Energy Changes in Electrochemistry

UNIT FOUR: ENERGY CHANGES IN ELECTROCHEMISTRY

Key Concepts

• Chemical Reactions: Processes that result in a chemical change and produce energy, e.g., combustion of fuels like carbon:
  • $C(s) + O<em>2(g) \rightarrow CO</em>2(g)$
  • $CH<em>4(g) + O</em>2(g) \rightarrow CO<em>2(g) + 2H</em>2O(g)$
• Energy Change (ΔE): Involvement in chemical reactions, particularly in combustion.
• System vs Surroundings:
  • System: The part being studied.
  • Surroundings: Everything else.
• Internal Energy (E): Total energy (potential + kinetic) within a system.
  • Formula: $E = E<em>{potential} + E</em>{kinetic}$

Energy Exchange in the Universe

• The total energy of the universe remains constant.
  • When a system loses energy, the surroundings gain it:
  • $\text{Total Energy} = \text{System} + \text{Surroundings}$
• Forms of Energy Transfer: Heat (q) and Work (w):
  • $\Delta E = q + w$

Types of Chemical Reactions

Exothermic Reactions

• Energy is released to the surroundings (Heat out).
• Examples:
  • Dissolution of NaOH in water:
    $NaOH(s) \rightarrow Na^+(aq) + OH^-(aq) + \text{Heat}$
  • Dissolution of CaCl2 in water:
    $CaCl<em>2(s) + 2H</em>2O(l) \rightarrow Ca(OH)_2(aq) + 2HCl(g) + \text{Heat}$
• Enthalpy Change (ΔH < 0): Heat of products < Heat of reactants.

Endothermic Reactions

• Energy is absorbed from the surroundings (Heat in).
• Examples:
  • Dissolution of KNO3 in water:
    $KNO<em>3(s) + H</em>2O + \text{Heat} \rightarrow K^+(aq) + NO_3^-(aq)$
• Enthalpy Change (ΔH > 0): Heat of products > Heat of reactants.

Measuring Energy Changes

• Calorimeter: Device to measure heat released or absorbed during reactions.
  • $q \propto \Delta T$ (where q is heat, and ΔT is temperature change)
  • Heat Capacity (C): The heat required to change temperature by 1 K, with units Joule per Kelvin (J/K).
• Specific Heat Capacity (c): Defined as:
  • $c = \frac{Heat : Capacity}{mass} = \frac{q}{mass \times \Delta T}$
  • Relation: $q = c \times mass \times \Delta T$

Chemical Changes Importance

• Chemical reactions: Key events for life processes (photosynthesis, digestion).
• Energy production: Essential for various biological processes.

Types of Conductivity

• Electrical Conductivity: Ability to transmit electricity (e.g., in solutions of ionic compounds).
  • Types:
  1. Metallic Conductivity: Movement of electrons (conductors like cooper, silver).
  2. Electrolytic Conductivity: Movement of ions in electrolytic solutions.
     • Strong electrolytes (like NaCl) dissociate completely in water.

Electrochemical Cells

• Definition: Devices that convert electrical energy to chemical energy and vice versa.
• Electrodes: Can be active (e.g., Zinc) or inert (e.g., Platinum).

Types of Electrochemical Cells

1. Galvanic Cells: Convert chemical energy to electrical energy (spontaneous redox).
   • Example: Daniell cell:
     • Anode (negative): $Zn \rightarrow Zn^{2+} + 2e^-$ (oxidation)
     • Cathode (positive): $Cu^{2+} + 2e^- \rightarrow Cu$ (reduction)
     • Overall reaction: $Zn(s) + Cu^{2+} \rightarrow Zn^{2+} + Cu(s)$
2. Electrolytic Cells: Convert electrical energy to chemical energy (non-spontaneous).
   • Oxidation occurs at the anode (+), reduction occurs at the cathode (-).

Additional Concepts

• Electrolysis: Techniquefor obtaining elements from their compounds using electrical energy.
• Secondary Cells: Rechargeable batteries (e.g., lead-storage battery) with reversible reactions.
• Non-electrolytes: Substances that do not conduct electricity (e.g., sugar, alcohol).

Summary of Key Processes & Terms

• Galvanic Cell: Converts chemical to electrical energy, spontaneous reactions.
• Electrolytic Cell: Converts electrical to chemical energy, non-spontaneous reactions.
• Heat Capacity: Amount of heat to change temperature by 1 K.
• Specific Heat Capacity: Heat capacity per mass unit.
• Endothermic and Exothermic Reactions: Defined by energy absorbed or released to/from the surroundings.