AQA A Level Chemistry

Revision Notes

5.4.2 Standard Electrode Potentials

Standard Electrode Potentials

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 100 kPa
  • Standard measurements are made using a high resistance voltmeter so that no current flows and the maximum potential difference is achieved
  • The electrode potentials are measured relative to 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 (Eis the potential difference ( sometimes called 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 Hydrogen Electrode

  • The standard hydrogen electrode is a half-cell used as a reference electrode and consists of:
    • Hydrogen gas in equilibrium with H+ ions of concentration 1.00 mol dm-3 (at 100 kPa)

2H+ (aq) + 2e ⇌ H2 (g)

    • An inert platinum electrode that is in contact with the hydrogen gas and H+ ions
  • When the standard hydrogen electrode is connected to another half-cell, the standard electrode potential of that half-cell can be read off a high resistance voltmeter

Standard Hydrogen Electrode, downloadable AS & A Level Chemistry revision notes

The standard electrode potential of a half-cell can be determined by connecting it to a standard hydrogen electrode

  • There are three different types of half-cells that can be connected to a standard hydrogen electrode
    • A metal / metal ion half-cell
    • A non-metal / non-metal ion half-cell
    • An ion / ion half-cell (the ions are in different oxidation states)

Metal / metal-ion half-cell

Metal_Metal Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Example of a metal / metal ion half-cell connected to a standard hydrogen electrode

  • An example of a metal/metal ion half-cell is the Ag+/ Ag half-cell
    • Ag is the metal
    • Ag+ is the metal ion
  • This half-cell is connected to a standard hydrogen electrode and the two half-equations are:

Ag+ (aq) + e⇌ Ag (s)        E= + 0.80 V

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

  • Since the Ag+/ Ag half-cell has a more positive Evalue, this is the positive pole and the H+/H2 half-cell is the negative pole
  • The standard cell potential (Ecell) is Ecell = (+ 0.80) – (0.00) = + 0.80 V
  • The Ag+ ions are more likely to get reduced than the H+ ions as it has a greater Evalue
    • Reduction occurs at the positive electrode
    • Oxidation occurs at the negative electrode

Non-metal / non-metal ion half-cell

  • In a non-metal / non-metal ion half-cell, platinum wire or foil is used as an electrode to make electrical contact with the solution
    • Like graphite, platinum is inert and does not take part in the reaction
    • The redox equilibrium is established on the platinum surface
  • An example of a non-metal / non-metal ion is the Br/ Br half-cell
    • Br2 is the non-metal
    • Br is the non-metal ion
  • The half-cell is connected to a standard hydrogen electrode and the two half-equations are:

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

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

  • The Br/ Br half-cell is the positive pole and the H/ H2 is the negative pole
  • The Ecellis: Ecell = (+ 1.09) – (0.00) = + 1.09 V
  • The Br2 molecules are more likely to get reduced than H+ as they have a greater Evalue

Non-Metal_Non-Metal Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Example of a non-metal / non-metal ion half-cell connected to a standard hydrogen electrode

Ion / Ion half-cell

  • A platinum electrode is again used to form a half-cell of ions that are in different oxidation states
  • An example of such a half-cell is the MnO4/ Mn2+ half-cell
    • MnO4 is an ion containing Mn with oxidation state +7
    • The Mn2+ ion contains Mn with oxidation state +2
  • This half-cell is connected to a standard hydrogen electrode and the two half-equations are:

MnO4 (aq) + 8H+ (aq) + 5e⇌ Mn2+ (aq) + 4H2O (l)       E = +1.52 V

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

  • The H+ ions are also present in the half-cell as they are required to convert MnO4into Mn2+ ions
  • The MnO4/ Mn2+ half-cell is the positive pole and the H+ / H2 is the negative pole
  • The Ecell is Ecell = (+ 1.09) – (0.00) = + 1.09 V

Ion_ Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Ions in solution half cell

Calculating EMF

Standard cell potential

  • Once the Eof a half-cell is known, the potential difference or voltage or emf of an electrochemical cell made up of any two half-cells can be calculated
    • These could be any half-cells and neither have to be a standard hydrogen electrode
  • The standard cell potential (Ecell) can be calculated by subtracting the less positive Efrom the more positive Evalue
    • The half-cell with the more positive Evalue will be the positive pole
      • By convention this is shown on the right hand side in a conventional cell diagram, so is termed  Eright
    • The half-cell with the less positive Eꝋ value will be the negative pole
      • By convention this is shown on the left hand side in a conventional cell diagram, so is termed  Eleft

Ecell = Erightꝋ – Eleftꝋ   

    • Since oxidation is always on the left and reduction on the right, you can also use this version

Ecell = Ereductionꝋ – Eoxidation

Worked Example

Calculating the standard cell potential

Calculate the standard cell potential for the electrochemical cell below and explain why the Cu2+ / Cu half-cell is the positive pole.

The half-equations are as follows:

Cu2+(aq) + 2e⇌ Cu(s)      E= +0.34 V

Zn2+(aq) + 2e⇌ Zn(s)      E= −0.76 V

Electrochemistry Calculations - Electrochemical Cell, downloadable AS & A Level Chemistry revision notes

Answer

Step 1: Calculate the standard cell potential. The copper is more positive so must be the right hand side.

EcellErightꝋ Eleftꝋ   

Ecell = (+0.34) – (-0.76)

= +1.10 V

The voltmeter will therefore give a value of +1.10 V

Step 2: Determine the positive and negative poles
The Cu2+ / Cu  half-cell is the positive pole as its Eis more positive than the Evalue of the Zn2+ / Zn half-cell

Exam Tip

A helpful mnemonic for remembering redox in cells

Lio the Lion, downloadable AS & A Level Chemistry revision notes

 

Lio the lion goes Roor! 

Lio stands for ‘Left Is Oxidation’ and he is saying ROOR because that is the order of species in the cell:

Reduced/Oxidised (salt bridge) Oxidised/Reduced

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