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: $O<em>2$ (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-}$</p> <ul> <li><p>O</p></li> <li><p>Oxidation State: O (-2), P (+5)</p></li></ul></li> <li><p>$LiAlH_4$</p> <ul> <li>Oxidation State: Li (+1), Al (+3), H (-1)</li></ul></li> </ul> <h4 id="balancingredoxreactionsinaqueoussolutions">Balancing Redox Reactions in Aqueous Solutions</h4> <ul> <li><p>Reactions must be balanced in terms of the number of atoms and types, and also in terms of charge (number of electrons).</p></li> <li><p>The number of electrons lost by the reductant must equal the number of electrons gained by the oxidant.</p></li> <li><p>Two Methods: The Half-Reaction Method</p> <ol> <li><p>Break the reaction into oxidation and reduction half-reactions.</p></li> <li><p>Balance each half-reaction separately.</p></li> <li><p>Combine the balanced half-reactions, ensuring the number of electrons cancels out.</p></li></ol></li> </ul> <h4 id="examplebalancingredoxreaction">Example: Balancing Redox Reaction</h4> <ul> <li><p>Unbalanced reaction:$HgCl2 + Fe \rightarrow Hg + FeCl3$</p></li> <li><p>Reduction:$HgCl_2 \rightarrow Hg + Cl^-$</p></li> <li><p>Oxidation:$Fe \rightarrow FeCl_3$</p></li> <li><p>Balanced half-reactions:</p> <ul> <li><p>Reduction:$3HgCl_2 + 6e^- \rightarrow 3Hg + 6Cl^-$</p></li> <li><p>Oxidation:$2Fe + 6Cl^- \rightarrow 2FeCl_3 + 6e^-$</p></li></ul></li> <li><p>Overall balanced reaction:</p> <ul> <li>$3HgCl2 + 2Fe \rightarrow 3Hg + 2FeCl3$</li></ul></li> </ul> <h4 id="balancingredoxreactionsinaqueoussolutionmorecomplex">Balancing Redox Reactions in Aqueous Solution (More Complex)</h4> <ul> <li>Involves$H_2O$,$H^+$, and$OH^-$ions.</li> </ul> <ol> <li><p>Divide the reaction into half-reactions; assign oxidation states.</p></li> <li><p>Balance atoms and charge in each half-reaction.</p> <ul> <li><p>Balance atoms other than O or H first.</p></li> <li><p>Balance O next (Acidic: add$H_2O$; Basic: add$OH^-$</p></li> <li><p>Balance H next (Acidic: add$H^+$; Basic: add$H_2O$).</p></li> <li><p>Balance charge by adding electrons ($e^-$).</p></li></ul></li> <li><p>Multiply half-reactions to balance electrons.</p></li> <li><p>Add the two half-reactions together.</p></li> </ol> <h4 id="examplebalancingredoxreactioninacidicconditions">Example: Balancing Redox Reaction in Acidic Conditions</h4> <ul> <li><p>Unbalanced reaction:$MnO4^- (aq) + Bi^{3+} (aq) \rightarrow Mn^{2+} (aq) + BiO3^- (aq)$</p></li> <li><p>Reduction half-reaction:$MnO_4^- \rightarrow Mn^{2+}$</p></li> <li><p>Oxidation half-reaction:$Bi^{3+} \rightarrow BiO_3^-$</p></li> <li><p>Balanced half-reactions (acidic conditions):</p> <ul> <li><p>Reduction:$5e^- + 8H^+ + MnO4^- \rightarrow Mn^{2+} + 4H2O$</p></li> <li><p>Oxidation:$3H2O + Bi^{3+} \rightarrow BiO3^- + 6H^+ + 2e^-$</p></li></ul></li> <li><p>Multiply to balance electrons:</p> <ul> <li><p>Reduction (x2):$10e^- + 16H^+ + 2MnO4^- \rightarrow 2Mn^{2+} + 8H2O$</p></li> <li><p>Oxidation (x5):$15H2O + 5Bi^{3+} \rightarrow 5BiO3^- + 30H^+ + 10e^-$</p></li></ul></li> <li><p>Overall balanced reaction:</p> <ul> <li>$14H2O + 2MnO4^- (aq) + 5Bi^{3+} (aq) \rightarrow 2Mn^{2+} (aq) + 5BiO_3^- + 16H^+ (aq)$$