5.5.3 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^θ)
• The standard electrode potential (E^θ) is 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^θ value

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

2H⁺(aq) + 2e^– ⇌ H₂(g)        E^θ = 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^θ value

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

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

 

• 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

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

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^ꝋ = + 0.80 V

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

• Since the Ag⁺/ Ag half-cell has a more positive E^ꝋ value, this is the positive pole and the H⁺/H₂ half-cell is the negative pole
• The standard cell potential (E_cell^ꝋ) is E_cell^ꝋ = (+ 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 E^ꝋ value
  • 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 Br₂ / 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^ꝋ = +1.09 V

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

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

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^ꝋ = +1.52 V

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

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

Ions in solution half cell

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

Worked Example

Answer

Al (s)∣Al³⁺ (aq) ∥ Zn²⁺ (aq)∣Zn (s)                  E _cell = +0.94 V

It is also acceptable to include phase boundaries on the outside of cells as well:

∣ Al (s)∣Al³⁺ (aq) ∥ Zn²⁺ (aq)∣Zn (s) ∣                  E _cell = +0.94 V

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