Concepts of Thermodynamics I Flashcards
Forms of energy
1) Potential energy - any form of stored energy
2) Kinetic energy - energy of moving objects
What is a System?
System in a flaks, flask is a boundary and surroundings is the enviroment outside
System + surroundings = universe
Types of systems
- Isolated system
- Closed system
- Open system
Zeroth Law of Thermodynamics
If all are in equilibrium energy can be transferred
Energy transferred from B to C through A
Different pathways
First Law of Thermodynamics
“Energy cannot be created or destroyed – it can only be transformed from one form to another”
Law of Conservation of Energy
“The change in internal energy of a closed system will be equal to the heat added or released by the system minus the work done by the system on its surrounding or the work done on the system by its surroundings”
Equation to first law
AU = Q - W
AU - change in internal energy
Q - Heat added to the system
W - work done by the system
Q positive
Heat added to the system
ENDOTHERMIC process
Q negative
Heat released from the system
ENDOTHERMIC process
Constant Volume (Isovolumeric)
∆U = q
No Work is done on or by the system
Any energy change is result of heat transfer
Constant Temperature (Isothermal)
q = w
No change in internal energy
Any heat transferred to the system is used to do work
No Heat Transfer (Adiabatic)
Internal energy changes only for work done or received by the system
Energy and Heat Capacity
ratio of heat absorbed by a material to the temperature change
What are the units of energy?
J - joules (SI units)
1 cal
heat needed to raise the temperature of 1 gram of water 1⁰C at standard pressure
Conversion of J to cal
1 cal = 4.184 J, 1 kcal = 4184 J
Heat Capacity (C):
amount of heat required to raise a substance’s temperature by 1 Kelvin
Specific Heat Capacity (Cₛ):
amount of heat required to raise 1 gram of a substance by 1 Kelvin
What is Enthalpy?
Enthalpy is the measurement of energy in a thermodynamic system (joules)
Equation of enthalpy
∆H = ∆U + ∆(PV)
What does each letter mean?
∆H = ∆U + ∆(PV)
P = Pressure, V = Volume
∆H = ∆U and ∆U = q, therefore ∆H = q
q > 0, so ∆H is positive
Endothermic process
q < 0, so ∆H is negative
Exothermic process
2nd Law of Thermodynamics
“The Universe is always moving towards maximum disorder”
Entropy (s)
Measure of disorder
Breakdown of units of 2nd Law of Thermodynamics
s = KB ln W
KB = Boltzmann constant (energy of an individual atom)
W = microstates (ways to arrange particles in a system)
∆s = sf - si
Increase in the number of micro-states
∆s = sf – si
then Wf > Wi and the entropy of the system increases ∆S > 0
Decrease in the number of micro-states
then Wf < Wi and the entropy of the system decreases ∆S < 0
More micro-states
More disorder
Two Bulb Experiment:
Spontaneous process where gas expands to fill
both bulbs equally, increasing entropy
More space for gas to go to…
More disorder
Wf < Wi
decrease in number of micro states
f = final
i = initial
2nd law - Increase Temperature:
Entropy will be greater as molecules have more kinetic energy and therefore more dispersion
Change state:
Gas > Liquid > Solid – more dispersion
ENTROPY increase - particles are more free
Mixing of particle types:
Increases entropy therefore dissolution increases entropy
3rd Law of Thermodynamics
“At absolute zero (0 K), the entropy of a perfect, crystalline substance is zero”
Why does entropy equal zero?
3rd Law
All vibrations and atomic movements stop
What temperature is absolute zero in ⁰C?
Answer: -273.15 ⁰C
Comes from the Triple Point of water, T3 = 273.15K
CO2 (s) > CO2 (g)
Entropy increase
Solid to gas
If there is more molecules of gas than the other side of equation…
the one with more has the higher temperature
