AQA A Level Chemistry

Revision Notes

5.4.1 Representing Cells

Representing Cells

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
    • For example, if you dipped a zinc metal rod into a solution which contained zinc ions, there would be zinc atoms losing electrons to form zinc ions and at the same time, zinc ions gaining electrons to become zinc atoms
    • This would cause a redox equilibrium
  • 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 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)        voltage = +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)        voltage = -2.71 V

Standard electrode potential

  • The position of equilibrium and therefore the electrode potential depends on factors such as:
    • Temperature
    • Pressure of gases
    • Concentration of reagents
  • So, to be able to compare the electrode potentials of different species, they all have to be measured against a common reference or standard
  • Standard conditions also have to be used when comparing electrode potentials
  • These standard conditions are:
    • Ion concentration of 1.00 mol dm-3
    • A temperature of 298 K
    • A pressure of 1 atm
  • The electrode potentials are measured relative to something called a standard hydrogen electrode
  • The standard hydrogen electrode is given a value of 0.00 V, and all other electrode potentials are compared to this standard
  • This means that the electrode potentials are always referred to as a standard electrode potential (E)
  • The standard electrode potential (E) is the voltage produced when a standard half-cell is connected to a standard hydrogen cell under standard conditions
  • For example, the standard electrode potential of bromine suggests that relative to the hydrogen half-cell it is more likely to get reduced, as it has a more positive E value

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

2H+(aq) + 2e⇌ H2(g)        E = 0.00 V

  • The standard electrode potential of sodium, on the other hand, suggests that relative to the hydrogen half-cell it is less likely to get reduced as it has a more negative E value

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

2H+ (aq) + 2e⇌ H2(g)        E = 0.00 V

Standard cell potential

  • Once the Eof a half-cell is known, the voltage of an electrochemical cell made up of two half-cells can be calculated
    • These could be any half-cells and neither have to be a standard hydrogen electrode
  • This is also known as the standard cell potential (Ecell)
    • The standard cell potential is the difference in E between two half-cells
    • For example, an electrochemical cell consisting of bromine and sodium half-cells has an Ecell of:

Ecell = (+1.09) – (-2.71)

= +3.80 V


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