Lecture Notes: Heat, Temperature, and Gas Laws

Concepts of Heat and Temperature

• Heating Ice:

  • Initially, temperature and kinetic energy increase as particles vibrate faster.
  • At melting point (0°C), temperature remains constant during phase change (melting).
  • Similar constant temperature behavior during boiling at 100°C.

• Change of State:

  • During a phase change, energy increases potential energy, not kinetic.
  • Use Specific Latent Heat (SLH) instead of specific heat capacity.

Specific Heat and Latent Heat

• Specific Heat Capacity (SHC):

  • Used for temperature change within the same state.
  • Equation relates change in temperature to heat energy.

• Specific Latent Heat (SLH):

  • Energy required to change the state of 1 kg of substance without changing temperature.

Kelvin and Celsius Scales

• Celsius vs. Kelvin:
  • Celsius is not absolute; Kelvin starts at absolute zero (0 K = -273°C).
  • Convert Celsius to Kelvin by adding 273.

Gas Laws

• Boyle's Law (Isothermal Change):

  • Pressure inversely proportional to volume at constant temperature: (P_1V_1 = P_2V_2).

• Charles's Law:

  • Volume directly proportional to temperature at constant pressure: (V_1/T_1 = V_2/T_2).

• Pressure Law (Gay-Lussac's Law):

  • Pressure directly proportional to temperature at constant volume.

• Combined Gas Law:

  • (PV \propto T), turns to equation with constant (PV = nRT).

Ideal Gas Assumptions - RAVED

• RANDOM: Particles move randomly.
• ATTRACTION: No attraction between particles.
• VOLUME: Particle volume is negligible.
• ELASTIC: All collisions are elastic.
• DURATION: Duration of collisions negligible.

Kinetic Theory of Gases

• Equation Derivation: Pressure, Volume, and Density:

  • Pressure derived from change in momentum and force.
  • Root Mean Square Speed (crms) is an average speed measure.

• Kinetic Energy:

  • (KE = \frac{3}{2}kT) for one molecule, where (k) is Boltzmann constant.

Thermodynamic Processes

• First Law of Thermodynamics:

  • (Q = \Delta U + W): Heat in equals change in internal energy plus work done.

• Types of Processes:

  • Adiabatic: No heat exchange, (Q = 0).
  • Isothermal: Constant temperature, (Q = W).
  • Isochoric: Constant volume, (W = 0).

• Work Done by/on Gas:

  • Positive if gas expands, negative if compressed.

Practical Applications and Problem Solving

• Problem Solving:

  • Use equations to equate total energy gained and lost in systems.
  • Consider whether processes are adiabatic, isothermal, or isochoric.

• Visualizing Processes:

  • Graphs can show compression/expansion and work done on gas.
  • Areas under curves represent work done.

These notes provide an overview of heat transfer, phase changes, gas laws, and related thermodynamic principles essential for understanding thermal physics at a higher academic level.