Thermodynamics
The relationships between heat, work, and internal energy.
Thermodynamics governs how thermal systems perform mechanical work. It is defined by fundamental laws describing energy conservation and natural direction of heat transfer.
This unit covers internal energy changes, ideal gas processes (isochoric, isobaric, isothermal, adiabatic), heat engines, and Carnot efficiency limits.
Key Takeaways
- •First Law (Energy Conservation): Change in internal energy equals heat added minus work done by the system.
- •Second Law (Entropy): Heat cannot spontaneously flow from colder to hotter bodies; entropy of closed systems always increases.
- •No heat engine can ever achieve 100% thermal efficiency.
Core Concepts & Definitions
1First Law of Thermodynamics
Statement of conservation of energy for thermal systems: ΔU = Q - W.
•ΔU is change in internal energy (temperature dependent).
•Q is heat added to system; W is work done by system on surroundings.
2Thermodynamic Gas Processes
Changes in pressure, volume, and temperature of an ideal gas.
•Isothermal: constant temperature (ΔT = 0, ΔU = 0).
•Isobaric: constant pressure (W = P * ΔV).
•Isochoric: constant volume (no work done, W = 0).
•Adiabatic: no heat exchange (Q = 0, ΔU = -W).
Quick Revision Notes
- •Ensure all temperatures are converted to Kelvin (K = °C + 273.15) before calculating efficiencies.
- •In adiabatic expansion, gas cools down because it performs work at the expense of its internal energy (Q=0, ΔU = -W).