Electrolysis Products and the Electrochemical Series

Introduction to Electrolysis

• Electrolysis: The process of using electrical energy to drive a non-spontaneous chemical reaction.
• Electrolytic cells are used to carry out electrolysis.
• Key components:
  • Electrolyte: A substance containing free ions that conducts electricity (molten or aqueous solution).
  • Electrodes: Conductors where oxidation and reduction occur.
    • Anode: Positive electrode where oxidation occurs.
    • Cathode: Negative electrode where reduction occurs.
• External Power Source: Provides the electrical energy needed for the reaction.

KEY TAKEAWAY: Electrolysis uses electrical energy to force non-spontaneous redox reactions to occur.

The Electrochemical Series

• Electrochemical Series: A list of redox half-equations arranged in order of their standard electrode potential ($E^\circ$).
• Provides information about the relative ease of oxidation or reduction.
• Stronger oxidants are at the top left of the series (more likely to be reduced).
• Stronger reductants are at the bottom right of the series (more likely to be oxidized).

REMEMBER: “LEO says GER” (Lose Electrons Oxidation, Gain Electrons Reduction) and “AN OX, RED CAT” (Anode Oxidation, Reduction Cathode).

Predicting Electrolysis Products

Molten Electrolytes

• Predicting products is straightforward.
• The cation (positive ion) is reduced at the cathode.
• The anion (negative ion) is oxidized at the anode.
• Example: Electrolysis of molten NaCl
  • Cathode: $Na^+ + e^- \rightarrow Na$
  • Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$
  • Overall: $2NaCl(l) \rightarrow 2Na(l) + Cl_2(g)$

Aqueous Electrolytes

• Predicting products is more complex due to the presence of water.
• Water can be either oxidized or reduced, competing with the ions from the salt.
• At the Cathode (Reduction):
  • Consider the reduction of the cation and the reduction of water:
    • $2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)$ $E^\circ = -0.83V$
  • The species with the more positive $E^\circ$ value is preferentially reduced.
• At the Anode (Oxidation):
  • Consider the oxidation of the anion and the oxidation of water:
    • $2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-$ $E^\circ = +1.23V$
  • The species with the more negative $E^\circ$ value is preferentially oxidized.
• Exceptions:
  • Overpotential: In some cases, the voltage required for a reaction to occur is higher than the standard electrode potential suggests. This is particularly common for the oxidation of water to oxygen.
  • Concentration: High concentrations of halide ions (Cl$^{-}$, Br$^{-}$, I$^{-}$) can favour their oxidation over water, even if the $E^\circ$ values suggest otherwise.

Example: Electrolysis of Aqueous NaCl

• Possible reduction reactions at the cathode:
  • $Na^+(aq) + e^- \rightarrow Na(s)$ $E^\circ = -2.71V$
  • $2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)$ $E^\circ = -0.83V$
  • Water is preferentially reduced because it has a more positive $E^\circ$ value.
• Possible oxidation reactions at the anode:
  • $2Cl^-(aq) \rightarrow Cl_2(g) + 2e^-$ $E^\circ = +1.36V$
  • $2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-$ $E^\circ = +1.23V$
  • Despite water having a slightly less positive $E^\circ$, high concentration of $Cl^-$ usually leads to $Cl_2$ being produced.
• Overall: $2H_2O(l) + 2Cl^-(aq) \rightarrow H_2(g) + Cl_2(g) + 2OH^-(aq)$

EXAM TIP: Always write out the possible half-equations for both the cation/anion and water at each electrode. Then, compare the $E^\circ$ values to determine the likely products. Consider exceptions like overpotential and concentration.

Inert vs. Reactive Electrodes

• Inert Electrodes: Do not participate in the electrolysis reaction (e.g., platinum, carbon). They only serve as a surface for electron transfer.
• Reactive Electrodes: Participate in the electrolysis reaction. The electrode itself can be oxidized or reduced.

Example: Electrolysis of $CuSO_4$ with Copper Electrodes

• Cathode: $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$
• Anode: $Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-$
• Overall: Copper is transferred from the anode to the cathode (electrorefining).

APPLICATION: Electroplating uses electrolysis with a reactive anode to coat an object with a thin layer of metal.

Limitations of the Electrochemical Series

• Standard Conditions: The $E^\circ$ values are measured under standard conditions (298 K, 1 M concentration, 1 atm pressure). Deviations from these conditions can affect the actual electrode potential.
• Kinetics: The electrochemical series predicts thermodynamic feasibility, but not the rate of the reaction. Some reactions with favorable $E^\circ$ values may be slow.
• Overpotential: As mentioned earlier, some reactions require a higher voltage than predicted due to kinetic factors.
• Concentration Effects: The Nernst equation can be used to calculate electrode potentials under non-standard concentrations, but the electrochemical series only provides a qualitative guide.

Factors Affecting Electrolysis Products

• Electrolyte Concentration: Higher concentrations of ions can favor their reaction over water, even if $E^\circ$ values suggest otherwise.
• Electrode Material: Reactive electrodes can participate in the reaction, changing the products.
• Temperature: Affects the reaction rate and equilibrium.
• pH: In some cases, the pH of the solution can affect the products formed.

COMMON MISTAKE: Forgetting to consider the possible oxidation or reduction of water in aqueous electrolysis.

Summary Table

VCAA FOCUS: VCAA exam questions often involve predicting the products of electrolysis in aqueous solutions, considering the electrochemical series, concentration effects, and electrode materials.