Electrochemistry Notes

• Electrochemical Reaction: Major type is Voltaic (galvanic) cells that generate electricity using spontaneous reactions.

• Reaction Overview:
  

  [ \text{Zn} + \text{CuSO}4 \rightarrow \text{ZnSO}4 + \text{Cu} + 2e^- ]

• Electricity Generation: Produced by spontaneous electron flow, recognized as a redox reaction.

Oxidation and Reduction

• Oxidation Number Rules: - LEO says GER or OIL RIG:

  • Loses Electrons = Oxidation

  • Gains Electrons = Reduction

  • Oxygen's oxidation state is typically -2.

  • Pure elements have an oxidation state of 0.

  • Hydrogen is +1 unless bonded to a metal, then it is -1.

  • In ions: oxidation states sum to the ion's charge.

  • In a neutral compound, they sum to zero.

Practice Problems

Problem 1: Oxidation and Reduction

• Given Reaction:
  

  [ \text{Cu(s)} + 2\text{AgNO}3(aq) \rightarrow \text{Cu(NO}3)_2(aq) + 2\text{Ag(s)} ]

• Oxidized: ( \text{Cu} ) (loses electrons)

• Reduced: ( \text{Ag}^+ ) (gains electrons)

Problem 2: Half-Reactions

• Reaction:
  

  [ \text{Cu(s)} + 2\text{Ag}^+(aq) \rightarrow \text{Cu}^{2+}(aq) + 2\text{Ag(s)} ]

• Oxidation half-reaction:
  

  [ \text{Cu}^{0} \rightarrow \text{Cu}^{2+} + 2e^- ]

• Reduction half-reaction:
  

  [ 2 \text{Ag}^{+} + 2e^- \rightarrow 2 \text{Ag}^{0} ]

Balancing Equations with Half-Reaction Method

• Split the overall equation into oxidation and reduction half-reactions.

• Balance elements (if Hydrogen, add H+; if Oxygen, add H2O).

• Balance charges by adding electrons.

• Ensure electrons lost equal electrons gained, then combine and simplify.

Acidic vs. Basic Reactions

• Acidic reactions: Include H+ ions.

• Basic reactions: Include OH- ions.

Electrochemical Cells

• Voltaic cells are designed to separate reactants into half-cells where oxidation occurs at the anode and reduction occurs at the cathode.

• Salt Bridge: Necessary to complete the circuit and balance charge, allowing ions to migrate.

• Cell setup:- Reactants separated into half-cells with a wire connecting them.

  • A salt bridge must connect both half-cells.

Standard Reduction Potentials (SRP)

• SRP values indicate the likelihood of reduction in volts.

• Hydrogen (H2) serves as a baseline, E° = 0.

• Higher SRP = more likely to be reduced; lower SRP = more likely to be oxidized.

Calculating Cell Potential

• Cell Potential Equation:
  

  [ E°{cell} = E°{cathode} - E°_{anode} ]

• Positive E°: reaction is spontaneous.

• Negative E°: reaction is non-spontaneous.

Gibbs Free Energy (ΔG)

• ΔG and spontaneity: - If E°cell is positive, the forward reaction is spontaneous.

  • If E°cell is negative, the reaction is non-spontaneous.

• Calculate ΔG:
  

  [ ΔG = -nF E° ]
  

  Where n = moles of electrons, F = Faraday's constant (96,485 C/mol e-).

Practice Problems - Difficult Scenarios

• Be aware of unit changes (e.g., converting minutes to seconds) and symbols (like Q for the reaction quotient).

• Analyze real-world applications (like the time it takes for electrolysis of a specific volume of gas or metal plating allowances).

Key Takeaways

• Understand oxidation-reduction concepts clearly: gain/loss of electrons, identifying oxidized and reduced species.

• Be methodical with your half-reaction balancing.

• Use established potential readings to predict outcomes in electrochemical systems.

• Efficiency calculations (percent yields) are essential, especially in practical applications.

• Mastery of calculations involving currents, time, and molar relationships will improve problem-solving skills in electrochemistry specifics.