Chem 2 Ch 20 Electrochemistry Notes

Electrochemistry Overview

  • Electrochemistry is the study of the interplay between electricity and chemical reactions.

  • It encompasses both spontaneous and nonspontaneous processes.

Oxidation Numbers

  • Basic Rules for Assigning Oxidation Numbers:

    • Elements: 0

    • Monatomic Ion: Its charge

    • Fluorine: -1

    • Oxygen: -2 (except in peroxides which is -1)

    • Hydrogen: +1 (except in metal hydrides which is -1)

    • The sum of oxidation numbers must equal the total charge of the compound, which is 0 for neutral compounds.

Oxidation and Reduction

  • Oxidation: An increase in oxidation number, representing the loss of electrons.

  • Reduction: A decrease in oxidation number, representing the gain of electrons.

  • Agents:

    • An oxidizing agent (e.g., $H^+$) facilitates oxidation.

    • A reducing agent (e.g., $Zn$) facilitates reduction.

Half-Reactions

  • Oxidation and reduction half-reactions are written separately for balancing redox reactions.

  • Example for oxidation:

    • Zn(s)
      ightarrow Zn^{2+}(aq) + 2e^-

  • Example for reduction:

    • 2H^+(aq) + 2e^-
      ightarrow H_2(g)

Balancing Redox Equations (Half-Reaction Method)

  1. Write out two half-reactions, oxidation and reduction.

  2. Balance non-oxygen and non-hydrogen atoms.

  3. Balance oxygen using water, and hydrogen using protons ($H^+$).

  4. Add electrons to both sides to balance the charges.

  5. If necessary, multiply to equalize electrons in both half-reactions.

  6. Add half-reactions together and simplify if needed.

Balancing in Basic Solutions

  • To balance a redox reaction in a basic solution, balance it as if it were acidic. After balancing:

    • Add hydroxide ($OH^-$) to each side to neutralize $H^+$ and create water. Remove any water molecules from both sides as appropriate.

Voltaic Cells

  • Function: Convert energy from spontaneous redox reactions into electrical energy.

  • Components:

    • Anode (oxidation occurs)

    • Cathode (reduction occurs)

    • Salt bridge to maintain charge balance.

  • Electrons flow from the anode to the cathode through an external circuit.

Electromotive Force (emf)

  • The potential difference between anode and cathode; indicates the tendency of electrons to flow.

  • Measured in volts (V): 1 V = 1 J/C.

Standard Reduction Potentials

  • Reduction potentials are compared to that of hydrogen, which is set as 0 V under standard conditions (1 M, 1 atm, 25°C).

  • Stronger oxidizing agents have more positive reduction potentials.

  • Standard cell potentials can be calculated using:

    • E{cell} = E{cathode} - E_{anode}

Free Energy and Redox Reactions

  • Relationship between cell potential and Gibbs free energy:


    • \Delta G = -nFE

    • Where $n$ is moles of electrons transferred and $F$ is Faraday’s constant (96,485 C/mol).

Nonstandard Conditions

  • Use Gibbs free energy under varying conditions with the Nernst equation:


    • E = E^\circ - \frac{0.0592}{n} \log Q

Applications of Electrochemistry

  1. Batteries: Portable sources of electrochemical power; can be primary (non-rechargeable) or secondary (rechargeable).

  2. Corrosion Prevention: Methods include cathodic protection using sacrificial anodes.

  3. Electrolysis: Nonspontaneous reactions driven by electrical energy for chemical synthesis.

Electrolysis Concepts

  • Nonspontaneous reactions can occur if an external voltage is applied.

  • Charge (Q) can be calculated as: Q = It = nF where:

    • $I$ = current (A), $t$ = time (s), $n$ = moles of electrons.

Conclusion

  • Understanding these principles of electrochemistry is essential for applications in batteries, corrosion prevention, and industrial electrochemical processes.