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)
2
Q
What occur within electrochemical cell?
A
Redox reactions occur within it
3
Q
What do electrochemical cells do?
A
Make electricity
4
Q
Describe what happens to zinc in a zinc/copper electrochemical cell
A
- Zinc loses electrons more easily than copper
- Zinc (from zinc electrode) is oxidised to from Zn2+(aq) ions
- = releases electrons into external circuit
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
6
Q
How are the 2 solutions connected in electrochemical cells
A
By a salt bridge
7
Q
What is a salt bridge made out of?
A
Filter paper soaked in KNO3(aq)
8
Q
What does the salt bridge enable?
A
Enables ions to flow through and balance out the charges
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
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
11
Q
A half-cell can involve solutions of 2 aq ions of same element. Give an example of ions.
A
Fe2+ / Fe3+
12
Q
Where does the conversion between these Fe2+ and Fe3+ occur?
A
On surface of platinum electrodes
13
Q
Why do you make electrodes out of platinum?
A
∵ it’s inert
14
Q
The reactions occurring at the electrodes are ______
A
reversible
15
Q
Write the half equations for a zinc/copper electrochemical cell
A
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)
17
Q
How easily metal is oxidised is measured using ____ ____
A
electrode potentials
18
Q
Metal easy to oxidise = ______ electrode potential
A
very negative electrode potential
(On the LHS)
19
Q
Metal harder to oxidise = ______ or ____ electrode potential
A
less negative or positive electrode potential
(On RHS)
20
Q
Write the overall equation for the zinc/copper electrochemical cell
A
21
Q
What are electrode potentials measured against?
A
Standard Hydrogen Electrodes
22
Q
Why do we use standard conditions to measure electrode potentials?
A
∵ Cell potential is affected by…
- Temperature
- Pressure
- Concentration
23
Q
State the standard conditions
A
- 1.00 mol dm-3
- 298K
- 100 kPa
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
25
State the overall equation in a cell allowing you to find the standard electrode potential of Zn2+/Zn half-cell
26
What is the standard hydrogen electrode made from?
Platinum
27
What solution is used in the half-cell with the standard hydrogen electrode?
An acid 1.00 mol dm-3 of H+(aq)
28
More negative electrode potentials mean that:
1. The right-hand substances are ...
2. The left-hand substances are ...
More negative electrode potentials mean that:
1. The right-hand substances are _more easily oxidised_
2. The left-hand substances are _more stable_
29
More postive electrode potentials mean that:
1. The right-hand substances are ...
2. The left-hand substances are ...
More postive electrode potentials mean that:
1. The right-hand substances are _more stable_
2. The left-hand substances are _more easily reduced_
30
State the equation you can use to calculate the standard cell potential (from standard electrode potential values)
* E*⦵cell = *E*⦵reduced — *E*⦵oxidised
i. e. RHS - LHS
31
Calculate the standard cell potential of a Mg/Fe electrochemical cell
32
State the form for drawing a standard convention
Half-cell with more negative potential goes on left
33
Write the standard convection for a Zn/Cu cell
34
State the standard convection for the standard hydrogen electrode
Pt(s) | H2(g) | H+(aq)
35
State the standard convection for
Ni2+(aq) + 2e- ⇌ Ni(s)
2H+(aq) + 2e- ⇌ H2(g)
Ni(s) | Ni2+(aq) || H+(aq) | H2(g) | Pt(s)
36
State the standard convection for
Fe2+(aq) + 2e- ⇌ Fe(s)
MnO4-(aq) + 8H+(aq) + 5e- ⇌ Mn2+(aq) + 4H2O
Fe(s) | Fe2+(aq) || MnO4-(aq), H+(aq), Mn2+(aq) | Pt(s)
37
State the standard convection for
2H+(aq) + 2e- ⇌ H2(g)
Cr2O72-(aq) + 14H+(aq) + 6e- ⇌ 2Cr3+(aq) + 7H2O
Pt(s) | H2(g) | H+(aq) || Cr2O72-(aq), Cr3+(aq), H+(aq) | Pt(s)
38
Predict whether zinc metal reacts with aqueous copper(II) ions
39
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)
H = 0 V & Th = -1.90 V
4H+ + Th → 2H2 + Th4+
40
Write the balanced equation for
Br2(l) + 2e- ⇌ 2Br-(aq) *E*⦵= 1.09V
2H+(aq) + 2e- ⇌ H2(g) *E*⦵= 0V
H2 + Br2 → 2H+ + 2Br-
41
What are batteries?
Types of electrochemical cell which provide electricity
42
Give an example of non-rechargeable batteries
Zinc/carbon cells
43
Describe zinc/carbon cells
* 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
What is the overall reaction occuring in zinc/carbon cells
Zn + 2MnO → ZnO + Mn2O3
45
Name 3 rechargeable batteries
* Lithium batteries
* Lead-acid batteries
* Nickel/cadmium batteries
46
Where are lithium batteries found?
Found in lots of devices e.g. phones, laptops, cars
47
Name what the electrodes are made out of in a lithium cell?
Lithium cobalt oxide (LiCoO2) electrode and a graphite electrode
48
Name what the electrolyte is in a lithium cell?
Lithium salt in organic solvent
49
State the half equations for lithium cells
50
**Lithium Batteries**
State the equation occuring at the negative electrode
Li → Li+ + e-
51
**Lithium Batteries**
State the equation occuring at the positive electrode
Li+ + CoO2 + e- → Li+[CoO2]-
52
**Lithium Batteries**
Calculate the *E*⦵cell
+3.60 V
53
How are rechargeable batteries recharged?
A current is supplied to force electrons to flow in the opposite direction around the circuit and reverse the reaction
54
Explain why non-rechargeable batteries cannot be recharged
Reactions that occur in non-rechargeable batteries are difficult/impossible to reverse
55
What are lead-acid batteries used for?
Used to operate the starter motor of cars
56
Lead-Acid Batteries
State the anode
Lead plate
57
Lead-Acid Batteries
State the cathode
Lead oxide coated lead plate
58
Lead-Acid Batteries
State the electrolyte
H2SO4
59
Lead-acid Batteries
Write the half-equation occuring at the positive electrode
PbO2 + 3H+ + HSO4- + 2e- ⇌ PbSO4 + 2H2O
| (V = +1.69)
60
Lead-acid Batteries
Write the half-equation occuring at the negative electrode
PbSO4 + H+ + 2e- ⇌ Pb + HSO4-
| (V = -0.36)
61
Lead-acid Batteries
Write the overall equation that occurs
PbO2 + Pb + 2H2SO4 ⇌ 2PbSO4 + 2H2O
| (Vcell = 2.05)
62
Lithium Batteries
Write the overall equation that occurs
Li + CoO2 → Li+[CoO2]-
63
What are nickel/cadmium batteries used for?
Used to replace zinc-carbon batteries
64
Nickel/cadmium Batteries
Write the half-equation occuring at the anode
Cd(OH)2(s) + 2e- ⇌ Cd(s) + 2OH-(aq)
65
Nickel/cadmium Batteries
Write the half-equation occuring at the cathode
NiO(OH)(s) + H2O(l) + e- ⇌ Ni(OH)2(s) + OH-(aq)
66
Nickel/cadmium Batteries
Write the overall equation that occurs
2NiO(OH)(s) + Cd(s) + 2H2O(l) ⇌ 2Ni(OH)2(s) + Cd(OH)2(s)
(e- flow from Cd → Ni)
67
What do fuel cells do?
Generate electricity from hydrogen and oxygen
68
In most cells, where are the chemicals used to generate electricity contained?
In electrodes and electrolyte
69
In fuels cells, where are the chemicals used to generate electricity stored?
Chemicals stored separately outside cell and fed in when electricity is required
70
Give an example of a fuel cell
Alkaline hydrogen fuel cell
71
What are fuel cells used for?
Used to power electric vehicles
72
Alkaline hydrogen fuel cell
Hydrogen and oxygen gases are fed into...
2 separate platinum-containing electrodes
73
Alkaline hydrogen fuel cell
What does platinum act as?
Catalyst
74
Alkaline hydrogen fuel cell
How are the electrodes separated and what does this allow?
Separated by anion-exchange membrane that allows anions (OH-) and water to pass through but NOT hydrogen and oxygen gas
75
Alkaline hydrogen fuel cell
State the electrolyte
Aqueous alkaline (KOH) solution
76
Describe what occurs in an alkaline hydrogen fuel cell (i.e. in terms of electrons and ions)
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
Alkaline hydrogen fuel cell
State the reaction that occurs at the negative electrode
78
Alkaline hydrogen fuel cell
State the reaction that occurs at the postive electrode
79
Alkaline hydrogen fuel cell
State the overall reaction that occurs
80
Name and describe 3 advantages of fuel cells
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
Name and describe 2 disadvantages of fuel cells
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
Describe how people are planning to store hydrogen in the future in fuel cells & the benefit
* Storing H2 absorbed into metals as metal hydrides mean if the cylinder is punctured, H2 escapes slowly
* Reduced explosion risk
83
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)
* More Fe3+ ions to accept electrons
* Fe3+/Fe2+ electrode becomes more postive
84
* B - Only Cu(I) undergoes disproportionation
85
Describe a standard hydrogen electrode (4)
* H2 gas
* 1.0 mol dm–3 HCl/H+
* At 298K and 100kPa
* Pt (electrode)
86
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)
* Fe3+ ions reduced to Fe2+
* Because E(Fe3+/Fe2+) \> E(H+/H2)
87
(6)
60.5%
88
The aluminium used as the electrode is rubbed with sandpaper prior to use. Suggest the reason for this. (1)
To remove oxide layer on aluminium
89
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)
* Carbonate ion react with acid (in the SHE)
* H+ concentration change/cell e.m.f altered
90
Draw labelled diagram of SHE
91
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)
Increases surface area
92
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)
Overall reaction is the same
93
Part 1: Describe how you would set up a cell
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
Part 2: Describe how you would compare electrode potentials of different metals
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
