Electrochemistry notes

Electrochemistry Overview

• Stored Electrical Energy

  • Common form is batteries.

  • Involves electrochemical reactions.

  • These are called REDOX reactions.

Oxidation-Reduction (Redox) Reactions

• Definition: Reactions in which electrons are transferred from one atom to another.

  • Atoms losing electrons are oxidized.

  • Atoms gaining electrons are reduced.

  • Example Reaction:

  • $2 ext{Na}(s) + ext{Cl}_2(g) ightarrow 2 ext{Na}^+ + 2 ext{Cl}^- (s)$

    • Sodium (Na) undergoes oxidation:

      • $ext{Na} <br>ightarrow ext{Na}^+ + 1 e^-$

    • Chlorine (Cl) undergoes reduction:

      • $ext{Cl}_2 + 2 e^- <br>ightarrow 2 ext{Cl}^-$

Identifying Redox Reactions

• The easiest way to identify a redox reaction is by checking if oxidation numbers change from reactants to products.

  • Example Reaction:

  • $ext{Mg} + 2 ext{HCl} <br>ightarrow ext{MgCl}_2 + ext{H}_2$

  • Change in oxidation numbers indicates a redox reaction.

Oxidation States

• Determining Oxidation States: A set of rules is used:

  1. The oxidation state of an atom in its elemental form is 0.

     • Examples: Silver (Ag) metal, Hydrogen (H2) gas.

  2. The oxidation state of a monotonic ion is its ionic charge.

     • Example: $ext{Cl}^- ext{ (ox. state -1)}, ext{Mg}^{2+} ext{ (ox. state +2)}$

  3. The sum of oxidation states in a molecule or ion equals the overall charge.

     • Example: In $ext{H}_2 ext{O}$: $2( ext{ox. state of H}) + ( ext{ox. state of O}) = 0$

     • Example: In $ext{SO}_4^{2-}$: $( ext{ox. state of S}) + 4( ext{ox. state of O}) = -2$

  4. Group 1A metals always have a +1 oxidation state, and Group 2A metals have a +2 oxidation state.

     • Example: $ext{NaCl (Na has +1 oxidation state)}, ext{MgCl}_2 ( ext{Mg has +2 oxidation state})$

  5. Oxidation states for nonmetals are assigned according to a specific table. Prioritize higher entries in case of conflict.

Identifying Changes in Oxidation States

• OIL RIG:

  • Oxidation is Losing (electrons)

  • Reduction is Gaining (electrons)

• Example: Reaction: $ext{Mg} + 2 ext{HCl} ightarrow ext{MgCl}_2 + ext{H}_2$

  • Mg loses electrons (oxidation) while hydrogens gain electrons (reduction).

  • Oxidation Numbers: H=+1, Cl=-1, Mg=0.

Oxidation and Reduction Definitions

• Oxidation occurs when:

  • The oxidation number increases.

  • An atom loses electrons.

  • A compound adds oxygen.

  • A compound loses hydrogen.

  • A half-reaction has electrons as products.

• Reduction occurs when:

  • The oxidation number decreases.

  • An atom gains electrons.

  • A compound loses oxygen.

  • A compound gains hydrogen.

  • A half-reaction has electrons as reactants.

Half-Reactions in Redox Reactions

• Oxidation and reduction processes occur as pairs (simultaneously).

• Example:

  • $ext{Zn} + ext{Cu}^{2+} <br>ightarrow ext{Zn}^{2+} + ext{Cu}$

  • Oxidation Half-Reaction:

    • Zinc loses two electrons:

      • $ext{Zn} <br>ightarrow ext{Zn}^{2+} + 2e^-$

  • Reduction Half-Reaction:

    • Copper gains two electrons:

      • $ext{Cu}^{2+} + 2e^- <br>ightarrow ext{Cu}$

Voltaic (Galvanic) Cells

• Spontaneous electrochemical reactions where electrons flow when the circuit is complete.

• Galvanic Cell Components:

  • Anode:

  • Where oxidation occurs. Example: $ext{Zn} ext{(s)} <br>ightarrow ext{Zn}^{2+}$

  • Cathode:

  • Where reduction occurs. Example: $ext{Cu}^{2+} + 2e^- <br>ightarrow ext{Cu (s)}$

  • Cell Notation:

  • $ext{Zn(s)} | ext{Zn}^{2+}(aq) || ext{Cu}^{2+}(aq) | ext{Cu(s)}$

Salt Bridge in Batteries

• The salt bridge:

  • Completes the circuit.

  • Provides a constant flow of charge.

  • Balances the charges in the solutions.

Balancing Redox Reactions

• Redox reactions must balance mass and charge accounting for electrons.

• Example Balancing Steps:

  1. Write half-reaction forms.

  2. Balance elements other than O and H.

  3. Use $ext{H}_2 ext{O}$ to balance oxygen and $ext{H}^+$ to balance hydrogen.

  4. Use electrons to balance charge on each side.

  5. Multiply half-reactions to have the same number of electrons.

  6. Add reactions cancelling identical species.

  7. Verify that both atoms and charges balance.

Redox in Acidic Solutions

• Example:

  • $ext{Fe}^{2+} + ext{Cr}_2 ext{O}_7^{2-} <br>ightarrow ext{Fe}^{3+} + ext{Cr}^{3+}$

• Remember to balance using electrons and ensuring that each charge and atom is accounted for.

Electrolytic Cells

• Involves non-spontaneous reactions driven by external energy.

• Oxidation still occurs at the anode, but it has a positive charge.

• Current passed dictates how much products are formed (quantitative electrolysis).

Electroplating Example

• Calculation:

  1. Calculate charge: $Q = I imes t$

  2. Calculate moles of electrons: $ext{moles} = rac{Q}{F}$

  3. Calculate mass deposited using molar mass of copper and stoichiometry.

• Example: Current of 2.40 A through Cu²⁺ for 30 minutes results in 1.42 g of copper deposited.

Applications of Electrochemistry

• Corrosion:

  • Occurs in moisture; iron oxidizes to rust (Fe₂O₃.xH₂O).

  • Salt accelerates the process, especially in water and on roads.

  • Aluminum forms a protective layer, preventing rusting when anodized.

Standard Cell Potentials and Nernst Equation

• Learn how to calculate cell potential and predict spontaneity of reactions.

• Utilize the Nernst Equation for non-standard conditions to adjust for concentration and temperature variations.

Comprehension Checks

• Important to periodically verify understanding through exercises related to the material covered.

• Utilize example problems and reactions to ensure competency in balancing and predicting electrochemical processes.