1. THE STATES OF GASES
Chapter 1

2. Bulk Variables Volume - m3 Pressure - Pa Temperature - K
Composition - moles

3. Volume length3 units molar volume Vm= V/n m3 or cm3 liter = 1000 cm3

4. Pressure Force/Area units Pascal = Newton/m2 = Joule/m3
atmosphere = Pa
bar = Pa – Standard Pressure
mm Hg
torr h

5. Mechanical Equilibrium
High P
Low P
High pressure gas will tend to compress a low pressure gas
Movable wall or
“piston”
Equal P

6. Temperature Thermometry - relate temperature to other properties.
Fahrenheit (°F)
Celsius (°C)
Kelvin (K)

7. Thermal Equilibrium
High T
Low T
Heat or energy will flow from the high T gas to the low T gas
Rigid diathermic wall
Equal T

8. Zeroth Law of Thermodynamics
Two systems that are separately in thermal equilibrium with a third system are also in thermal equilibrium with one another.
The zeroth law is the basis for thermometry, i.e. the use of a third body (the thermometer) to measure an equilibrium property (the temperature) of other systems.

9. Zeroth Law of Thermodynamics
Thermal
Equilibrium
Thermal
Equilibrium B C
Thermal
Equilibrium

10. CHARLES’S LAW

11. Absolute zero = °C
T/K = θ/°C
CHARLES’S LAW

12. Composition moles: ni S ni = n mole fraction: xi S xi = 1
partial pressure: pi S pi = p

13. Gas Laws
Boyle’s Law: pV = constant at constant n, T

14. Equations of State P,V,T, and n are not independent.
Any three will determine the fourth.
An equation of state is an equation that relates P,V,T, and n for a given substance.
Gases have the simplest equations of state.
The simplest equation of state is the ideal gas law, pV = nRT

15. Ideal Gas Law p = pressure V = volume n = moles T = temperature
R = universal gas constant = L-atm/mol-K (Table 1.2)

16. Ideal Gas Model
Molecules may be treated as point masses relative to the volume of the system.
Molecular collisions are elastic, i.e. kinetic energy is conserved.
Intermolecular forces of attraction and repulsion have negligible on the molecular motion.
Real gases approximately behave as ideal gases at higher temperatures and low pressures.

17. Ideal Gas Model

18. Ideal Gas Problem
Gas in a vessel of constant volume is heated to 600 K. If it initially entered the vessel at 300 K and 1 atm, what is the pressure after heating?

19. Ideal Gas Problem
Gas in a vessel of constant volume is heated to 600 K. If it initially entered the vessel at 300 K and 1 atm, what is the pressure after heating?

20. Ideal Gas Problem
Gas in a vessel of constant volume is heated to 600 K. If it initially entered the vessel at 300 K and 1 atm, what is the pressure after heating?

21. Partial Pressure pA = partial pressure of gas A V = total volume
nA = moles of gas A
T = temperature
R = universal gas contant = L-atm/mol-K

22. Dalton’s Law
The total pressure is the sum of all the partial pressure.

23. Dalton’s Law Problem
Earth’s atmosphere is ~ 75.5% N2, 23.2 % O2, and 1.3 % Ar by mass. What is the partial pressure of each component when the total pressure is 1.00 atm?

24. Dalton’s Law Problem
Earth’s atmosphere is ~ 75.5% N2, 23.2 % O2, and 1.3 % Ar by mass. What is the partial pressure of each component when the total pressure is 1.00 atm?
STOP HERE

25. Real Gases Real gases do not obey the ideal gas law.
Deviations at low temperature and high pressure.
Especially near point of condensation.

28. Molecular Interactions
Real gases are not point masses – they have volume.
Real gases do interact.
Two forces of interaction
Attractive forces
Repulsive forces.

29. Molecular Interactions
Attractive forces are “long” range forces
Long = several molecular diameters
Beyond this attractive forces are not significant
“Low” temperature
Low = T near condensation point.

30. Molecular Interactions
Repulsive forces are “short” range forces
Short = ~ less than 1 molecular diameter
At low T, attractive forces overcomes repulsive forces.

31. Intermolecular Forces

32. Summary
At low pressures and large volumes, there is little interaction ~ Ideal Gas.
At moderate pressure, attractive forces start to dominate – gas more compressible.
At high pressure, repulsive forces dominate – gas less compressible.

33. Compressibility
The compressibility of a gas is defined by:

34. Compressibility
If the gas behaves ideally, then Z=1 at all pressures and temperatures.
For real gases, however, Z varies with pressure, and deviates from its ideal value.

35. Compressibility

36. Isotherms, isobars and isochores

37. Ideal Gas Isotherm