Specific Heat and Thermochemistry

Specific Heat

  • Definition: Specific heat (c) is the amount of energy required to raise the temperature of 1 gram of material by 1 degree Celsius.
    • Units: J/(g·°C)

Heat Transfer

  • Key Calculation: The heat transfer for a substance can be calculated with the formula:

    • Q = m × c × ΔT
      • Q: heat (J)
      • m: mass (g)
      • ΔT: change in temperature (°C)
      • c: specific heat (J/g·°C)
    • Where ΔT = Tf - Ti.

  • Example: To understand the time taken to heat materials:

    • For 50g each of Aluminum (Al) and Copper (Cu):
      • Al has a higher specific heat than Cu, so it takes longer to heat to 100°C and will also take longer to cool down.

Practice Calculations

Cooling of a Silver Spoon

  • Problem: A 32-g silver spoon cools from 60°C to 20°C. How much heat is lost?
    • Given:
      • m = 32 g
      • Ti = 60°C
      • Tf = 20°C
      • c = 0.235 J/g·°C
    • Solution:
      • ΔT = Tf - Ti = 20°C - 60°C = -40°C
      • Q = m · c · ΔT
      • Q = (32 g)(0.235 J/g·°C)(-40°C)
      • Q = -300.8 J (heat lost)

Warming Water

  • Problem: How much heat is required to warm 230 g of water from 12°C to 90°C?
    • Given:
      • m = 230 g
      • Ti = 12°C
      • Tf = 90°C
      • c = 4184 J/g·°C
    • Solution:
      • ΔT = Tf - Ti = 90°C - 12°C = 78°C
      • Q = m · c · ΔT
      • Q = (230 g)(78°C)(4.184 J/g·°C)
      • Q = 75,061 J

Thermochemistry

  • Definition: The study of heat released or required by chemical reactions and phase changes.
    • Example: Combustion of fossil fuels (e.g. CH4 + 2O2 → CO2 + 2H2O)

Enthalpy (∆H)

  • Definition: Change in heat content of a system, often measured as the heat change in the surroundings during a reaction.

Systems and Surroundings in Thermodynamics

  • System: The part of the universe we study (e.g. a chemical reaction).
  • Surroundings: Everything else outside the system.

Types of Systems

  • Open System: Exchanges both matter and energy with surroundings (e.g., open reaction flask).
  • Closed System: Exchanges only energy (fixed matter) (e.g., sealed reaction flask).
  • Isolated System: Exchanges neither energy nor matter (e.g., a thermos flask).

Measuring Heat in Reactions

  • Exothermic Reaction: Heat is given off; temperature of surroundings rises.
  • Endothermic Reaction: Heat is absorbed; temperature of surroundings drops.

Reaction Enthalpies

  • Exothermic Reactions: Reactants → Products + energy released (-ΔH).
  • Endothermic Reactions: Reactants + energy → Products (+ΔH).
    • Examples:
      • Exothermic: Burning fossil fuels.
      • Endothermic: Photosynthesis.

Heat of Phase Changes

  • Heat of Vaporization (ΔHvap): Energy required to convert a liquid into vapor; an endothermic process (always positive).
  • Heat of Fusion (ΔHfus): Energy required to change a solid to a liquid; also endothermic (always positive).
    • Freezing is exothermic (ΔH negative).

Calorimetry

  • Calorimeter: Device used to measure changes in thermal energy.
    • Coffee cup calorimeter: In an insulated system, heat gained = heat lost.

Energy Transformations

  • Chemical to Thermal: Chemical reactions (like combustion) transform chemical energy into light and heat.
  • Automobile Example: Gasoline combustion transforms chemical energy into mechanical energy that drives pistons, generating electricity.

Understanding Thermal Energy

  • Definition: Total potential and kinetic energy in a system.

Endothermic and Exothermic Examples

  • Endothermic:

    • Melting a candle
    • Dew formation (condensation)
    • Breaking open a cool pack

  • Exothermic:

    • Combining metal with acid (test tube gets hot).