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Physical chemistry. > 3.1.11 Electrode potentials and electrochemical cells > Flashcards

3.1.11 Electrode potentials and electrochemical cells Flashcards

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1
Q

3.1.11 Electrode potentials and electrochemical cells

Where do redox reactions occur?

A

Take place in electrochemical cells where electrons are transferred from the reducing agent to the oxidising agent indirectly via an external circuit.

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2
Q

3.1.11 Electrode potentials and electrochemical cells

What is created in a redox reaction inside an electrochemical cell and what can it cause?

A
  1. A potential difference aka: voltage.
  2. can cause an electrical current to do work.
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3
Q

3.1.11 Electrode potentials and electrochemical cells

Define reduction.

A
  1. loss of oxygen
  2. gain of e-
  3. decrease in oxidation number
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4
Q

3.1.11 Electrode potentials and electrochemical cells

Define oxidation.

A
  1. gain of oxygen
  2. loss of e-
  3. increase in oxidation number.
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5
Q

3.1.11 Electrode potentials and electrochemical cells

Define reducing agents.

A

Electron donors: lose electrons.

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6
Q

3.1.11 Electrode potentials and electrochemical cells

Define oxidising agents.

A

Electron acceptors: gain electrons.

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7
Q

3.1.11 Electrode potentials and electrochemical cells

Ranking the power of reducing agents.

1 is the most reducing agent

A
  1. Li
  2. K
  3. Ca
  4. Mg
  5. Zn
  6. Ni
  7. Pb
  8. Cu
  9. Ag

Reducing agent will reduce anything below it.

Lies Kill Calm Minds Zen Night Please Calm Ants.

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8
Q

3.1.11 Electrode potentials and electrochemical cells

Ranking the power of oxidising agents.

1 is the most oxidising agent.

A
  1. Ag
  2. Cu
  3. Pb
  4. Ni
  5. Zn
  6. Mg
  7. Ca
  8. K
  9. Li

Oxidising agent can oxidise anything below it.

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9
Q

3.1.11 Electrode potential and electrochemical cells

How is an electrochemical cell formed?

A
  1. is made from two different metals dipped in salt solutions of their own ions and are connected by a wire (an external circuit) called a salt bridge.
  2. the metal acts as an electrode.
  3. there are always 2 reactions within an electrochemical cell - redox process.
  4. The 2 half cells produce a small voltage if connected to a circuit (battery or cell.)
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10
Q

3.1.11 Electrode potential and electrochemical cells

Using the example of zinc/copper, explain how an electrochemical cell / voltage is formed.

A

Zn (s) ⇌ Zn2+ (aq) + 2e-
Cu2+ (aq) + 2e- ⇌ Cu(s)

  1. The Zn half cell has more tendency to oxidise to the Zn2+ ion and release electrons than the Cu half cell.
  2. allows more electrons to build up on the Zn electrode than the Cu electrode.
  3. Voltage or PD is created between the two electrodes (can test using voltmeter in external circuit to measure the voltage.) This is known as cell potential / EMF / E cell.
  4. Zn cell is negative terminal and Cu cell is positive terminal.

metal that is easy to oxidise has a negative electrode potential and metal that is hard to oxidise has a positive electrode potential.

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11
Q

3.1.11 Electrode potential and electrochemical cells

What is the overall reaction that will happen in the cell?
Zn2+ (aq) / Zn(s) = -0.76
Cu2+(aq) / Cu(s) = +0.34

A

Zinc is oxidised so the half-equation goes backwards } more negative.
(Zn (s) ⇌ Zn2+ (aq) + 2e-)
Copper is reduced so the half-equation goes foward } more positive.
Cu2+ (aq) + 2e- ⇌ Cu(s)

OVERALL REACTION
Cu2+(aq) + Zn ⇌ Cu(s) + Zn2+(aq)

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12
Q

What is the purpose of the salt bridge?

A

Connects up the circuit. The free moving ions conduct the charge.

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13
Q

3.1.11 Electrode potential and electrochemical cells

How is the salt bridge made?

A

Made from a piece of filter paper soaked in salt solution } KNO3.

The salt bridge must be unreactive with the electrodes and electrode solutions i.e. KCl will not be suitable for copper systems as chloride can form complexes with copper ions.

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14
Q

3.1.11 Electrode potential and electrochemical cells

What happens if a current is allowed to flow through?

A

If voltmeter is removed and replaced with a bulb or if circuit is short-circuited, a current can flow.

Reactions occur seperately at each electrode.

Voltage fall to 0 as reactants are used up.

+ electrodes = reduction = e- are used up

  • electrodes = oxidation = e- are given off.
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15
Q

3.1.11 Electrode potential and electrochemical cells

Drawing electrochemical cells using example of zinc and copper.

EQ may ask conventional representation instead.

A

Reduced|Oxidised||Oxidised|Reduced
| = different phases seperated.
|| = salt bridge.

  1. More negative electrode potential / oxidation = left (zinc)
  2. Oxidised form goes in the centre of diagram.

Zn(s) + Zn2+ (aq) ⇌ Cu2+ (aq) + Cu(s)

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16
Q

3.1.11 Electrode potential and electrochemical cells

How would you work out the standard electrode potential of a system that does not include metals and what does the replacement provide?

A

Plantinum electrode is replacing a metal that can act as an electrode.
It is unreactive and can conduct electricity
provides conducting surface for e- transfer.

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17
Q

3.1.11 Electrode potential and electrochemical cells

Outline an example to highlight to work out the standard electrode potential of a system that does not include two metals.

Use Iron.

A

Fe2+ (aq) ⇌ Fe3+ (aq) + e-
there is no solid conducting the surface thus a plantinum electrode must be used.

Fe3+(aq), Fe2+(aq) | Pt(s)

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18
Q

3.1.11 Electrode potential and electrochemical cells

Why is it not possible to measure the electrode potential of a cell?

A

Cannot measure absolute potential of a half electrode on its own. } can only measure EPD between two electrodes.

Convenetion: assign relative potential to each electrode comparing it to standard hydrogen electrode which is given potential of 0 volts.

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19
Q

3.1.11 Electrode potential and electrochemical cells

Why are standard hydrogen electrodes (SHE) used?

A

Electrode potential of all electrodes are measured by comparing their potential to SHE.

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20
Q

3.1.11 Electrode potential and electrochemical cells

What value is assigned to the SHE?

A

0 volts

21
Q

3.1.11 Electrode potential and electrochemical cells

Outline the hydrogen electrode equilibrium.

A

H2(g) ⇌ 2H+(aq) + 2e-

22
Q

3.1.11 Electrode potential and electrochemical cells

Outline the cell diagram / conventinal representation of hydrogen electrode.

A

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

Pt used as it provides a conductibe surface for electron transfer.

23
Q

3.1.11 Electrode potential and electrochemical cells

What are the conditions for the standard electrode potential?

A

Temperature: 298K
Pressure: 100KPa
Concentration: 1.00 moldm-3 solution of ions. usually HCl.
Plantinum electrode

24
Q

3.1.11 Electrode potential and electrochemical cells

Why are standard conditions important?

A

Ensure the electrode potential value doesn’t change so you can compare values for different cells.

25
Q

3.1.11 Electrode potential and electrochemical cells

Why is a planinum wire / electrode used?

A

The equilibrium does not include a conducting metal surface. Plantinum can absorb hydrogen gas as it is porous.

26
Q

3.1.11 Electrode potential and electrochemical cells

What is the standard electrode potentials?

A

A LIST OF ELECTROCHEMICAL SERIES

when an electrode system is connected to the hydrogen electrode system.
Potential difference measured = standard electrode potential

27
Q

Why are electrochemical cells used?

A

Electrode potentials cannot be measured directly.

28
Q

How do you calculate cell potentials?

A

E(cell) = E(reduced) - E(oxidised)
or
E(cell) = E(positive) - E(negative)
or
E(cell) = E(RHS) - E(LHS)

R = reduction = positive = moves forward
O = oxidised = negative = goes backwards

29
Q

Do very reactive metals have more negative or positive electrode potentials?

A

More negative

30
Q

Do very reactive non-metals have more negative or positive electrode potentials?

A

More positive

31
Q

How do you know whether a reaction is feasible?

A

The EMF will be positive.

32
Q

How can e values be used to predict the direction of a simple redox reaction?

A

The equilibrium with more negative E value = move to the left.

The equilibrium with more positive E value = move to the right.

33
Q

Effect of conditons on cell voltage / Ecell:

A

If current is allowed to flow, cell reaction will occur and the Ecell will fall to 0 as the reaction continues and the reactant concentrations drop.

34
Q

Effect of concentration on cell voltage / Ecell :

A

Increase concentration of reactants would increase the E cell.

35
Q

Effect of temperature on cell voltage / E cell:

A

(Most E cells are exothermic) Increase in temperature = decrease in E cell as equilibrium reaction would shift backwards.

E cell positive = a reaction might occur (may occur slowly)

High activation energy: reaction will not

36
Q

What can electrochemical cells be used as?

A

Commerical source of electrical energy.
Can be non-rechargeable, rechargeable and fuel cells.

37
Q

What are non-rechargeable cells?

A

Reactions that occur are non-reversible.

38
Q

Outline a primary non-rechargeable cell:
Zn2+ + 2e- -> Zn = -0.76
2MnO2 + 2NH4+ + 2e- -> Mn2O3 + 2NH3 + H2O = 0.75

Dry cell

A

Zn2+ + 2e- -> Zn = -0.76
2MnO2 + 2NH4+ + 2e- -> Mn2O3 + 2NH3 + H2O = 0.75

  1. Zn -> Zn2+ + 2e- } more negative E value = reaction goes backwards.
  2. OVERALL EQUATION:
    2MnO2 + 2NH4+ + Zn -> Mn2O3 + 2NH3 + H2O + Zn2+
  3. EMF of cell = 0.75 - - 0.76 = +1.51v.
39
Q

What are rechargeable cells?

A

Can be recharged and used mutliple times and reversing the cell reaction. This is done by applying an external voltage greater than the voltage of the cell to move the electrons in opposite direction.

40
Q

Example of rechargeable cell:

A

Lead-acid batteries = operate starter motors of cars.
+ plate = lead coated with lead oxide = PbO2
- plate = lead

41
Q

PbSO4 + 2e- -> Pb + SO42- E= -0.356V

PbO2 + SO42- + 4H+ + 2e- -> PbSO4 + 2H2O E= +1.685V

Calculate overall equation and EMF.

A
  1. PbSO4 + 2e- -> Pb + SO42- = E= -0.356V
    } Pb(s) + SO42- -> PbSO4 + 2e-
  2. PbO2 + 4H+ + SO42- + 2e- -> PbSO4 + 2H2O E= +1.685V
  3. OVERALL EQUATION: PbO2 + 4H+ + 2SO42- + Pb -> 2PbSO4 +2H2O
    EMF = +2.04V
42
Q

Example secondary nickel–cadmium cells: what are they used for and what type of cell are they?

A

Power electrical equipment such as drills or shavers.
Rechargeable cells.

43
Q

Example secondary nickel–cadmium cells:
NiO(OH) + H2O + e- -> Ni(OH)2 + OH– E = +0.52 V (Ni will reduce changing oxidation state from 3 to 2)
Cd(OH)2 + 2e- -> Cd + 2OH– E = –0.88 V (Cd will oxidise changing oxidation state from 0 to 2)

A
  1. NiO(OH) + H2O + e- -> Ni(OH)2 + OH– = +0.52V
    Cd(OH)2 + 2e- -> Cd + 2OH– E = –0.88 V
  2. REVERSE THE MOST NEGATIVE:
    Cd + 2OH- - E -> Cd(OH)2 + 2e-
  3. BALANCE E- AND CANCEL OUT:
    NiO(OH) + H2O + e- -> Ni(OH)2 + OH– X2
    2NiO(OH) + 2H2O + 2e- -> 2Ni(OH)2 + 2OH–

Cd + 2OH- -> Cd(OH)2 + 2e-

2NiO(OH) + 2H2O + -> 2Ni(OH)2
Cd + -> Cd(OH)2

  1. OVERALL DISCHARGE EQUATION:
    2NiO(OH) + 2H2O + CD+ -> 2Ni(OH)2 + Cd(OH)2

5.EMF = POS - NEG
0.52 - - 0.88 = 1.40V

  1. OVERALL EQUATION WOULD BE REVERSED IN RECHARGING STATE
    2Ni(OH)2 + Cd(OH)2 -> 2NiO(OH) + Cd + 2H2O
44
Q

Example secondary Lithium ion cells: What are they used for and what type of cell are they?

A

Power cameras and mobile phones, tablets, latpots etc.

45
Q

Example secondary Lithium ion cells: What is the positive and negative electrode made out of?

A

Positive electrode: Lithium cobalt oxide = LiCoO2
Negative electrode: Carbon

46
Q

Example secondary Lithium ion cells: What reaction occurs at the negaitve electrode?

A

Li+ + e- -> Li E=-3V

47
Q

Example secondary Lithium ion cells: What reaction occurs at the positive electrode?

A

Li+(CoO2)- -> Li+ + CoO2 + e- E= +1v

48
Q
A

Physical chemistry. (13 decks)

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