Intro to Thermochemistry and First Law

Introduction to Thermochemistry

• Thermochemistry: The study of the heat change in chemical reactions.

• Switch from theoretical chemistry to quantitative studies regarding thermochemistry and the first law of thermodynamics.

Definitions

Energy

• Definition: Capacity to do work (work = directed energy change)

• Types of Energy:

  • Kinetic Energy: Energy of motion.

  • Potential Energy: Energy associated with position.

  • Chemical Energy: A form of potential energy, related to positions of nuclei and electrons in atoms.

    • Example: Chemical energy in gasoline is converted to mechanical energy when a car engine operates.

Heat vs Thermal Energy

• Heat (q): The transfer of thermal energy between two bodies at different temperatures.

• Thermal Energy: The total energy due to random motion of molecules; an extensive property (depends on mass/volume).

• Temperature: An intensive property, a measure of the average kinetic energy of molecules.

Thermal Equilibrium

• Example: In a swimming pool at 80°F, air temperature is also 80°F.

• At this point, no net energy transfer occurs between the pool and the air (thermal equilibrium reached).

• When water is in a bucket versus the pool, both can be at 80°F, but the pool contains more thermal energy.

• If both were frozen, the bucket would freeze faster due to lesser thermal energy.

Systems and Surroundings

System Definition

• System: The part of the universe being studied; everything else is the surroundings.

• Types of Systems:

  • Open System: Mass and energy can transfer between system and surroundings (e.g. water evaporating).

  • Closed System: Mass cannot escape; energy can enter/leave (e.g. sealed flask of water).

  • Isolated System: Neither matter nor energy can escape (e.g. thermally insulated system).

First Law of Thermodynamics

• The first law states: Energy is conserved. [ \Delta E_{universe} = \Delta E_{system} + \Delta E_{surroundings} = 0 ]

• This means you can convert energy from one form to another, but cannot create or destroy it.

Exothermic and Endothermic Processes

Exothermic Processes

• Definition: System releases heat to surroundings.

  • Surroundings feel hotter; the system feels cooler.

  • Examples: Combustion and freezing.

Endothermic Processes

• Definition: System absorbs heat from surroundings.

  • Surroundings feel cooler; system feels hotter.

  • Examples: Melting and boiling.

Internal Energy

• Internal energy (E or U): Total energy inside the system, dependent on heat and work.

• ΔU = q + w

• ΔU is a state function; q and w are not state functions (dependent on the process).

State Functions

• Definition: Variables or properties that are path independent (e.g. potential energy).

• Example: Different paths taken by hikers result in the same change in potential energy, but work and heat can vary based on the path taken.

Sign Conventions

Heat (q)

• Positive when heat is absorbed (endothermic).

• Negative when heat is released (exothermic).

Work (w)

• Positive when work is done on the system (compression).

• Negative when work is done by the system (expansion).

PV Work

• Formula: [ w = -P \Delta V ]

  • Related to pressure and volume changes in a gas.

• Important for internal combustion engines and other systems where gases are under pressure.

First Law of Thermodynamics in Practice

• Matrix for energy transfers in systems.

• Measurements typically only apply to the system.

• Work done by the system depends on pressure and volume changes.

• Example Problems:

  1. Gas cools, losing 65 Joules of heat (q = -65 J), contracts with 22 Joules of work done on it (w = +22 J).

     • Calculate ΔE: [ \Delta E = -65 + 22 = -43 J ]

  2. Gas gains 50 Joules of heat (q = +50 J) and expands from 1 L to 3 L against 1 atm pressure.

     • Calculate work and ΔE after converting units properly.

Conclusion

• Understanding the first law of thermodynamics and its principles is critical for future lessons and examinations in thermochemistry and chemical processes.