AS/A-Level Chemistry- Electrode potential

🔎 🔎 first we look at half cells in AS/A-level Chemistry to help understanding electode potential:

Voltaic cells: 

a system which converts chemical energy into electrical. Voltaic cells are made of two half cells which each contain a chemical species present in a redox half equation.

There are two main types of voltaic cells:

• metal/metal ion half cells- contain a 
• metal rod in a solution of its aqueous ions
• ion/ion half cells- 

ions of the same element in different oxidation states with an inert metal electrode

Standard electrode potential

The tendency to be reduced and gain electrons, or the difference between the reduction potentials is measured as standard electrode potential: E, under standard condition (1 atm, 298K/25^oC, 1 mol dm^-3)

E_cell = -E_cathode - E_anode

The more negative the electrode potential value:

• The 

greater the tendency to lose electrons and undergo oxidation

• the less tendency to gain electrons and undergo reudction
• the 

less spontaneous is the reaction (A positive E°_cell means that the reaction will occur spontaneously as written. A negative E°_cell means that the reaction will proceed spontaneously in the opposite direction.)  

The standard electrode potential, E^θ, for a half-cell is when a standard half-cell (1 mol dm^-3, 25^oC) is connected to a standard hydrogen electrode (1 mol dm^-3 H⁺, 25^oC, 1 atm) using a high resistance voltmeter. The hydrogen electrode is defined as having a standard electrode potential of 0.00 Volts.

Standard electrode potential of a metal – The metal dips into a 1M solution of its ions with a salt bridge to an standard hydrogen electrode.

E^θ_cell = E^θ(standard hydrogen electrode) - E^θ(metal ion(aq)/metal(s)

👇🏻 Example with zinc electrode and a reference electrode (the standard hydrogen electrode).

2H+(aq)+2e−→H2(g)E°cathode=0V Cathode:

2H⁺ (aq) +2e^- ➔ H₂ (g) ⇢ (E⁰ _cathode= 0 V)

Anode:

Zn(s) ➔Zn²⁺ (aq) + 2e^-  ⇢ (E⁰ _cathode= -0.76 V)

E⁰_cell = -E_cathode - E_anode

E⁰_cell = 0.76 V

Standard electrode potential of a gas – E.g. chlorine gas, platinum plate dips into a 1M solution of NaCl with chlorine gas, at 1.0 atm, bubbling over the platinum.

Standard electrode potential of an ion pair – E.g. Manganate (VII), one electrode is made of a piece of platinum dipping into a solution that is 1M in MnO₄^-, H⁺ and Mn²⁺ ions. This is connected to an SHE.

‼️ ‼️ If the conditions are not standard, the value of the electrode potential will alter. The direction of change can be predicted using Le Chatelier’s principle in AS/A-level Chemistry:

• Change in concentration – 

E.g. Cr₂O₇^2- (aq) + 14H⁺ (aq) + 6e-  <=> 2Cr³⁺(aq) + 7H₂O (l)              E^θ = +1.33V

If the concentrations of dichromate (VI) and hydrogen ions are increased above 1M, the position of equilibrium is driven to the right. This causes the value of the electrode potential, E, to be higher than the standard value.

• Change in pressure - 

Affects gaseous reactants only, e.g. O₂(g) + 2H₂O(l) + 4e- <=> 4OH-       E^θ = +0.40V

However, in air the partial pressure of oxygen is 0.2atm, not 1atm. Therefore, the equilibrium is driven to the left, making E < +0.40V

• Change in temperature – 

Effect depends on whether the redox half-equation is exothermic or endothermic. If it is exothermic an increased temperature drives the position of equilibrium to the left. This makes the value of the electrode potential less positive.

When doing calculations/answering questions:

1. Identify the two reactants– one reactant will be on the LHS of one half-equation and the other on the RHS of the second half-equation.

2. The half-equation with reactant on the RHS must be reversed. This alters the sign of its E° value.

3. If necessary, the two half-equations are multiplied by integers to give the same number of electrons in each equation and this does NOT alter the E° value.

4. The overall equation is obtained by adding these half-equations together.

5. The E° value of the reversed half-equation is then added to the E° value of the unchanged half-equation to give the E°_cell value.

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Drafted by Eunice Wong (Chemistry)

References:

• https://byjus.com/chemistry/standard-electrode-potential/
• https://chem.libretexts.org/Courses/Mount_Royal_University/Chem_1202/Unit_6%3A_Electrochemistry/6.2%3A_Standard_Electrode_Potentials