3.1.11 Electrode potentials & electrochemical cells

What are electrochemical cells?

Two different metals dipped in salt solutions of their own ions & connected by a wire (the external circuit)

Describe this electrochemical cell

• Copper electrode is dipped in a solution of Cu²⁺(aq) ions & a zinc electrode is dipped in a solution of Zn²⁺(aq) ions

• Zinc loses electrons more easily than copper, so in the left half-cell, zinc is oxidised to Zn²⁺(aq) ions, releasing electrons into the external circuit

• In the right half-cell, the same number of electrons are taken from the external circuit, reducing the Cu²⁺ ions to copper atoms

What connects the two solutions in an electrochemical cell?

A salt bridge:

• a strip of filter paper soaked in potassium nitrate solution, KNO₃(aq)

• Allows the ions to flow between the half-cells & completes the circuit (balances out the charges)

Which way do electrons flow through the wire in an electrochemical cell?

From most reactive metal to the least reactive metal

How is the voltage/cell potential between two half-cells measured?

By putting a voltmeter in the external circuit

Describe this half-cell

A half-cell with solutions of two aqueous ions of the same element:

• a platinum electrode is dipped into the solution, as it is inert & conducts electricity

• conversion from Fe²⁺(aq) to Fe³⁺(aq), or vice versa, happens on the surface of the electrode

• the direction of the conversion depends on the other half-cell in the circuit

  • if the other cell contains a metal that is less reactive than iron, then Fe²⁺ will be oxidised to Fe³⁺ at the electrode

  • if the other cell contains a more reactive metal, Fe³⁺ will be reduced to Fe²⁺ at the electrode

The reactions that occur at each electrode in a cell are reversible. What determines which direction each reactions goes in?

How easily each metal is oxidised (measured using electrode potentials):

• a metal that’s easily oxidised has a very negative electrode potential

• a metal that’s harder to oxidise has a more positive electrode potential

Which metals are reduced/oxidised & why?

• Oxidised: zinc (zinc half-cell has a more negative electrode potential)

• Reduced: copper (copper half-cell has a more positive electrode potential)

How are electrochemical cells drawn (conventional representation) ?

ROOR → Reduced I Oxidised II Oxidised I Reduced:

1. Use the electrode potentials to work out which half-cell goes on the left & which goes on the right

2. Write out the left-hand half-equation as an oxidation reaction. Then write out the right-hand half-equation as a reduction reaction

3. Write out the reactants & products of the oxidation reaction, followed by the reactants & products of the reduction reaction

4. Add in a salt bridge by drawing two vertical lines between the half-cells

5. Draw a vertical line between any reagents that are in different phases. Put a comma between any reagents that are in the same phase

6. If either of the half-cells have a separate electrode (e.g. platinum), put that on the outside & separate it from the half-cell with a vertical line

Draw the conventional representation of this electrochemical cell

Write half-equations for the reactions occurring at the positive & negative electrodes of this cell

Positive electrode is on the right-hand side & the negative electrode is on the left-hand side:

• reduction occurs at the positive electrode: Pb²⁺ (aq) + 2e^- → Pb (s)

• oxidation occurs at the negative electrode: Al³⁺ (aq) → Al (s) + 3e^-

  • remember → think “NO PRoblem”

What is the equation to calculate the cell potential (E^θ_cell) ?

(will always be a positive voltage → more negative E^θ value is being subtracted from the more positive E^θ value)

Calculate the standard electrode potential of the aluminium half-cell

What factors affect the electrode potential?

Temperature, pressure & concentration:

• half-cell reactions are reversible → equilibrium position is affected by same changes just like any other reversible reaction

Changing the equilibrium position changes the cell potential. What is done to get around this issue?

Standard conditions are used to measure electrode potentials:

• you measure the electrode potential of a half-cell against a standard hydrogen electrode

  • hydrogen gas is bubbled through a solution of aqueous H⁺ ions

  • a platinum electrode is used as a platform for the oxidation/reduction reactions

What are the standard conditions to be used when measuring electrode potentials using the standard hydrogen electrode?

• Any solutions of ions (H⁺) must have a concentration of 1.00 moldm^-3 (e.g. acid, such as HCl)

• Temperature must be 298K

• Pressure must be 100kPa

What is the standard electrode potential of a half-cell?

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

What is the electrode potential of the standard hydrogen electrode half-cell?

0.00V

What is the conventional representation of this electrochemical cell?

What is an electrochemical series?

A list of electrode potentials in numerical order for different electrochemical cells:

• electrode potentials are written starting from the most negative & going down to the most positive

• half-equations are always written as reduction reactions

• oxidising agents are on the left & reducing agents are on the right

• best reducer is always the most negative (think if it’s negative, it’s reduced) & best oxidiser is always the most positive

When two half-equations are put together in an electrochemical cell, which directions do they go in?

• The half-equation with the more negative electrode potential goes in the direction of oxidation

• The half-equation with the more positive electrode potential goes in the direction of reduction

What is the equation to calculate the standard cell potential when two half-cell are connected together?

Work out which half-equation is going in the direction of oxidation & which half-equation is going in the direction of reduction

Calculate the cell potential of an Mg/Ag electrochemical cell using the two redox reaction equations

How can electrode potentials be used to predict whether a redox reaction will happen & which direction it will go in?

1. Find the two half-equations for the redox reaction & write them both as reduction reactions

2. Use an electrochemical series to work out which half-equation has the more negative electrode potential

3. Write out the half-equation with the more negative electrode potential going in the direction of oxidation & the half-equation with the more positive electrode potential going in direction of reduction

4. Combine the two half-equations & write out a full redox equation

This is the feasible direction of the reaction & will give a positive overall E^θ value

Work out the direction of the reaction when a Zn/Zn²⁺ half-cell is connected to a Cu²⁺/Cu half-cell

Deduce whether the following redox reaction is feasible or not: Cu(s) + Ni²⁺ (aq) → Cu²⁺(aq) + Ni(s)

Why do the half-equations have non-reversible arrows?

It’s impractical to reverse them in a battery:

• they can be made to run backwards under the right conditions, but this can make the battery leak or explode

• zinc electrode forms the casing of the battery, so becomes thinner as the zinc is oxidised

How are batteries recharged?

• A current is supplied to force electrons to flow in the opposite direction around the circuit & reverse the reactions

• This is possible as none of the substances in a rechargeable battery escape or are used up

What are fuel cells?

• Re-chargeable batteries in which a fuel donates electrons at one electrode & oxygen gains electrons at the other electrode

• As the fuel enters the cell it becomes oxidised, which sets up a potential difference/voltage within the cell

What is an example of a fuel cell & how does it work?

The hydrogen-oxygen fuel cell:

• hydrogen & oxygen gas are fed into two separate platinum-containing electrodes

• electrodes are separated by an anion-exchange membrane that allows anions (OH⁻ ions) & water to pass through it, but not hydrogen & oxygen gas. The electrolyte is an aqueous alkaline solution (KOH). Hydrogen is fed to the negative electrode

• reaction at negative electrode is: H₂(g) + 2OH⁻(aq) → 2H₂O(l) + 2e⁻

• the electrons flow from the negative electrode through an external circuit to the positive electrode. The OH⁻ ions pass through the anion-exchange membrane towards the negative electrode. Oxygen is fed to the positive electrode

• reaction at positive electrode is: O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)

• overall reaction: 2H₂(g) + O₂(g) → 2H₂O(l)

What are the advantages of fuel cells?

• Much more efficient than an internal combustion engine in machines (waste a lot of energy as thermal energy) → more energy is converted into kinetic energy

• Unlike batteries, fuel cells don’t need to be recharged → just need a continuous supply of oxygen & hydrogen gas

• The only waste product is water & no carbon dioxide is emitted directly from the cell, unlike a combustion engine

What are the disadvantages of fuel cells?

• Hydrogen is highly flammable & must be stored & transported correctly

• It is expensive to transport & store hydrogen → storage of hydrogen is in pressurised containers

• Energy is required to make the hydrogen & oxygen gas in the first place

  • fossil fuels are generally used to pass water through an electrolysis process → contributes to CO₂ emissions

Describe the zinc-carbon fuel cell

Non-rechargeable battery:

• uses Zn_(s)/Zn²⁺_(aq) & Cu_(s)/Cu²⁺_(aq) reactions

• zinc rod is the negative electrode & copper rod is the positive electrode

• at the negative electrode (oxidation): Zn(s) → Zn²⁺(aq) + 2e^-

• at the positive electrode (reduction): Cu²⁺(aq) + 2e^- → Cu(s)

• overall equation: Zn(s) + Cu²⁺(aq) ⇌ Zn²⁺(aq) + Cu(s)

How did Alessandro Volta discover non-rechargeable batteries?

• Used alternating discs of copper & zinc with damp pads between them

• Found that by connecting wires to the top & bottom disc, an electric charge passed around the external circuit

  • the voltage was proportional to the number of discs used → called this device a ‘pile’

What is the dry cell?

Non-rechargeable battery:

• contains a paste of manganese dioxide & ammonium chloride around a carbon rod (positive electrode)

• surrounded by a zinc case (negative electrode)

• Zn(s) → Zn²⁺(aq) + 2e^-

• MnO₂ + H₂O + e^- → MnO(OH) + OH^-

• Overall equation: Zn + 2MnO₂ + H₂O → Zn²⁺(aq) + 2MnO(OH) + 2OH^-

What is the lead-acid battery?

Re-chargeable battery:

• acting as electrodes are plates that separate the cells, immersed in dilute sulphuric acid

• negative electrodes are plates made from lead & positive electrodes are plates made from lead coated in lead (IV) oxide

• negative electrode: Pb (s) + SO₄^2- (aq) ⇌ PbSO₄ (s) + 2e^-

• positive electrode: PbO₂ (s) + SO₄^2- (aq) + 4H⁺ (aq) + 2e^- ⇌ PbSO₄ (s) + 2H₂O (l)

What are the simplified electrode reactions in a lithium cell?

• Positive electrode: Li⁺ + CoO₂ + e^- → Li⁺[CoO₂]^-

• Negative electrode: Li → Li⁺ + e^-

List the reactivity series & the mnemonic to remember it

(Most reactive)

• Please → Potassium

• Stop → Sodium

• Calling → Calcium

• Me → Magnesium

• A → Aluminium

• Careless → Carbon

• Zebra → Zinc

• Instead → Iron

• Try → Tin

• Learning → Lead

• How → Hydrogen

• Copper → Copper

• Saves → Silver

• Gold → Gold

(Least reactive)