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Chemistry Y13 > Electrochemistry > Flashcards

Electrochemistry Flashcards

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1
Q

Electrochemistry

A
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2
Q

Redox reactions can be studied electrically using…?

A

electrochemical cells

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3
Q

What is an electrochemical cell?

A

an electrochemical cell is a simple system that generates a voltage

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4
Q

Draw a diagram to show an electrochemical cell made up of 2 half-cells connected by a salt bridge

A
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5
Q

What are the metal strips?

A

these are called electrodes (metallic conductors)

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6
Q

What does each beaker contain?

A

each beaker contains a half-cell made up of redox couple

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7
Q

What does a redox couple consist of?

A

a redox couple consist of metal atoms in equilibrium with it’s aqueous solution of ions

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8
Q

In this experiment, which two things need to be connected in order for a reaction to occur?

A

No reaction can occur until the 2 half cells are connected by an external circuit (and a salt bridge)

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9
Q

What is the external circuit?

A

The external circuit is a conducting wire through which electrons can pass from one electrode to the other.

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10
Q

What is the salt bridge?

A

The salt bridge is a strip of filter paper soaked in a solution of soluble ionic compound (e.g potassum nitrate or potassium chloride)

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11
Q

What is the salt bridge usually soaked in?

A

Potassium Nitrate or Potassium Chloride

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12
Q

Electrode Potentials

A
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13
Q

What is the external Circuit?

A

The external circuit is a conducting wire through which electrons can pass from one electrode to the other

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14
Q

What is the salt bridge?

A

the salt bridge is often a strip of filter paper soaked in a solution of a soluble ionic compound(e.g potassium nitrate or potassium chloride) that does not react with the contents of either beaker

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15
Q

What is the role of the salt bridge?

A
  • the salt bridge enables the cell to work by completing the circuit without the 2 solutions mixing
  • the salt bridge allows ions to move between the 2 half-cell compartments
  • it prevents ion saturation at the electrodes that would stop the cell from working
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16
Q

when does a voltmeter measure a reading?

A

when there is a potential difference between the 2 half-cells

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17
Q

What is a voltmeter reading called?

A

A cell potential

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18
Q

Draw and label a cell diagram

A
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19
Q

The cell diagram can be written in a shorthand notation. Write the shorthand notation for RH magnesium electrode and a LH copper electrode. Describe what each thing notation represents?

A
  • Positive electrode = written/drawn on the right hand side
  • single vertical lines = represent phase (state) boundaries
  • double vertical lines = salt bridge
  • metal electrodes = always on the outside
  • oxidation occurs in the Left hand half-cell and reduction occurs in the Right hand half-cell
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20
Q

Describe the flow of electrons according to the cell diagram and the shorthand cell diagram?

A

electrons flow from the left hand half-cell to the right hand half-cell via the external circuit

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21
Q

What type of reaction occurs in the Left-hand half cell and what type of reaction occurs in the right hand half-cell?

A

Left = Oxidation
Right = reduction

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22
Q

Factors that affect the cell potential

A
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23
Q

What are the factors that affect the cell potential?

A
  • cell current
  • cell concentration
  • cell temperature
  • cell pressure
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24
Q

What is the cell current/ what is it measure by?

A

a high resistance digital voltmeter must be used to measure the true cell potential

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25
Q

What does a high resistance voltmeter provide?

A

a high resistance voltmeter provides zero-current conditions

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26
Q

What is the electromotive force (e.m.f)?

A

is the cell potential measured under zero-current conditions (Ecell)

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27
Q

What happens when current if lowing to the voltmeter reading?

A

the cell potential reading drops

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28
Q

What are the standard conditions used when measuring cell potentials?

A
  • Cell concentration = 1 mol dm-3
  • Cell temperature = 298K
  • Cell pressure = 100KPa
  • (Cell current = Zero-current)
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29
Q

The cell pressure is significantly affected by what type of electrodes?

A

Gas electrodes

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30
Q

What are gas electrodes?

A

gas electrodes consist of an inert metal surrounded by gas in equilibrium with a solution of it’s ions

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31
Q

The inert metal is usually what metal?

A

Platinum

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32
Q

What does the inert metal/ platinum act as?

A

Acts as a source or a sink for electrons

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33
Q

What is the most important gas electrode?

A

Hydrogen electrode

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34
Q

What redox couple does the hydrogen electrode correspond to?

A

H+ (aq) / H2(g)

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35
Q

draw the left and right shorthand diagram to show the hydrogen cell in a cell diagram?

A

Left: Pt (s) / H2 (g) / H+ (aq)
Right: H+ (s) / H2 (g) / Pt (s)

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36
Q

What is the oxidation half equation for hydrogen?

A

H2 (g) —-> 2H+ (aq) + 2e-

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37
Q

What is a redox electrode?

A

at a redox electrode, 2 oxidation states of a given element undergo a redox reaction at an inert metal surface

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38
Q

For example the redox couple Fe3+ (aq) / Fe2+, write out the right and left shorthand notation for this Iron gas electrode?

A

Left: Pt (s) / Fe2+ or Fe3+ (aq)
Right : Fe2+ or Fe3+ (aq) / Pt (s)

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39
Q

The Standard Hydrogen Electrode and The Standard Electrode Potentials

A
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40
Q

What are standard potentials?

A

Cell potentials are called standard potentials when they are carried out under standard conditions

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41
Q

What are standard conditions?

A

Standard conditions are:
- 1 mol dm-3
- 298K
- 100 KPa
- Zero current

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42
Q

What can standard potentials for particular electrochemical cell are given by what symbol?

A

E°cell

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43
Q

What is the electrode potential?
What is it a measure of?

A
  • the electrode potential is the potential difference in volts between the metal and a solution of metal ions.
  • it is a measure of the tendency of the metal to lose electrons
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44
Q

What is it not possible to measure the potential of?

A

it is not possible to measure the potential of a single electrode

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45
Q

So how are potentials measured?

A

Instead electrode potentials are listed relative to a standard reference electrode.

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46
Q

What is the Primary Standard electrode reference?

A

the primary standard electrode used is the Standard Hydrogen Electrode (SHE)

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47
Q

What value has the Standard Hydrogen electrode been given?

A

a value of Zero E°= 0.00V by scientists

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48
Q

So all electrode potentials are what to the Standard Hydrogen potential?

A

all other electrode potentials are listed relative to the Standard Hydrogen electrode

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49
Q

What is the standard Hydrogen Electrode?

A

the standard Hydrogen electrode is the hydrogen electrode operating under standard conditions (298K, 100 KPa, 1 mol dm-3)

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50
Q

Describe the conditions surrounding the standard hydrogen electrode?
Draw a diagram to show this.

A

in the hydrogen electrode pure hydrogen gas at 100KPa pressure is bubbled across an inert platinum electrode placed in 1M solution of H+ at 298K

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51
Q

Write the notion used to represent the standard hydrogen electrode?

A

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

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52
Q

Draw the reduction half equation to show this reaction?

A

2H+ (aq) + 2e- ⇌ H2 (g)

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53
Q

The direction of the above reaction depends on what?

A

the direction of the above reaction depends on the other half cell to which the hydrogen electrode is connected

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54
Q

The standard electrode potential of any redox couple can be predicted, how?

A

by connecting the redox couple to the standard hydrogen electrode

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55
Q

Why must these reactions be carried out in standard conditions

A

Standard conditions so that E°cell values for different redox couples can be compared

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56
Q

Describe the flow of electrons in the redox couple below:
Pt (s) / H2(g) / H+ (aq) // Cu2+ (aq) / Cu (s)

A

electrons flow from the hydrogen electrode to the Cu2+ (aq) / Cu (s) redox couple

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57
Q

What is the positive electrode?

A

Cu (s)

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58
Q

So the E °Cell for Cu2+ (aq) / Cu (s) redox couple is also…?
State the E ° Value of this redox couple ?

A
  • Positive -
  • the E ° value for the copper half cell = + 0.34 V
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59
Q

If the copper half cell is replaced by Magnesium half cell Mg(s) /Mg2+ (aq), describe the flow of electrons now?

A

the electrons will flow from the magnesium redox couple towards the hydrogen electrode

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60
Q

So the magnesium electrode will be..?
So the E ° value for the magnesium half cell would be ?

A
  • Negative
  • E ° value = -2.36V
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61
Q

How are standard electrode potential of redox couples determined normally?

A

Many SEP of redox couples have been determined experimentally (using the SHE or secondary standard) and are listed in chemistry data books

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62
Q

What can these values be used to calculate..?

A

standard e.m.f values practically

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63
Q

Why are standard electrode potentials always quoted for the reduction process?

A

by convention

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64
Q

All electrode reactions are..?

A

reversible equilibrium reactions

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65
Q

the direction in which the forward and backward reaction in a half cell will go depends on..?

A

which half-cell it is connected to

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66
Q

What is the standard convection?

A

the standard convection is to place the half-cell with the more positive E° on the right hand side (positive electrode) and the half cell with the more negative E° on the left hand side (negative electrode)

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67
Q

This results in electrons flowing?

A

From the left hand half cell to the right hand half cell

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68
Q

Using this convention an expression can be applied to any pair of redox couples, what is the expression?

A

E°Cell = E°right - E°left

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69
Q

Calculate the standard e.m.f of a cell with Zinc and Copper electrodes in solutions of 1.0M Zinc sulphate and Copper (II) sulphate respectively

Zn2+ (aq) + 2e- ⇌ Zn (s) E° = -0.76V
Cu2+ (aq) + 2e- ⇌ Cu(s) E° = +0.34V

A
  • use expression: E°Cell = E°right - E°left
  • remember the reaction with the more positive value is the right hand half-cell
  • E°Cell = +0.34 - (-0.76)
  • E°Cell = +1.10V
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70
Q

Consider the cell made by connecting
Zn2+ (aq) / Zn (s) to Mg2+ (aq) / Mg (s)

a) write the equation for the reaction that occurs (the spontaneous or feasible reaction) and draw the cell notation
b) calculate the standard e.m.f of the cell

A

a) Two half equations are:
Zn2+ (aq) + 2e- ⇌ Zn (s) E° = -0.76V
Mg2+ (aq) + 2e-⇌ Mg(s) E° = -2.38V
- zinc is the positive electrode because E° = is more positive
- electrons are received here at this redox couple from the external circuit
- so the half cell reaction for this reaction goes FORWARD:
Zn2+ (aq) + 2e- ——–> Zn (s)
- Magnesium is the negative electrode
- electrons leave this electrode and enter the external circuit
- so magnesium half equation = in reverse
Mg (s) ——-> Mg2+ + 2e-
- so overall equation:
Zn2+ (aq) + Mg (s) —-> Zn (s) + Mg2+(aq)
- Cell representation:
- Mg (s) / Mg2+ (aq) // Zn2+ (aq) / Zn (s)

b) = +1.62V

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71
Q

Secondary Standards

A
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72
Q

What are the disadvantages of using the hydrogen electrode?

A

the hydrogen electrode is bulky, difficult to use and a little hazardous

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73
Q

How do we resolve this?

A

so sometimes a different standard is used instead of the hydrogen electrode. Called the secondary standard

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74
Q

Why is using a secondary standard better?

A

secondary standards are more convenient to use than the hydrogen electrode

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75
Q

what is the secondary standard calibrated against?

A

the secondary standard is calibrated against the standard hydrogen electrode and produce equally valid standard potentials

76
Q

What are the two secondary standards?

A
  • Calomel cell and
  • Silver/silver chloride electrode
77
Q

What is the Calomel cell?
Draw the shorthand notation of this cell?

A
  • this is the most commonly used Secondary standard
  • Pt (s) / Hg(l) / Hg2Cl2 (s), KCl (aq)
78
Q

What is the reduction half equation of the calomel cell?

A

Hg2Cl2 (s) + 2e- —> 2Hg (l) + 2Cl - (aq)

79
Q

What is the E° value calibrated against the hydrogen half-cell?

A

E° value = +0.27

80
Q

What is the silver/silver chloride cell?

A

second most commonly used secondary standard
Ag (s) / AgCl (s) / Cl- (aq)

81
Q

What is the reduction half-equation?

A

AgCl (s) + e- —–> Ag (s) + Cl- (aq)

82
Q

What is the E° value calibrated against the Hydrogen half-cell?

A

E° = +0.22

83
Q

What happens when reaction conditions are altered?

Why is there a change in equilibrium positions?

A
  • then a change in equilibrium position occurs and the electrode potential value
  • this is because the reaction between the species in a redox couple is a reversible reaction and involves an equilibrium
84
Q

Consider the standard electrode potential of Zn2+ (aq) / Zn (s) redox couple:

Zn2+ (aq) + 2e- ⇌ Zn (s) E° = -0.76V

a) When is the value (E°) valid?

A

a) only valid in standard conditions (298K, 1.0M solution of Zn ions, 100KPa)

85
Q

What happens when the concentration of the above reaction increases?

A
  • if the concentration of zinc ions is increased, then the electrode potential will become more positive
  • this is because of the reduction of the Zinc Ions is now more likely to occur (Le Chatlier’s Principle)
86
Q

Eventually (non-rechargeable) batteries stop working. Which electrochemical cell was constructed during the early day of batter development?

A

Zn (s) / Zn2+ (aq) // Cu2+ (aq) / Cu (s) = Daniel cell

87
Q

What are the two half equations for this reaction?

A
  • Zn2+ (aq) + 2e- ⇌ Zn (s) E° = -0.76V
  • Cu2+ (aq) + 2e- ⇌ Cu (s) E° = +0.34V
88
Q

What is the overall equation for this recation?

A

Zn(s) + Cu2+ (aq) –> Zn2+ (aq) + Cu (s) E° cell = +1.10V

89
Q

Using Le Chatelier’s Principle, explain what heppens when:
1. [Cu2+] is increased
2. [Cu2+] is decreased
3. [Zn2+] is decreased
4. [Zn2+] is increased

A
  1. If [Cu2+] is increased, E°cell goes up, equilibrium goes to the right
  2. [Cu2+] is decreased, E° goes down, equilibrium goes to the left
  3. E° goes up, equilibrium goes to the right
  4. E° goes down, equilibrium goes to the left
90
Q

Describe what happen overtime when the [Zn2+] increases and [Cu2+] decreases?

A
  • In the Daniel cell, over time, the [Zn2+] increases, E° goes down
  • and [Cu2+] decreases, E° goes down
  • As a result the E°cell value will gradually decrease to zero volts
91
Q

The Electrochemical Series

A
92
Q

What is the Electrochemical series?

A

The electrochemical series is a list of standard electrode potentials (E°) placed in order according to the value of the reduction process

93
Q

Describe the main trend in the electrochemical series?

A

the following series places the most positive electrode potential at the top and the other electrode potentials are placed in descending order of electrode potentials down the series

94
Q

Why do we need to know how the ordering works?

A

In the exam, the order is often reversed

95
Q

What equation can we use to think about these reactions?

A

The reaction can be thought of as an oxidising agent accepting electrons to form a reducing agent.
Ox + e- —-> Red

96
Q

The strongest oxidising agents accept…?

A

electrons more easily

97
Q

The strongest reducing agents lose…?

A

electrons more easily

98
Q

Using the Electrochemical Series and Predicting the Direction of Simple Redox Reactions

A
99
Q

When is a reaction spontaneous or feasible?

A
  • when the E°cell is positive, so E°Cell is greater than 0
  • If reduction occurs at the right hand cell
  • the more positive the cell potential (E°cell), the further the equilibrium lies to the right
100
Q

The electrochemical series half-equation with the more positive E° value oxidises (reverses), the one….?

A

with a lower value

101
Q

What direction does this equation go then in a spontaneous/feasible reaction?

A

this means that the half equation with the more positive E° value always goes FORWARD (as written in the series)

102
Q

What happens to the other half equation?

A

The other half equation get’s reversed

103
Q

State the direction of the reaction, when the E° value is more positive?

A

More positive E° values - Spontaneous direction is FORWARDS

104
Q

State the direction of the reaction, when the E° value is more negative?

A

More negative E° Value - Spontaneous direction is BACKWARDS

105
Q

When a reaction is feasible, Gibbs free energy is..?

A

Negative -ΔG° < 0

106
Q
  • ΔG Corresponds to ..?
A

Positive +E°

107
Q

Consider the cell made by Connecting
Ag+(aq)/ Ag(s) to Mg2+ (aq)/ Mg (s)

The 2 half-equations are:
Ag+ (aq) + e- ⇌ Ag(s) E°= +0.80V
Mg2+ (aq) + 2e- ⇌ Mg (s) E°= -2.38V

a) Write the equation for the spontaneous/ feasible reaction that occurs?
b) Draw the cell representation and calculate the standard e.m.f value of the cell, and state is the reaction is feasible?

A

a) Silver is the more positive electrode because the E° value for Ag+ (aq)/ Ag (s) redox couple is the ore positive

So the half cell for the reaction goes forwards as written:
Ag+ (aq) + e- —–> Ag (s)

Magnesium is the negative electrode
So magnesium half equation goes in reverse:
Mg (s) —–> Mg2+ (aq) + 2e-

So the spontaneous/ feasible reaction that occurs is:
2ag+ (aq) + Mg (s) —-> 2Ag (s) + Mg2+

B) Cell representation is obtained by placing the more positive E° value on the right:
Mg (s) / Mg2+ (aq) //Ag+ (aq) / Ag (s)

Standard e.m.f = +0.80 - (-2.38)
°e..m.f = +3.18V = reaction is feasible

108
Q

a) Using the electrochemical series to predict is Zinc will reduce V3+ (aq) to V2+

The 2 half-equations are:
V3+ (aq) + e- ⇌ V2+(aq) E°= -0.26V
Zn2+ (aq) + 2e- ⇌ Zn (s) E°= -0.76V

b) Draw the cell representation and calculate the standard e.m.f value of the cell?

A

V3+/V2+ is the positive electrode because the E° value for this redox couple is more positive
So half equation for this couple goes forwards as written:
V3+ (aq) + e- —-> V2+ (aq)

Zinc is the more negative electrode
So the half-equation goes in reverse:
Zn (s) —–> Zn2+ (aq) + 2e-

overall equation:
2V3+ (aq) + Zn (s) —-> 2V2+ (aq) + Zn2+

b) Cell representation:
Zn (s) / Zn2+ (aq) //V3+ (aq) , V2+ / Pt (s)

Standard e.m.f = -0.26-(-0.76)
°e.m.f = + 0.50V
The value is positive therefore the reaction is feasible. Zinc Ions can reduce V3+ ions to V2+ ions

109
Q

E3) Using the electrochemical series to predict is Fluorine can oxidise Fe2+ (aq) to Fe3+

The 2 half-equations are:
F2+ (g) + 2e- ⇌ 2F-(aq) E°= +2.87V
Fe3+ (aq) + e- ⇌ Fe2+ (aq) E°= +0.77V

b) Draw the cell representation and calculate the standard e.m.f value of the cell?

A

F2/F- is the positive electrode as the Ea) Using the electrochemical series to predict is Zinc will reduce V3+ (aq) to V2+

The 2 half-equations are:
V3+ (aq) + e- ⇌ V2+(s) E°= -0.26V
Zn2+ (aq) + 2e- ⇌ Zn (s) E°= -0.76V

b) Draw the cell representation and calculate the standard e.m.f value of the cell? value s more positive
So the half-cell equation for this redox couple goes forwards as written:
F2 (g) + 2e- —> 2F- (aq)

Fe3+/Fe2+ is the negative electrode
So this half equation goes in reverse:
Fe2+ (aq) ——> Fe3+ (aq) + e-

Overall equation:
F2 (g) + 2Fe2+ (aq) —-> 2F- (aq) + 2Fe3+ (aq)

Cell representation:
Pt (s) / Fe2+ (aq), Fe3+ (aq) // F2 (g)/ F- (aq) / Pt (s)

Standard e.m.f = + 2.87V - (+0.77)
Standard e.m.f = +2.10V

E° cell is positive therefore the reaction is feasible.
Fluorine can oxidise Fe2+ ions to Fe3+ ions

110
Q

Calculate the e.m.f of the following electrochemical cell and write an equation for the overall cell reaction:
Al (s) / Al3+ (aq) // V3+ (aq), V2+(aq) / Pt (s)

The 2 half equations are:
V3+ (aq) + e- ⇌ V2+ (aq) E°= -0.26V
Al3+ (aq) + 3e- ⇌ Al (s) (s) E°= -1.66V

A

V3+/V2+ is the positive electrode because E° value is more positive
So the half-cell equation for the reaction goes forwards as written:
V3+ (aq) + e- —-> V2+ (aq)

Al (s)/Al3+ is the negative electrode
So this half equation is in reverse:
Al (s) —–> Al3+ (aq)+ 3e-

Overall equation:
Al (s) + 3V3+ (aq) —> Al3+ (aq) + 3V2+ (aq)

Standard e.m.f = +1.40V
Therefore the reaction is feasible

111
Q

Predicting Redox Equations

A
112
Q

What values can be used to predict whether or not a redox reaction occurs (is feasible) when reactants are mixed together?

A

Standard electrode potentials

113
Q

How should we go about predicting redox equations when reactants are mixed together?

A
  • Identify Appropriate Half Equations
    -redox involves both oxidation+ reduction
    -So one of the half-equations must show electron gain (Reduction) and other electron loss (oxidation)
  • Using the E° Values
    -E° Values for two half equation are identified (one forward and one reversed)
    -E° Values are then added together
    -Note that the values are not adjusted in terms of the different number of electrons involved
    -if overall E°reaction = is positive reaction is feasible and reaction will occur spontaneously
    -E° Reaction is negative then a reaction can not occur, thermodynamically impossible under standard conditions
    -Not that Half equations are equivalent to the half-cells discussed previously and E°reaction is the same as E°cell
114
Q

Example 1.
Predict the products if any, when Cu2+ (aq) ions are mixed with V2+ (aq) ions. Use the E° values to explain your answer.

The 2 Half equations are:
V3+ (aq) + e- ⇌ V2+ (aq) E°= -0.26V
Cu2+ (aq) + 2e- ⇌ Cu (s) E°= +0.24V

A

Cu2+ ions are the reactants and must therefore be on the left-hand side of the equation (written as forwards):
Cu2+ (aq) + 2e- ⇌ Cu (s) E° = +0.24

V2+ ions are also reactants and must therefore be on the left-hand side of the equation.
This involves reversing the half-equation and changing the E° sign:
V3+ (aq) + e- ⇌ V2+ (aq) E°= + 0.26V

We can now add together the 2 E° values to get an E° Reaction value
E° reaction = +0.26 + (+0.34)
E° r = +0.60V
The reaction is therefore feasible (a reaction occurs) and therefore products will be formed

Overall equation:
2V2+ (aq) + Cu2+ (aq) —> 2V3+ (aq) + Cu (s)
Therefore the products are V3+ (aq) and Cu (s)

115
Q

Example 2.
Predict the products if any, when Zn metal is mixed with Mg2+ (aq) ions. Use E° Values to explain your answer

The 2 Half equations are:
Mg2+ (aq) + 2e- ⇌ Mg (s) E°= -2.36V
Zn2+ (aq) + 2e- ⇌ Zn (s) E°= -0.76V

A

Mg2+ ions are the reactants, so therefore must be on the left hand side of the equation (forwards as written):
Mg2+ (aq) + 2e- ⇌ Mg (s) E°= -2.36V

Zinc metal is also a reactant and must therefore be on the left-hand side of the half-equation.
This involves reversing the half-equation and changing the E° Sign:
Zn (s) —–> Zn2+ + 2e- E°= -0.76V

We can now add the two E°values to get a E°reaction value.
E°r = -1.60V
The reaction is therefore not feasible and therefore no products will be formed

116
Q

ELECTROCHEMICAL CELLS

A
117
Q

What are adapted chemical cells used for?

A

Cn be used as commercial source of electrical energy

118
Q

These electrochemical cells can be divided into 3 main groups, what are they?

A
  • primary cells
    -secondary cells
    -Fuel cells
119
Q

What are electrochemical cells that provide a convenient source of electrical energy commonly called?

A

batteries

120
Q

Where does the term battery originate from?

A

The term battery originates from an array of more than 1 cell joined together in series

121
Q

What is a primary cell?

A

a primary cell is irreversible and is not intended to be recharged by an electric current (non-rechargeable)

122
Q

what is a secondary cell?

A

a secondary cell is reversible and is specifically designed to be recharged by an electric current (rechargeable)

123
Q

What is a Fuel cell?

A

A fuel cell generates electricity from the continuous oxidation of an external source of fuel

124
Q

Fuel cells can also be considered what type of cells?

A

to be primary cells

125
Q

Why is this?

A

since they do not need to be electrically recharged

126
Q

What are fuel cells charged by?

A

Fuel cells are charged by the continuous supply of fuel

127
Q

Primary Cells

A
128
Q

What is a primary cell?

A

a primary cell is irreversible and not intended to be recharged by an electrical current

129
Q

What is the single function of a primary cell?

A

The single function of a primary cell is to provide current to an external circuit while discharging (galvanic action)
- once discharged, primary cells are discarded

130
Q

Explain what is meant by the term disposable?

A

no longer useful and needs to get ridden of

131
Q

What is a Galvanic Cell?

A

a galvanic cell is one with a positive e.m.f in which the spontaneous forward cell reaction can be used to provide an electrical current to an external circuit

132
Q

What are some features of a primary cell?

A
  • disposable
  • long-life / high power
133
Q

Secondary Cells

A
134
Q

what is a secondary cell?

A

a secondary cell si reversible and is designed to be recharged by an electricla current (rechargeable)

135
Q

secondary cells combine two opposing functions of all reversible cells what are they?

A
  • They provide current to an external circuit while discharging (galvanic action) –> chemical into electrical energy
  • They use current form an external circuit while charging (electrolytic action) – electrical into chemical energy
136
Q

What is an electrolytic cell?

A

An electrolytic cell is one with a negative e.m.f, which makes the forward reaction impossible, but which is then made possible by the flow of electric current from an external circuit

137
Q

In order for the system to be completely reversible, what is important that needs to be maintained?

A

It is important that the products resulting from both Galvanic and electrolytic action are not dispersed in the cell but (being insoluble) remain attached to the cell electrodes

138
Q

What is an example of a rechargeable cell?

A

Lithium ion cell

139
Q

what examples of technology contain Lithium ion cell batteries?

A

mobile phones, tablets and laptops contain Lithium ion batteries - these are rechargeable batteries

140
Q

What type of reaction do these cells use?

A

they use reversible reactions

141
Q

What is a Lithium ion cell made up of?

A
  • one type of Lithium ion cell is made u of LiCoO2 electrode and a graphite electrode
  • the electrolyte is a lithium salt in an organic solvent
142
Q

What are the 2 half equations for the reaction?

A
  • Li+ (aq) + CoO2 (s) + e- ⇌ Li+[CoO2]- (s) E° = +0.56V
  • Li+ (aq) e- ⇌ Li (s) E° = -3.04V
143
Q

Describe using equation what occurs when the battery supplies power (discharging)?

A
  • At the positive electrode = - Li+ (aq) + CoO2 (S) + e- ⇌ Li+[CoO2]- (s)
  • At the Negative electrode = Li+ (aq) e- ⇌ Li (s) E° = -3.04V

-Overall reaction = Li (s) + CoO2 (S) –> Li+ [CoO2]- (s)

144
Q

Reverse reactions occur on…?
What does this do?

A

Charging - replenishing the battery

145
Q

What are two advantages of the Lithium ion cell?

A
  • the EMF value is really high (due to the lithium having the highest negative standard electrode potential
  • Another advantage of the lithium ion cell is that Lithium is ver light (great for mobile devices)
146
Q

What are the features of some common secondary cells?

A
147
Q

Fuel cells

A
148
Q

What are Fuel cells?

A

Fuel cells use a supply of hydrogen or organic fuel, together with a supply of oxygen, to provide a source of electrical power

149
Q

The simplest fuel cell uses energy from the reaction of … to provide electrical power

A

Hydrogen and Oxygen

150
Q

What is the percentage of chemical energy of combustion can be directly converted into electrical energy by Fuel cells?

A

85%

151
Q

FUEL CELLS

A
152
Q

why can fuel cells also be considered as primary cells?

A

Because they do not need to be recharged

153
Q

What is a fuel cell?

A

Fuel cells use a supply of hydrogen or organic fuel, together with a supply of oxygen, to provide a source of electrical power.

154
Q

How does a fuel cells make electricity?

A

A fuel cell makes elctricity by introducing fuel (on the anode catalyst side) and an oxidant (on the cathode catalyst side), which react together in the presence of an electrolyte

155
Q

Describe the flow of reactants and products in a fuel cell?
This means that..?

A

The reactants flow in and the products flow out, while the electrolyte stays the same. This means that the fuel cells can flow continuously as long as the various flows are maintained

156
Q

Why are fuel cells very reliable?

A

Because there are no moving parts

157
Q

What are fuel cells used in combination with and why?

A

Fuel cells cannot store energy like batteries but they are sometimes used in combination with batteries

158
Q

What can these batteries do?

A

These batteries can electrolyse the products during charging

159
Q

What is the difference between fuel cells and batteries?

A

Fuel cells consume reactants that must be replenished, whereas batteries store and release energy in a closed system

160
Q

What is the simplest of the fuel cells?

A

The Hydrogen-Oxygen fuel cell

161
Q

SO what is the role of the Hydrogen and Oxygen fuel cell?

A

Uses energy from the reaction of oxygen and hydrogen to provide electrical power

162
Q

What is used as fuel and what is used an an oxidant in this reaction/fuel cell?

A

Hydrogen is used as the fuel (on the anode catalyst)
Oxygen is used as the Oxidant (on the cathode catalyst side)

163
Q

What must this fuel cell occur in?

A

An electrolyte

164
Q

What catalyst is typically used in a hydrogen and oxygen fuel cell?

A

Usually a platinum- group metal or a alloy

165
Q

Why are the platinum electrodes Inert and porous?

A

to allow the passage of reactant and product gases

166
Q

The 2 platinum electrodes are immersed in an electrolyte. what could this electrolyte be?

A
  • Acidic (e.g Phosphorus acid)
  • Alkaline (e.g commercial cells use hot aqueous potassium hydroxide
167
Q

Describe the effect of platinum on the rate of reaction?

A

Platinum increases the rate of the electrode reactions (faster reactions provide higher current)

168
Q

How does the ACIDIC hydrogen-Oxygen fuel cell work?

A

-Hydrogen diffuses to the anode platinum catalyst - here it splits up to make H+ ions (protons) and electrons
- The hydrogen ions are conducted through the electrolyte to the cathode
- the electrons travel through the external circuit
- At the cathode platinum catalyst surface, oxygen molecules react with the electrons (that have travelled through the external circuit) and the hydrogen ions to form water
- Any excess gases are recycled, so water is the only waste product

169
Q

How does the ALKALINE Hydrogen-Oxygen Fuel cell work?

A
  • at the anode, hydrogen reacts with hydroxide ions(that have been conducted through the electrolyte), water and electrons are produced - the electrons enter the external circuit
  • At the cathode, oxygen reacts with the electrons arriving from the external circuit to make hydroxide ions (that are conducted through the electrolyte
170
Q

What are Acidic condition half-equations?

A

Red: O2 (g) + 4H+ () + 4e- —> 2H2O (l)
Ox: H2 (g) –> 2H+ (aq) + e-
Overall: O2 (g) + 2H2 (g) —> 2H2O (l)

171
Q

What are alkaline half-equations

A

Red: O2 (g) + 2H2O (l) + 4e- —>4OH- (aq)
Ox: H2 (g) + 2OH- (aq) —> 2H2O (lo) + 2e-
Overall: 2H2 (g) + O2 (g) –> 2H2O (l)

172
Q

Why is the EMF value of the two equations the same?

A

This is because the overall reaction is the same

173
Q

Give the cell representation of an acidic Hydrogen-oxygen fuel cell?

A

Pt (s) / H2 (g) / H+ (aq) // O2 (g)/ H20 (l), H+ (aq)/ Pt (s)

174
Q

Give the cell representation of the alkaline hydrogen-oxygen fuel cell

A

Pt (s)/ H2 (g)/ OH- (aq), H2O (l) // O2 (g) / H2O (l), OH- (aq) / Pt (s)

175
Q

What are the reaction conditions of the hydrogen-oxygen fuel cell and explain what happens if we do not control these?

A
  • Low electric current is produced - therefore the rate of flow of electrons is low
  • the cell operates at 200’ - this increased temperature increases the overall rate of reaction and hence the available current
  • however, since the reaction is highly exothermic, increases the temperature reduces the cell e.m.f
  • therefore an increased pressure of 40 bar is used to compensate for the increased temperature
176
Q

How is Hydrogen and oxygen obtained?

A

easily and cheaply obtained from the air

177
Q

How is most hydrogen produced?

A
  • Most hydrogen needs to be made.
  • made by the reaction of steam and methane
  • CH4 (g) + H20 (g) —-> CO (g) + 3H2 (g)
  • CO (g) + H2O (g) —–> CO2 (g) + H2
    Overall equation: CH4 (g) + 2H2O (g) —> CO2 (g) + 4H2 (g)
178
Q

Why is this reaction of making hydrogen not really good?

A

Both reactions demand a high input of energy and the methane leaves behind a large carbon footprint (Co2 is made)

179
Q

Why is the Hydrogen-fuel cell not considered green?

A

appear to be green as it makes water as the only waste product.

180
Q

How can Hydrogen be made ‘green’ or ‘Cleanly’?

A
  • but can be considered green if the hydrogen is manufactured by a means that does not itself produce carbon dioxide
  • Hydrogen can be made ‘cleanly’ by the electrolysis of acidified water
181
Q

Where does energy to make hydrogen green/electrolysis come from, what is the downside of this?

A

However the energy required for electrolysis must come from carbon neutral sources e.g solar. And solar is expensive.

182
Q

ORGANIC FUEL CELLS

A
183
Q

What is an organic fuel cell?

A

In organic fuel cells, an organic fuel such as ethanol mixes with water (on the anode catalyst) and reacts with oxygen as the oxidant (on the cathode side pf the catalyst)

184
Q

Write equations to show what happens at the anode and the cathode, as well as an overall equation

A

Anode: c2h6+ 3h20 –> 12H+ + 12e- + 2CO2
Cathode O2+ 4H+ + 4e- —> 2H2O
overall: c2h6 + o2 –> 2co2 + 3h2o

185
Q

DONE!!!

A

Chemistry Y13 (13 decks)

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