A-level Chemistry: 3.1.9 Electrode Potentials

Made from 2 different metals dipped in salt solutions of their own ions and connected by wire (external circuit)

Redox reactions occur within it

Make electricity

1. Zinc loses electrons more easily than copper

2. Zinc (from zinc electrode) is oxidised to from Zn2+(aq) ions

3. = releases electrons into external circuit

Same no. of electrons (as zinc releases) are taken from external circuit, reducing Cu2+ ions to copper atoms

By a salt bridge

Filter paper soaked in KNO3(aq)

Enables ions to flow through and balance out the charges

Electrons flow through wire from more reactive metal to less reactive one

Cell potential or EMF, known as Ecell

Fe2+ / Fe3+

On surface of platinum electrodes

∵ it’s inert

reversible

how easily each metal loses electrons (i.e. how easily it’s oxidised)

electrode potentials

very negative electrode potential

(On the LHS)

less negative or positive electrode potential

(On RHS)

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

∵ Cell potential is affected by…

• Temperature

• Pressure

• Concentration

• 1.00 mol dm-3

• 298K

• 100 kPa

Voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode

Platinum

An acid 1.00 mol dm-3 of H+(aq)

More negative electrode potentials mean that:

1. The right-hand substances are more easily oxidised

2. The left-hand substances are more stable

More postive electrode potentials mean that:

1. The right-hand substances are more stable

2. The left-hand substances are more easily reduced

• E*⦵cell = E⦵reduced — E⦵oxidised
  i. e. RHS - LHS

The standard cell potential of the Mg/Fe electrochemical cell is +2.35 V.

Half-cell with more negative potential goes on left

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Fe(s) | Fe2+(aq) || MnO4-(aq), H+(aq), Mn2+(aq) | Pt(s)

Pt(s) | H2(g) | H+(aq) || Cr2O72-(aq), Cr3+(aq), H+(aq) | Pt(s)

H = 0 V & Th = -1.90 V

4H+ + Th → 2H2 + Th4+

H2 + Br2 → 2H+ + 2Br-

Types of electrochemical cell which provide electricity

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

Zn + 2MnO → ZnO + Mn2O3

• Lithium batteries

• Lead-acid batteries

• Nickel/cadmium batteries

Found in lots of devices e.g. phones, laptops, cars

Lithium salt in organic solvent

Li → Li+ + e-

Li+ + CoO2 + e- → Li+[CoO2]-

+3.60 V

A current is supplied to force electrons to flow in the opposite direction around the circuit and reverse the reaction

Reactions that occur in non-rechargeable batteries are difficult/impossible to reverse

Used to operate the starter motor of cars

Lead plate

Lead oxide coated lead plate

H2SO4

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

| (V = +1.69)

PbSO4 + H+ + 2e- ⇌ Pb + HSO4-

| (V = -0.36)

PbO2 + Pb + 2H2SO4 ⇌ 2PbSO4 + 2H2O

| (Vcell = 2.05)

Li + CoO2 → Li+[CoO2]-

Used to replace zinc-carbon batteries

Cd(OH)2(s) + 2e- ⇌ Cd(s) + 2OH-(aq)

NiO(OH)(s) + H2O(l) + e- ⇌ Ni(OH)2(s) + OH-(aq)

2NiO(OH)(s) + Cd(s) + 2H2O(l) ⇌ 2Ni(OH)2(s) + Cd(OH)2(s)

(e- flow from Cd → Ni)

Generate electricity from hydrogen and oxygen

In electrodes and electrolyte

Chemicals stored separately outside cell and fed in when electricity is required

Alkaline hydrogen fuel cell

Used to power electric vehicles

2 separate platinum-containing electrodes

Catalyst

Separated by anion-exchange membrane that allows anions (OH-) and water to pass through but NOT hydrogen and oxygen gas

Aqueous alkaline (KOH) solution

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

Aqueous alkaline (KOH) solution

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

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

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

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

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

• Storing H2 absorbed into metals as metal hydrides mean if the cylinder is punctured, H2 escapes slowly

• Reduced explosion risk

To remove oxide layer on aluminium

• Carbonate ion react with acid (in the SHE)

• H+ concentration change/cell e.m.f altered

Increases surface area

Overall reaction is the same

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

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