Representing Cells (AQA A Level Chemistry)

Representing Cells

• Electrochemical cells generate electricity from spontaneous redox reactions

• For example:

Zn (s)  + CuSO₄ (aq)→ Cu (s)  + ZnSO₄ (aq)

• Instead of electrons being transferred directly from the zinc to the copper ions, a cell is built which separates the two redox processes

• For example:

Zn (s)  ⇌  Zn²⁺ (aq) + 2e^– 

• If a rod of metal is dipped into a solution of its own ions, an equilibrium is set up

• Each part of the cell is called a half cell

When a metal is dipped into a solution containing its ions an equilibrium is established between the metal and its ions. This is the basis of a half cell in an electrochemical cell.

• This is a half cell and the strip of metal is an electrode

• The position of the equilibrium determines the potential difference between the metal strip and the solution of metal

• The Zn atoms on the rod can deposit two electrons on the rod and move into solution as Zn²⁺ ions:

Zn(s) ⇌ Zn²⁺(aq) + 2e^– 

• This process would result in an accumulation of negative charge on the zinc rod

• Alternatively, the Zn²⁺ ions in solution could accept two electrons from the rod and move onto the rod to become Zn atoms:

Zn²⁺(aq) + 2e^–  ⇌ Zn(s)

• This process would result in an accumulation of positive charge on the zinc rod

• In both cases, a potential difference is set up between the rod and the solution

• This is known as an electrode potential

• A similar electrode potential is set up if a copper rod is immersed in a solution containing copper ions (e.g. CuSO₄), due to the following processes:

Cu²⁺(aq) + 2e^–  ⇌ Cu(s)   – reduction (rod becomes positive)

Cu(s) ⇌ Cu²⁺(aq) + 2e^–    – oxidation (rod becomes negative)

Oxidation of copper atoms

Reduction of copper(II) ions

• Note that a chemical reaction is not taking place – there is simply a potential difference between the rod and the solution

• The potential difference will depend on

  • the nature of the ions in solution

  • the concentration of the ions in solution

  • the type of electrode used

  • the temperature

Electrode potential

• The electrode (reduction) potential (E) is a value which shows how easily a substance is reduced

• These are demonstrated using reversible half equations

  • This is because there is a redox equilibrium between two related species that are in different oxidation states

• When writing half equations for this topic, the electrons will always be written on the left-hand side (demonstrating reduction)

• The position of equilibrium is different for different species, which is why different species will have different electrode (reduction) potentials

• The more positive (or less negative) an electrode potential, the more likely it is for that species to undergo reduction

  • The equilibrium position lies more to the right

• For example, the positive electrode potential of bromine below, suggests that it is likely to get reduced and form bromide (Br^-) ions

Br₂(l) + 2e^- ⇌ 2Br^-(aq)        E = +1.09 V

• The more negative (or less positive) the electrode potential, the less likely it is that reduction of that species will occur

  • The equilibrium position lies more to the left

• For example, the negative electrode potential of sodium suggests that it is unlikely that the sodium (Na⁺) ions will be reduced to sodium (Na) atoms

Na⁺(aq) + e^- ⇌ Na(s)        E = -2.71 V

Conventional Representation of Cells

• Chemists use a type of shorthand convention to represent electrochemical cells

• In this convention:

  • A solid vertical (or slanted) line shows a phase boundary, that is an interface between a solid and a solution

  • A double vertical line (sometimes shown as dashed vertical lines) represents a salt bridge

    • A salt bridge has mobile ions that complete the circuit

    • Potassium chloride and potassium nitrate are commonly used to make the salt bridge as chlorides and nitrates are usually soluble

    • This should ensure that no precipitates form which can affect the equilibrium position of the half cells

  • The substance with the highest oxidation state in each half cell is drawn next to the salt bridge

  • The cell potential difference is shown with the polarity of the right hand electrode

• The cell convention for the zinc and copper cell would be

Zn (s)∣Zn²⁺ (aq) ∥Cu²⁺ (aq)∣Cu (s)                  E _cell = +1.10 V

• This tells us the copper half cell is more positive than the zinc half cell, so that electrons would flow from the zinc to the copper

• The same cell can be written as:

Cu (s)∣Cu²⁺ (aq) ∥Zn²⁺ (aq)∣Zn (s)                  E _cell = -1.10 V

• The polarity of the right hand half cell is negative, so we can still tell that electrons flow from the zinc to the copper half cell