UAB Thermochemistry Notes

Module 4: Thermochemistry - UAB Notes

Learning Outcomes

  • Describe the First Law of Thermodynamics and its relationship to work and energy.
  • Summarize energy conversion from potential to kinetic.
  • Compare open, closed, and isolated systems.
  • Calculate energy changes in a chemical reaction.
  • Discuss how enthalpy quantifies chemical reactions.
  • Define enthalpies of formation and write balanced chemical equations.
  • Use enthalpies of formation to calculate heats of reaction.
  • Apply Hess’ Law to calculate reaction enthalpies.
  • Perform calculations based on Calorimetry.

Thermochemistry Overview

  • Study of energy related to chemical reactions.
    • Usually involves release of energy as heat/light.
  • Energy: Capacity to do work.
    • Chemical reactions can either release or absorb energy.

Energy Forms

  • Thermal Energy: Caused by atomic/molecular motion; higher motion = greater thermal energy.
  • Radiant Energy: Carried by light (e.g., microwave, infrared, X-ray).
  • Electrical Energy: From flow of charged particles.
  • Nuclear Energy: Stored within atomic nuclei.
  • Chemical Energy: Based on atomic arrangements in molecules.
  • Energy can convert between forms.

Systems in Thermodynamics

  • System: Area of interest in thermodynamics.
  • Surroundings: Everything outside the system.
  • The combination of the system and surroundings is the universe.
Types of Systems
  • Open System: Can exchange matter & energy with surroundings.
  • Closed System: Can exchange energy, but not matter.
  • Isolated System: No exchange of energy or matter with surroundings.

First Law of Thermodynamics

  • Total energy in the universe is constant: Energy can only change forms, never created or destroyed.
  • Mathematically expressed as: [ \Delta E = q + w ]
    • where (E) is energy, (q) is heat, (w) is work.
  • Internal Energy (ΔE): ( \Delta E = E{final} - E{initial} )
    • The total internal energy of an isolated system is constant.
Energy Components
  • Internal Energy = sum of kinetic & potential energy in a system.
  • Units of Energy:
    • 1 Joule (J) = 1 kg m²/s²
    • 1 Calorie (cal) = energy to raise 1.00 g of water by 1°C (1 cal = 4.184 J).

Heat Calculations

  • Heat transfer (q) calculated as: [ q = mc\Delta T ]
    • where (m) = mass, (c) = specific heat, (\Delta T) = temperature change.
Calorimetry
  • Measurement of heat flow during reactions using calorimetry.
  • Based on First Law of Thermodynamics: [ q{sol} = -q{rxn} ]
    • (q) can be estimated using:
      [ q = m * C * \Delta T ]
  • Example problems involve measuring temperature changes during reactions.

Enthalpy

  • Enthalpy (H): Heat flow under constant pressure.
  • Enthalpy change: ( \Delta H = H{final} - H{initial} )
  • Reactions categorized:
    • Exothermic: Energy released (( \Delta H < 0 ))
    • Endothermic: Energy absorbed (( \Delta H > 0 ))
  • Dependent on states of reactants/products.
Enthalpies of Formation
  • ΔHf°: Energy required/released during the formation of a compound from its elements.
  • Elements in standard state have ΔHf° = 0.
  • Calculating using: [ \Delta H{rxn} = \Sigma n H{f, products} - \Sigma n H_{f, reactants} ]
    • where coefficients are moles for specific reactions.

Hess's Law

  • Hess's Law: Total enthalpy change is the sum of all steps in a reaction pathway.
  • Useful for calculating reaction enthalpy using tabulated values; can combine multiple reactions to find overall change.

Conclusion

  • Thermochemistry provides a framework for understanding energy in chemical processes, allowing for energy changes to be quantified and predicted in various systems and reactions.