3.1.11 - Electrode potentials and electrochemical cells

RECAP:

What is a redox reaction?

One in which reduction and oxidation occurs in the same reaction

RECAP:

What is a disproportionation reaction?

reaction in which an element is both oxidised and reduced

RECAP:

What is a redox reaction?

One in which reduction and oxidation occurs in the same reaction

RECAP:

What is a disproportionation reaction?

reaction in which an element is both oxidised and reduced

RECAP:

What is an oxidising agent (oxidant)?

A substance that reduces itself to oxidise another species

RECAP:

What is a reducing agent (reductant)?

a substance that oxidises itself to reduce another species

When is an equilibrium set up in an electrochemical cell?

when a piece of metal (electrode) is dipped into a solution of its metal ions

What does it mean if the electrode has a negative potential?

the equilibrium lies to the left and the metal has a negative charge due to a build up of electrons on the metal

AANN: Oxidation/Anode/Negative

What does it mean if the electrode has a positive potential?

The equilibrium lies to the right and the metal has a positive charge as electrons have been used up to form metal from the metal ions

What does the position of the equilibrium depend on?

• explain

Depends on the metal

• reactive metals tend to form M2+ ions, so negative charge builds up on the metal = negative potentials

• Unreactive metals tend to have positive charge on metal = positive potentials

Definition of a half-cell/electrode

A metal dipping into a solution of its ions

What metal electrode is used in a half-cell with no metal involved in the half-equation?

• why?

platinum

• inert

• conducts electricity

If the surface area of the platinum electrode was doubled, what would happen to the emf of the cell?

it would remain unchanged

Definition of the electrochemical series

list of electrode potentials in numerical order

How do you measure the potential of an electrode?

combine two half-cells together (electrochemical cell) and find the potential difference between the two half-cells measured

What do you use to join two half cells together? why do you use these?

1. Voltmeter = to measure potential difference

2. Salt bridge = allows movement of ions to complete circuit

What is a salt bridge?

either:

• a piece of filter paper soaked with a solution of unreactive ions

• A tube containing unreactive ions in an agar gel

Why is a salt bridge needed?

It allows the movement of ions between electrodes to complete the circuit

What compounds are used for the salt bridge and why?

KNO₃

• Does not react with electrodes or electrolytes

What does SHE stand for? What is it?

SHE = Standard hydrogen electrode

An electrode assigned with 0 volts of potential and is used in comparison with other half-cells

What do electrochemical cells use?

uses electron transfer reactions i.e. redox to produce energy

Draw the electrochemical cell of Cu(s)/Cu²⁺(aq) and Zn(s)/Zn²⁺(aq)

1. write the half equation of each half cell

2. Which direction do the electrons flow?

3. Name with half cell is oxidised and which is reduced

4. Name the anode and the cathode

5. What happens to the mass of each electrode

6. Give the overall equation

1. Cu²⁺(aq) + 2e^- ⇌ Cu(s)

   Zn²⁺(aq) + 2e^- ⇌ Zn(s)

2. Electrons flow from Zn → Cu

3. Zn²⁺(aq) + 2e^- ⇌ Zn(s) = OXIDATION

   Cu²⁺(aq) + 2e^- ⇌ Cu(s) = REDUCTION

4. ANODE = Zn (supply e-)

   CATHODE = Cu (uses e-)

5. Zn decreases in mass

   Cu forms on positive electrode = increases in mass

6. Zn + Cu²⁺ → Zn²⁺ + Cu

If given cell potential values how do you figure out which cell is oxidised and which is reduced?

NO PROBLEM

MORE NEGATIVE = oxidised

MORE POSITIVE = reduced

How do you calculate the total E°cell value?

Reduced E° cell value - Oxidised E° cell value

What happens to the total E°cell value/EMF if the concentration of ions in one electrode is increased?

• why?

value: increases

• more ions (be specific) to accept/donate e- (depends on either ions are reduced or oxidised)

• Electrode becomes more positive/negative

What are the standard conditions of a half-cell?

• cell concentration = 1 mol/dm3 of ions involved in half equation

• Cell temperature = 298K

• Cell pressure = 100kPa (only affects half-cells with gases)

Why is the potential exactly 0V under the standard conditions?

By definition

Explain what happens to the equilibrium of this cell that is not under stand conditions?

• state which cell is oxidised and which is reduced giving the equation for both

• more concentrated Cu²⁺ ions on left = eqm shifts to right

  Left cell = reduction (+ve electrode) = Cu + 2e^- → Cu²⁺

• Less concentrated Cu²⁺ ions on right = eqm shifts to left

  Right cell = oxidation (-ve electrode) = Cu²⁺ → Cu + 2e^-

Order in which you write a conventional cell notation

1. Oxidation is on left in order of being oxidised

1. Reduction is on right in order of being reduced

2. || means salt bridge

3. | separates phases (e.g. s, aq)

4. , between same phase

What is the exception to this order?

The SHE (standard hydrogen electrode) is always on the LHS

Write the cell notation for the following:

• Cu/Cu2+ = +0.34V

• Zn/Zn2+ = -0.76V

Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu (s)

Write the cell notation for the following:

• Fe²⁺/Fe³⁺ = +0.77V

• H₂/H⁺ = 0V

Pt (s) | H₂ (g) | H⁺ (aq) || Fe³⁺ (aq), Fe²⁺ (aq) | Pt (s)

From the previous cell:

• name the solutions used in each cell

• H+/H2 = HCl

• Fe2+/Fe3+ = FeCl2, FeCl3

When is a reaction feasible?

• +ve value = feasible

• -ve value = not feasible

From the electrochemical series explain the trend of:

1. Reducing agents

2. Oxidising agents

1. Best reducing agents = bottom (reducing ability increases down the series)

2. Best oxidising agents = top (oxidising ability decreases down the series)

Name 3 types of commercial cells

• non-rechargeable

• Rechargeable

• Fuel cells

Briefly explain what a non-rechargeable cell is

• Cell where the chemicals are used up over time and the emf drops

• When one reactant is completely used up, the cell is flat and the emf is 0 volts

• Cannot be recharged and have to be disposed after a single use

Name the two types of non-rechargeable cells you need to know

• zinc-carbon

• Alkaline

Zinc-carbon cell:

1. Characteristics of this cell

2. Give the two half equations involved in the cell

3. State which is oxidised/reduced

4. give overall equation

5. Cell emf

1. cheap but short life

2. Zn(NH₃)₂²⁺ + 2e^- ⇌ Zn + 2NH₃

   2MnO₂ + 2H⁺ + 2e^- ⇌ Mn₂O₃ + H₂O

3. Zinc equation is oxidised

   Second equation is reduced (more positive value in series)

4. 2MnO₂ + 2H⁺ + Zn + 2NH₃ → Zn(NH₃)₂²⁺ + Mn₂O₃ + H₂O

5. 0.70 - (-0.80) = +1.50V

Alkaline cell:

1. Characteristics of this cell

2. Give the two half equations involved in the cell

3. State which is oxidised/reduced

4. give overall equation

5. Cell emf

1. Higher cost cell, longer life

2. Zn²⁺ + 2e^- ⇌ Zn

   MnO₂ + H₂O + e^- ⇌ MnO(OH) + OH^-

3. Zinc equation is oxidised

   Second equation is reduced (more positive value in electrochemical series)

4. Zn + 2MnO₂ + 2H₂O → 2MnO(OH) + 2OH^- + Zn²⁺

5. 0.84 - (-0.76) = +1.60V

Name the three types of rechargeable cells you need to know

• lithium ion

• Lead-acid

• Nickel-cadmium

Lithium ion cell:

1. Uses

2. Give the two half equations involved in the cell

3. State which is oxidised/reduced

4. overall equation during discharge

5. Overall equation during re-charge

6. Cell emf

7. Cell notation

1. Phones, cameras

2. Li⁺ + CoO₂ + e^- ⇌ LiCoO₂

   Li⁺ + e^- ⇌ Li

3. First equation is reduced

   Second equation is oxidised

4. Li + CoO₂ → LiCoO₂ (normal overall equation (Li⁺ cancels out))

5. LiCoO₂ → Li + CoO₂ (reverse equation)

6. 0.60 - (-3.00) = 3.60V

7. Li (s) | Li⁺ (aq) || CoO₂ (aq) | LiCoO₂ (s) | Pt (s)

LITHIUM ION CELL:

Write half equation for reaction at the negative electrode during operation

Li → Li+ +e-

LITHIUM ION CELL:

Suggest why water is not used as a solvent in this cells

Water would react with the lithium and create an explosion

LITHIUM ION CELL:

Suggest why the recharging of a lithium cell may lead to the release of carbon dioxide in the atmosphere

Electricity for recharging cell may come from power stations burning fossil fuels

Lead-acid cell:

1. Uses

2. Give the two half equations involved in the cell

3. State which is oxidised/reduced

4. overall equation during discharge

5. Overall equation during re-charge

6. Cell emf

7. In this reaction what does sulfuric acid act as?

1. Cars

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

   PbSO₄ + H⁺ + 2e^- ⇌ Pb + HSO4^-

3. First equation = reduced

   Second equation = oxidised

4. 2H⁺ + Pb + 2HSO4^- + PbO₂ → 2PbSO₄ + 2H₂O

5. 2PbSO₄ + 2H₂O → 2H⁺ + Pb + 2HSO4^- + PbO₂

6. 1.68 - (-0.36) = +2.04V

7. Electrolyte

Nickel-cadmium cell:

1. Give the two half equations involved in the cell

2. State which is oxidised/reduced

3. overall equation during discharge

4. Overall equation during re-charge

5. Cell emf

1. NiO(OH) + 2H₂O + 2e^- ⇌ Ni(OH)₂ + 2OH^-

   Cd(OH)₂ + 2e^- ⇌ Cd + 2OH^-

2. First equation = oxidation

   Second equation = reduction

3. Cd + NiO(OH) + 2H₂O → Ni(OH)₂ + Cd(OH)₂

4. Ni(OH)₂ + Cd(OH)₂ → 2H₂O + NiO(OH) + Cd

5. 0.52 - (-0.88) = +1.40V

What are fuel cells?

• most common

• continuous supply of chemicals into the cell so do not need recharging as never run out of chemicals

• Most common fuel cell = hydrogen-oxygen fuel cell (which can be run in alkaline and acidic conditions)

Hydrogen-oxygen fuel cell (alkaline conditions)

1. equation at negative electrode

2. Equation at positive electrode

3. Overall equation

4. Cell emf

5. Cell notation

1. H₂ + 2OH^- → 2H₂O + 2e^-

2. O₂ + 2H₂O + 4e^- → 4OH^-

3. 2H₂ + O₂ → H₂O

4. +1.23V

5. Pt (s) | H₂ (g) | OH^- (aq), H₂O (l) || O₂ (g) | H₂O (l), OH^- (aq) | Pt (s)

Hydrogen-oxygen fuel cell (acidic conditions)

1. equation at negative electrode

2. Equation at positive electrode

3. Overall equation

4. Cell emf

5. Cell notation

1. H₂ → 2H⁺ + 2e^-

2. O₂ + 4H⁺ + 4e^- → 2H₂O

3. 2H₂ + O₂ → H₂O

4. +1.23V

5. Pt (s) | H₂ (g) | H⁺ (aq) || O₂ (g) | H⁺ (aq), H₂O (l) | Pt (s)

Benefits and risks of:

• using non-rechargeable cells

BENEFITS = cheap

RISKS = waste issues

Benefits and risks of:

• using rechargeable cells

BENEFITS = Less waste, cheaper in long run, lower environmental impact

RISKS = some waste issues after useful life

Benefits and risks of:

• using hydrogen fuel cells

BENEFITS = only waste product is water, do not need re-charging, very efficient

RISKS = need constant supply of fuels, hydrogen is flammable and explosive, hydrogen usually made using fossil fuels, high cost of fuel cells

Why do hydrogen fuel cell not need recharging

because fuel supplied continuously to cell so voltage output does not change

Explain why the current in the circuit of this cell falls to zero after the cell has operated for some time

(In a cell that has solutions with different concentrations)

Eventually the ions will be the same in each electrode

Suggest one reason why waste disposal centres have a separate section for cells and batteries?

pollution from toxic chemicals

If a question asks to give a reason (other than cost) why a cell is not recharged, what should you do?

See if it is reversible, if its not then it will not be rechargged

With a table of E°cell values how would you know which is the weakest oxidising/reducing agent?

Weakest oxidising agent = most negative value

Weakest reducing agent = most positive value

EXAM QUESTION: Use the data in the table to explain why copper does not react with most acids but does react with nitric acid

1. 0.00 - 0.34 = -0.34V

2. E°cell for Cu with most acids (-0.34) is negative showing Cu is a less powerful reducing agent than H2

3. 0.96 - 0.34 = +0.62V

4. E°cell for Cu with nitric acid is positive

State how you would change an electrochemical cell apparatus to allow the cell reaction to go to completion

remove voltmeter and add ammeter

Why would any electrode potentials that you calculate in experiments be different from the actual standard electrode potential for that electrode?

non standard conditions