Lecture 3 Examples and Problems - University Of Illinois
Lecture 3 Examples and Problems
Mechanics & thermodynamics Equipartition First Law of Thermodynamics Ideal gases Isothermal and adiabatic processes
Reading: Elements Ch. 1-3
Lecture 3, p 1
William Thomson (1824 ? 1907) a.k.a. "Lord Kelvin"
First wrote down Second Law of Thermodynamics (1852) Became Professor at University of Glasgow at age 22!
(not age 1.1 x 1021)
Physics 213: Lecture 3, Pg 2
Ideal Gas p-V, p-T Diagrams
p vs V at various constant T's
p = NkT V
increasing T
Isotherms
p vs T at constant V
Pressure
Pressure
Volume
For an ideal gas at constant T, p is inversely proportional to the volume.
0
0
Temperature
Pressure zero as T absolute zero, because the thermal kinetic energy of the molecules vanishes.
Lecture 3, p 3
Last time: The First Law of Thermodynamics
Energy is conserved !!!
U = Q + Won
change in total internal energy
heat added to system
work done on the system
alternatively:
U = Q - Wby
Note: For the rest of the course, unless explicitly stated, we will
ignore KECM, and only consider internal energy that does not contribute to the motion of the system as a whole.
Lecture 3, p 4
Heat Capacity
Look at Q = U + Wby
If we add heat to a system, there are two general destinations for the energy: ? It will "heat up" the system (i.e., raise T). ? It can make the system do work on the surroundings.
Heat capacity is defined to be the heat required to raise the temperature of a system by 1K (=1? C). Its SI units are J/K.
C
Q T
(for
small
T)
The heat capacity will depend on whether energy goes into work, instead of only increasing U. Therefore, we distinguish between:
? Heat capacity at constant volume (CV), for which W = 0. ? Heat capacity at constant pressure (Cp), for which W > 0
(most systems expand when heated).
Lecture 3, p 5
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