Electrolysis Notes

Objectives

• Define the term “electrolysis.”
• Illustrate the use of reduction potentials to predict the products of electrolysis.
• Define Faraday’s Laws of Electrolysis.

Electrolysis

• In an electrolytic cell, an electric current is used to drive a non-spontaneous chemical reaction.
• Electrical energy is converted into chemical energy.
• Electrolysis is the process by which the decomposition of an electrode occurs due to the passage of an electrical current through it.

Electrolytic vs. Voltaic Cells

• Electrolytic Cell: Uses electrical energy to drive a non-spontaneous reaction.
• Voltaic Cell: A spontaneous redox reaction releases energy.
• Reversing the direction of electron flow changes a voltaic cell into an electrolytic cell.
• Voltaic Cell:
  • Anode: Oxidation (A → A⁺ + e⁻) (negative)
  • Cathode: Reduction (B⁺ + e⁻ → B) (positive)
  • Overall: A + B⁺ → A⁺ + B
  • \Delta G < 0
  • Spontaneous redox reaction releases energy.
  • The system (the cell) does work on the surroundings.
• Electrolytic Cell:
  • Anode: Oxidation (B⁻ → B + e⁻) (positive)
  • Cathode: Reduction (A⁺ + e⁻ → A) (negative)
  • Overall: A⁺ + B⁻ → A + B
  • \Delta G > 0
  • Nonspontaneous redox reaction absorbs energy to drive it.
  • The surroundings (the source of energy) do work on the system.

Downs Cell: Electrolysis of Molten NaCl

• The battery acts as an “electron pump,” driving electrons to the cathode (reduction) and away from the anode (oxidation).
• $Na^+$ ions migrate to the cathode, accept electrons, and are reduced: $Na^+ + e^- \rightarrow Na$
• $Cl^-$ ions migrate to the anode, lose electrons, and are oxidized: $2Cl^- \rightarrow Cl_2 + 2e^-$
• Overall reaction: $2Na^+(aq) + 2Cl^-(aq) \rightarrow 2Na(s) + Cl_2(g)$
• Molten sodium chloride is electrolyzed to yield sodium metal and chlorine gas.

Electrolysis of NaCl (aq)

• Electrolysis of molten sodium chloride involves $Na^+$ and $Cl^-$ ions.
• Electrolysis of aqueous sodium chloride involves $Na^+, Cl^-$, and $H_2O$.

Electrolysis of NaCl (aq) - Cathode Reaction

• Possible Cathode Reactions:
  • $2H^+ + 2e^- \rightarrow H<em>2$ $E°</em>{red} = 0.00V$
  • $Na^+ + e^- \rightarrow Na$ $E°_{red} = -2.71V$
  • $2H<em>2O + 2e^- \rightarrow H</em>2 + 2OH^-$ $E°_{red} = -0.83V$
• Due to self-ionization of $H_2O$, $H^+$ is produced.
• $H^+$ has the highest $E°_{red}$ and is most likely to be reduced at the cathode.
• However, $H_2O$ has a pH of 7, so there is a low concentration of $H^+$.
• Therefore, $H<em>2O$ is reduced at the cathode, evolving $H</em>2$ gas and leaving an excess of $OH^-$ ions.

Electrolysis of NaCl (aq) - Anode Reaction

• Possible Anode Reactions:
  • $2Cl^- - 2e^- \rightarrow Cl<em>2$ $E°</em>{ox} = -1.36V$
  • $2H<em>2O - 4e^- \rightarrow O</em>2 + 4H^+$ $E°_{ox} = -1.23V$
• $H<em>2O$ should be more readily oxidized than $Cl</em>2$ at the anode based on $E°_{ox}$.
• However, $Cl^-$ is oxidized at the anode.
• Forming $O<em>2$ from $H</em>2O$ requires an extra voltage (overpotential) that is too high.

Electrolysis of NaCl (aq) - Electrode Material

• The reaction at the electrode depends on the electrode material.
  • Platinum cathode: $2H^+ + 2e^- \rightarrow H_2$
  • Mercury cathode: $Na^+ + e^- \rightarrow Na$

Faraday’s Laws of Electrolysis

• The mass of a substance altered at an electrode during electrolysis is directly proportional to the quantity of electricity (electrical charge) transferred at that electrode.
  • $m \propto Q$
• For a given quantity of electricity, the mass of an elemental material altered at an electrode is directly proportional to the element's equivalent weight.
• Current is measured in Amperes, A
  • $Q = It$

Faraday’s Laws of Electrolysis Formula

• $m = \frac{MIt}{zF} = \frac{MQ}{zF}$
  • Where:
    • m = mass of the substance altered at an electrode
    • Q = total electric charge passed through the substance
    • F = Faraday constant (96,500 C/mol)
    • M = molar mass of the substance
    • z = valency number of ions of the substance (electrons transferred per ion)

Question 1

• Chromium metal is plated out from an acidic solution containing $CrO_3$:
  • $CrO<em>3 + 6H^+ + 6e^- \rightarrow Cr(s) + 3H</em>2O$
• Calculate:
  • (a) How many grams of chromium will be plated out by 24000 C?
  • (b) How long will it take to plate out 1.5g of chromium by using 12.5 ampere current? Atomic mass of Cr = 52, $Q = It$
• Answers:
  • 2.15 g
  • 22 min, 16.1 s

Question 2

• A current of 0.452 A is passed through an electrolytic cell containing molten $CaCl_2$ for 1.50 hrs.
• Write the electrode reactions and calculate the quantity of products (in grams) formed at the electrodes.
• Mass of Ca = 0.507 g
• Mass of $Cl_2$ = 0.897 g