Redox Reactions and Electrochemistry

Redox Reactions and Electrochemistry

Chapter 4: Types of Chemical Reactions

Oxidation-Reduction Reactions

• Oxidation State: The charge assigned to atoms in compounds, assuming bonding is 100% ionic. This is a formalism that aids in counting electrons.

• Reduction: The gaining of electrons, resulting in a decrease in oxidation state.

• Oxidation: The loss of electrons, resulting in an increase in oxidation state.

• Reductant (Reducing Agent): A compound that donates electrons and gets oxidized in a redox reaction.

• Oxidant (Oxidizing Agent): A compound that accepts electrons and gets reduced in a redox reaction.

• Redox Reaction: A reaction involving electron transfer.

  • OIL RIG: Oxidation Is Loss, Reduction Is Gain.

  • LEO the lion goes GER: Loss of Electrons is Oxidation, Gain of Electrons is Reduction.

Oxidation State Assignments

1. The sum of oxidation states equals the charge of the compound or ion.

   • This relates to the concept of formal charge.

2. Atoms in their elemental forms have oxidation states equal to 0.

3. In compounds:

   • Group 1 elements (Li group) have an oxidation state of +1.

   • Group 2 elements (Be group) have an oxidation state of +2.

   • Halogens (F and below) typically have an oxidation state of -1.

   • Hydrogen:

     • +1 when bound to a non-metal.

     • -1 when bound to a metal (forming a hydride).

   • Oxygen:

     • -2, unless bound to itself (e.g., in peroxides O_2^{2-}, where it's -1) or to fluorine (where it would be positive).

     • Examples: O2 (0), O2^{2-}$ (-1)

Examples of Assigning Oxidation States and Formal Charges

• HCl:

  • Formal Charge: H (0), Cl (0)

  • Oxidation States: H (+1), Cl (-1)

• CH4:

  • Formal Charge: C (0), H (0)

  • Oxidation States: C (-4), H (+1)

• PO_4^{3-}

  • O

  • Oxidation State: O (-2), P (+5)

• LiAlH_4

  • Oxidation State: Li (+1), Al (+3), H (-1)

Balancing Redox Reactions in Aqueous Solutions

• Reactions must be balanced in terms of the number of atoms and types, and also in terms of charge (number of electrons).

• The number of electrons lost by the reductant must equal the number of electrons gained by the oxidant.

• Two Methods: The Half-Reaction Method

  1. Break the reaction into oxidation and reduction half-reactions.

  2. Balance each half-reaction separately.

  3. Combine the balanced half-reactions, ensuring the number of electrons cancels out.

Example: Balancing Redox Reaction

• Unbalanced reaction: HgCl2 + Fe \rightarrow Hg + FeCl3

• Reduction: HgCl_2 \rightarrow Hg + Cl^-

• Oxidation: Fe \rightarrow FeCl_3

• Balanced half-reactions:

  • Reduction: 3HgCl_2 + 6e^- \rightarrow 3Hg + 6Cl^-

  • Oxidation: 2Fe + 6Cl^- \rightarrow 2FeCl_3 + 6e^-

• Overall balanced reaction:

  • 3HgCl2 + 2Fe \rightarrow 3Hg + 2FeCl3

Balancing Redox Reactions in Aqueous Solution (More Complex)

• Involves H_2O, H^+, and OH^- ions.

1. Divide the reaction into half-reactions; assign oxidation states.

2. Balance atoms and charge in each half-reaction.

   • Balance atoms other than O or H first.

   • Balance O next (Acidic: add H_2O; Basic: add OH^-

   • Balance H next (Acidic: add H^+; Basic: add H_2O).

   • Balance charge by adding electrons (e^-).

3. Multiply half-reactions to balance electrons.

4. Add the two half-reactions together.

Example: Balancing Redox Reaction in Acidic Conditions

• Unbalanced reaction: MnO4^- (aq) + Bi^{3+} (aq) \rightarrow Mn^{2+} (aq) + BiO3^- (aq)

• Reduction half-reaction: MnO_4^- \rightarrow Mn^{2+}

• Oxidation half-reaction: Bi^{3+} \rightarrow BiO_3^-

• Balanced half-reactions (acidic conditions):

  • Reduction: 5e^- + 8H^+ + MnO4^- \rightarrow Mn^{2+} + 4H2O

  • Oxidation: 3H2O + Bi^{3+} \rightarrow BiO3^- + 6H^+ + 2e^-

• Multiply to balance electrons:

  • Reduction (x2): 10e^- + 16H^+ + 2MnO4^- \rightarrow 2Mn^{2+} + 8H2O

  • Oxidation (x5):15H2O + 5Bi^{3+} \rightarrow 5BiO3^- + 30H^+ + 10e^-

• Overall balanced reaction:

  • 14H2O + 2MnO4^- (aq) + 5Bi^{3+} (aq) \rightarrow 2Mn^{2+} (aq) + 5BiO_3^- + 16H^+ (aq)$$