happens to reducing power as the value of potential increases?

A

It decreases

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57
Q

Which elements in the electrochemical series are the best reducing agents?

A

Those with very negative potentials

58
Q

In terms of electrochemical cells, when do electrons flow?

A

When two different half reactions are connected as part of a complete circuit

59
Q

In an electrochemical cell, which direction do electrons flow?

A

From the more reactive metal to the less reactive metal

60
Q

In terms of electrons, what does it mean if a metal is more reactive?

A

It is more likely to give up electrons (oxidation)

61
Q

In terms of electrons, what does it mean if a metal is less reactive?

A

It is more likely to gain electrons (reduction)

62
Q

Which is the reducing half cell?

A

The one with the most negative value

63
Q

Which is the oxidising half cell?

A

The one with the least negative value

64
Q

What happens to electrons at the anode?

A

The metal loses electrons - so is oxidised to metal ions

65
Q

What happens to electrons at the cathode?

A

The metal ions gain electrons - so is reduced to metal atoms

66
Q

Is the anode oxidised or reduced?

A

Oxidised

67
Q

Why is the anode negative?

A

The metal loses electrons - oxidation, so is oxidised to metal ions

68
Q

Is the cathode oxidised or reduced?

A

Reduced

69
Q

Why is the cathode positive?

A

As electrons are used up - metal ions gain electrons and so it is reduced to metal atoms

70
Q

What must be noted about the charges of the electrodes in the electrochemical cell?

A

The positive and negative electrodes have opposite charges to those during electrolysis

71
Q

What equation is used to find the voltage of a cell?

A

E cell = E₊ - E₋ = E right - E left

72
Q

When calculating the voltage of a cell, what is subtracted from what?

A

The oxidation value (most negative value) is taken away from the reduction value

73
Q

When will a reaction not take place in terms of electrochemical cells?

A

If a stated reaction has a negative electrode potential value

74
Q

What happens if a stated reaction has a negative electrode potential value?

A

The reaction will not take place

75
Q

What does || represent in the convention for writing cells?

A

A salt bridge

76
Q

What represents a salt bridge in the convention for writing cells?

A

||

77
Q

In the convention for writing cells, which electrode is always on the right hand side?

A

The reducing electrode (more negative)

78
Q

In the convention for writing cells, which electrode is always on the left hand side?

A

The oxidising electrode (more positive)

79
Q

In the convention for writing cells, which side does the reducing electrode go on?

A

Right

80
Q

In the convention for writing cells, which side does the oxidising electrode go on?

A

Left

81
Q

When will the oxidising electrode not be written on the left hand side?

A

When measuring potentials against primary or secondary standards, when the standard electrode is always the left hand electrode

82
Q

In the convention for writing cells, which species is put closest to the salt bridge in the middle?

A

The most oxidised species

83
Q

Do state symbols need to be included when writing out cells?

A

Yes

84
Q

In the convention for writing cells, what happens if there are two species in the same state?

A

They are separated using a coma

85
Q

In the convention for writing cells, what represents the boundary between two phases?

A

|

86
Q

In the convention for writing cells, where is the metal electrode written?

A

On the outside

87
Q

In the convention for writing cells, what happens if there is no solid metal in the half equation?

A

Platinum is used and included in cell notation

88
Q

What do batteries contain?

A

Two separate half cells

89
Q

What does the salt bridge in a battery do?

A

Allows ions to flow through without allowing complete mixing of the solutions

90
Q

In a battery, what are the electrodes connected to?

A

The terminals of the battery

91
Q

What is a non-rechargeable cell?

A

When the reactions taking place in the half cells are irreversible and has to be disposed of after their single use

92
Q

What happens to the emf in non-rechargeable cells?

A

Emf drops over time as chemicals are used up

93
Q

What happens, in a non-rechargeable cell, when one or more of the chemicals have been completely used up?

A

The cell is flat and emf is 0V

94
Q

What are non-rechargeable cells useful for?

A

Gadgets which do not use a lot of power e.g. TV remotes and smoke alarms

95
Q

Example of a non-rechargeable cell?

A

Zinc/carbon cells

96
Q

What are the two half equations in a zinc/carbon cell?

A
  • Zn(NH₃)₂²⁺ + 2e⁻ ⇋ Zn + 2NH₃

* 2MnO₂ + 2H⁺ +2e⁻ ⇋ Mn₂O₃ + H₂O

97
Q

What is equation for the overall reaction in a zinc/carbon cell?

(Zn(NH₃)₂²⁺ + 2e⁻ ⇋ Zn + 2NH₃ and

2MnO₂ + 2H⁺ +2e⁻ ⇋ Mn₂O₃ + H₂O)

A

2MnO₂ + 2H⁺ + Zn + 2NH₃ → Mn₂O₃ + H₂O + Zn(NH₃)₂²⁺

98
Q

When is a battery rechargeable?

A

When the reactions taking place in the half cells are reversible

99
Q

Example of a rechargeable cell?

A

Lithium cells

100
Q

What are rechargeable cells useful for?

A

Gadgets which use power frequently e.g. mobile phones and cars

101
Q

How is a rechargeable battery recharged?

A
  • another power supply with a larger emf is connected
  • electrons are forced around the circuit in the opposite direction
  • reversing the spontaneous chemical reaction regenerating the chemicals
102
Q

Examples of the electrodes and electrolyte in a lithium cell?

A
  • positive lithium cobalt oxide electrode
  • negative graphite electrode
  • electrolyte - lithium salt in organic solvent
103
Q

What are the two half equations for a lithium cobalt oxide/graphite electrode lithium cell?

A
  • Li⁺ + CoO₂ + e⁻ ⇋ LiCoO₂

* Li⁺ + e⁻ ⇋ Li

104
Q

In a lithium cell, what is the overall reaction during discharge?

(Li⁺ + CoO₂ + e⁻ ⇋ LiCoO₂ and

Li⁺ + e⁻ ⇋ Li)

A

CoO₂ + Li → LiCoO₂

105
Q

In a lithium cell, what is the overall reaction during re-charge?

(Li⁺ + CoO₂ + e⁻ ⇋ LiCoO₂ and

Li⁺ + e⁻ ⇋ Li)

A

LiCoO₂ → CoO₂ + Li

106
Q

When discharging lithium cells, which direction do electrons flow?

A

Negative to positive electrode

107
Q

When re-charging lithium cells, which direction do electrons flow?

A

Positive to negative electrode (forced through the external circuit)

108
Q

When re-charging lithium cells, what happens to lithium ions?

A

They flow through the electrolyte towards the positive electrode to maintain the balance of charge

109
Q

What type of battery is lead-acid?

A

Rechargeable

110
Q

What type of battery is nickel-cadmium?

A

Rechargeable

111
Q

What is a fuel cell?

A

A cell in which a chemical reaction between a fuel and oxygen is used to create a voltage

112
Q

What are the energy transfers in a fuel cell?

A

Chemical to electrical energy

113
Q

Why do fuel cells not run out or need charging?

A

They have a continuous supply of chemicals to the cell

114
Q

What does a fuel cell consist of?

A

Two half cells connected by a semi-permeable membrane

115
Q

What is the most common fuel cell?

A

The hydrogen fuel cell

116
Q

What does the semi-permeable membrane in a hydrogen fuel cell allow for?

A

OH⁻ ions and water to pass through it but not the oxygen or hydrogen gas

117
Q

What electrodes are used in hydrogen fuel cells?

A

Platinum

118
Q

What do the platinum electrodes act as in hydrogen fuel cells?

A

A catalyst

119
Q

Why do the platinum electrodes act as a catalyst in hydrogen fuel cells?

A

They provide a surface for the reaction

120
Q

What the current available in the hydrogen fuel cell depend on?

A

The rate, hence the advantage of a catalyst

121
Q

As well as a catalyst, what will affect the current generated from a hydrogen fuel cell?

A

Temperature and pressure

122
Q

Does the hydrogen-oxygen fuel cell run in alkaline or acidic conditions?

A

Can run in either

123
Q

For a hydrogen fuel cell run in alkaline conditions, what is the half equation for the negative electrode (anode)?

A

H₂ + 2OH⁻ → 2H₂O + 2e⁻

124
Q

In electrochemical cells, is the anode positive or negative?

A

Negative

125
Q

In electrochemical cells, is the cathode positive or negative?

A

Positive

126
Q

For a hydrogen fuel cell run in alkaline conditions, what is the half equation for the positive electrode (cathode)?

A

O₂ + 2H₂O + 4e⁻ → 4OH⁻

127
Q

For a hydrogen fuel cell run in alkaline conditions, what is the overall equation?

A

2H₂ + 4OH⁻ + O₂ + 2H₂O → 4H₂O + 4OH⁻

128
Q

For a hydrogen fuel cell run in acidic conditions, what is the half equation for the negative electrode (anode)?

A

H₂ → 2H⁺ + 2e⁻

129
Q

For a hydrogen fuel cell run in acidic conditions, what is the half equation for the positive electrode (cathode)?

A

O₂ + 4H⁺ + 4e⁻ → 2H₂O

130
Q

For a hydrogen fuel cell run in acidic conditions, what is the overall equation?

A

2H₂ + 4H⁺ + O₂ + 2H₂O → 2H₂O + 4H⁺

131
Q

For fuel cells working in alkaline conditions, what is used as the electrolyte?

A

A base e.g. KOH

132
Q

For fuel cells working in acidic conditions, what is used as the electrolyte?

A

An acid e.g. H₃PO₄

133
Q

Advantages of cells?

A

Portable source of electrical energy

134
Q

Advantages of non-rechargeable cells?

A

Cheap

135
Q

Advantages of rechargeable cells?

A
  • less waste
  • cheaper in the long run
  • lower environmental impact
136
Q

Advantages of fuel cells?

A
  • water is the only waste product
  • don’t need recharging
  • very efficient
137
Q

Disadvantages of cells?

A

Waste issues

138
Q

Disadvantages of non-rechargeable cells?

A

Waste issues

139
Q

Disadvantages of rechargeable cells?

A

Some waste issues (at end of life)

140
Q

Disadvantages of hydrogen fuel cells?

A
  • need constant supply of fuels
  • hydrogen is flammable and explosive
  • hydrogen is usually made using fossil fuels
  • high cost of fossil fuels

AQA A Level Physical Chemistry (12 decks)

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