5.3.3 Standard Electrode & Cell Potentials

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^ꝋ)
• The standard electrode potential (E^ꝋ) 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^ꝋ value

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

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

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

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

Standard cell potential

• Once the E^ꝋ of 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 (E_cell^ꝋ)
• The standard cell potential can be determined by two methods:
  1. Using the equation E_cell^ꝋ = E_reduction^ꝋ – E_oxidation^ꝋ
     • Use of this equation does require knowledge of which reaction is reduction and which is oxidation
     • The reduction reaction has the higher / more positive E^ꝋ value
  2. E_cell^ꝋ is the difference in E^ꝋ between two half-cells
• For example, an electrochemical cell consisting of bromine and sodium half-cells has an E_cell^ꝋ of:
  • E_cell^ꝋ = (+1.09) - (-2.71)
  • E_cell^ꝋ = +3.80 V

• When a metal rod is placed in an aqueous solution, a redox equilibrium is established between the metal ions and atoms
  • For example, the copper atoms get oxidised and enter the solution as copper ions

Cu(s) → Cu²⁺(aq) + 2e^-

Oxidation of copper ions

• • The copper ions gain electrons from the metal rod and deposit as metal atoms on the rod

Cu²⁺(aq) + 2e^- → Cu(s)

Reduction of copper ions

• • When equilibrium is established, the rate of oxidation and reduction of copper is equal

• The position of the redox equilibrium is different for different metals
  • Copper is more easily reduced, thus the equilibrium lies further over to the right

Cu²⁺ (aq) + 2e^- ⇌ Cu (s)

• • Vanadium is more easily oxidised, thus the equilibrium lies further over to the left

V²⁺ (aq) + 2e^- ⇌ V(s)

• The metal atoms and ions in solution cause an electric potential (voltage)
• This potential cannot be measured directly however the potential between the metal/metal ion system and another system can be measured
• This value is called the electrode potential (E) and is measured in volts
  • The electrode potential is the voltage measured for a half-cell compared to another half-cell
  • Often, the half-cell used for comparison is the standard hydrogen electrode

Standard hydrogen electrode

• The standard hydrogen electrode is a half-cell used as reference electrodes and consists of:
  • Hydrogen gas in equilibrium with H⁺ ions of concentration 1.00 mol dm^-3 (at 1 atm)

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 voltmeter

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