Disclaimer:
Mistakes on this review sheet are possible but NOT intentional. You should ALWAYS crosscheck all information.
Thermodynamics: the study of heat transfer and work
First Law of Thermodynamics:
Energy can be neither created nor destroyed – but can be converted from one form to another or
passed from one object to another.
Define a system and consider both it and its surroundings (the whole rest of the universe)
The universe = the system + the surroundings
There are different types of systems, know what you are working with:
Open: exchange energy & matter (Example: hot air balloon)
Closed: exchange energy only (Example: sealed ice pack)
Isolated:
NO exchange of energy or matter (Example: Thermos flask)
Newtonian Physics: the total energy = potential energy plus kinetic energy.
Often they are interested in potential energy due to gravity, PE = mgh (where the gravitational
constant g = 9.81 m/s
2
) and kinetic energy, KE =
½
mv
2
For chemical thermodynamics the internal energy of the system, U or E, (two letters, used to
represent the same thing) includes both heat energy and work energy.
Because we cannot measure or calculate the absolute energy of the system we talk about changes in
the internal energy of the system.
Δ
E or
Δ
U = q + w
State functions do not depend on the path taken to reach its particular “state of being” – They
describe how the system is. Examples: internal energy, E, enthalpy, H, pressure, P, Volume, V,
entropy, S, All state functions are (should be) represented by capital letters.
work and heat are path-dependant and thus are not state functions: w = -P
Δ
V =
Δ
nRT
notes: when calculating
Δ
n only consider the number of moles of gas (condensed phases are small
and minimally compressible – all the action is in the gases). Also: here we see,
l
.
atm is a unit of
energy. One can reconstruct the conversion factor to joules, by dividing two versions of the ideal
gas constant: 1 = R(J) / R(
l
.
atm) = 8.314 J/mole
.
K / 0.08206
l
.
atm/mole
.
K = 101.3 J/
l
.
atm
Isobaric (constant pressure) conditions: Whenever we work in open containers the pressure is
constant and the same as the pressure of the surroundings. At constant pressure, it is best to
consider the enthalpy of a system, H, which is the internal energy plus the product P
.
V:
H = E + PV
The change in enthalpy is
Δ
H =
Δ
E + P
Δ
V = q + w + P
Δ
V
When only PV work is allowed,
w = -P
ext
Δ
V,
thus
Δ
H = q “heat of (reaction or process)”
Absolute enthalpy can not be known so arbitrary zero point for enthalpy values,
Δ
H