1. Gas Laws 1

2. Properties of Gases Variable volume and shape
Expand to occupy volume available
Volume, Pressure, Temperature, and the number of moles present are interrelated
Can be easily compressed
Exert pressure on whatever surrounds them
Easily diffuse into one another 2

3. Mercury Barometer Used to define and measure atmospheric pressure
On the average at sea level the column of mercury rises to a height of about 760 mm.
This quantity is equal to 1 atmosphere
It is also known as standard atmospheric pressure 3

4. Barometer
The mercury barometer was the basis for defining pressure, but it is difficult to use or to transport
Furthermore Mercury is very toxic and seldom used anymore
Most barometers are now aneroid barometers or electronic pressure sensors, 4

5. Pressure Units & Conversions
The above represent some of the more common units for measuring pressure. The standard SI unit is the Pascal or kilopascal. (kPa)
The US Weather Bureaus commonly report atmospheric pressures in inches of mercury.
Pounds per square inch or PSI is widely used in the United States.
Most other countries use only the metric system. 5

6. Standard Temperature and Pressure
Standard Temperature and Pressure or STP = 0oC or Kelvin and a pressure of 1 atmosphere.
This is used as a reference point when comparing quantities of gases
Gases are seldom measured at exactly these conditions.
We need to be able to compute the volume at various temperature and pressures

7. Sample Problem 1:
If the pressure of helium gas in a balloon has a volume of 4.00 dm3 at 210 kPa, what will the pressure be at 2.50 dm3?
P1 V1 = P2 V2
(210 kPa) (4.00 dm3) = P2(2.50 dm3)
P2 = (210 kPa) (4.00 dm3)
(2.50 dm3)
= 340 kPa 7

8. Boyle’s Law
According to Boyle’s Law the pressure and volume of a gas are inversely proportional at constant pressure.
PV = constant.
P1V1 = P2V2 8

9. Boyle’s Law A graph of pressure and volume gives an inverse function
A graph of pressure and the reciprocal of volume gives a straight line 9

10. Standard Temperature and Pressure (STP)
The volume of a gas varies with temperature and pressure. Therefore it is helpful to have a convenient reference point at which to compare gases.
For this purpose standard temperature and pressure are defined as:
Temperature = 0oC K
Pressure = 1 atmosphere
= torr
= kPa
This point is often called STP 10

11. Charles’ Law
According to Charles’ Law the volume of a gas is proportional to the Kelvin temperature as long as the pressure is constant
V = kT
Note: The temperature for gas laws must always be expressed
in Kelvin where Kelvin = oC (or 273 to 3 significant digits) V1 = T1 V2 T2 11

12. Charles’ Law A graph of temperature and volume yields a straight line.
Where this line crosses the x axis (x intercept) is defined as absolute zero 12

13. Sample Problem 2
A gas sample at 40 oC occupies a volume of 2.32 dm3. If the temperature is increased to 75 oC, what will be the final volume?
V1 = V2
T T2
Convert temperatures to Kelvin. 40oC = 313K
75oC = 348K
2.32 dm3 = V2
313 K K
(313K)( V2) = (2.32 dm3) (348K)
V2 =
2.58 dm3 13

14. Gay-Lussac’s Law P1 = P2 T2 T1
Gay-Lussac’s Law defines the relationship between pressure and temperature of a gas.
The pressure and temperature of a gas are directly proportional
P = P2 T2 T1 14

15. Sample Problem 3:
The pressure of a gas in a tank is 3.20 atm at 22 oC. If the temperature rises to 60oC, what will be the pressure in the tank?
P1 = P2
T T2
Convert temperatures to Kelvin. 22oC = 295K
60oC = 333K
3.20 atm = P2
295 K K
(295K)( P2) = (3.20 atm)(333K)
P2 =
3.6 atm 15

16. The Combined Gas Law
1. If the amount of the gas is constant, then Boyle’s Charles’ and Gay-Lussac’s Laws can be combined into one relationship
P1 V = P2 V2 T1 T2 16

17. Sample Problem 4:
A gas at 110 kPa and 30 oC fills a container at 2.0 dm3. If the temperature rises to 80oC and the pressure increases to 440 kPa, what is the new volume?
P1V1 = P2V2
T T2
Convert temperatures to Kelvin. 30oC = 303K
80oC = 353K
V2 = V1 P1 T2
P2 T1
= (2.0 dm3) (110 kPa ) (353K)
(440 kPa ) (303 K)
V2 = 0.58 dm3 17

18. Advogadro’s Law
Equal volumes of a gas under the same temperature and pressure contain the same number of particles.
If the temperature and pressure are constant the volume of a gas is proportional to the number of moles of gas present
V = constant * n
where n is the number of moles of gas
V/n = constant
V1/n1 = constant = V2 /n2
V1/n1 = V2 /n2 18

19. Universal Gas Equation
Based on the previous laws there are four factors that define the quantity of gas: Volume, Pressure, Kelvin Temperature, and the number of moles of gas present (n).
Putting these all together: PV nT
= Constant = R
The proportionality constant R is known as the
universal gas constant 19

20. Universal Gas Equation
The Universal gas equation is usually written as
PV = nRT
Where P = pressure
V = volume
T = Kelvin Temperature
n = number of moles
The numerical value of R depends on the pressure unit (and perhaps the energy unit) Some common values of R include:
R = dm3 torr mol-1 K-1
= dm3 atm mol-1 K-1
= dm3kPa mol-1 K-1 20

21. Sample Problem 5 Example: What volume will 25.0 g O2 occupy
at 20oC and a pressure of atmospheres? :
(25.0 g)
n = = mol
(32.0 g mol-1)
Data
Formula
Calculation
Answer
V =? P = atm; T = ( )K = 293K
R = dm-3 atm mol-1 K-1
PV = nRT so V = nRT/P
V = (0.781 mol)( dm-3 atm mol-1 K-1)(293K)
0.880 atm
V = 21.3 dm3 21

22. d is the density of the gas in g/L
Universal Gas Equation –Alternate Forms
Density (d) Calculations = PM RT
m is the mass of the gas in g m V
d =
M is the molar mass of the gas
Molar Mass (M ) of a Gaseous Substance
dRT P
d is the density of the gas in g/L
M = 22

23. Sample Problem 6 dRT m M = d = = = 2.21 V P x 0.0821 x 300.15 K M =
A 2.10 dm3 vessel contains 4.65 g of a gas at 1.00 atmospheres and 27.0oC. What is the molar mass of the gas?
dRT P
M =
d = m V
4.65 g
2.10 dm3 =
= 2.21 g
dm3
2.21 g
dm3
1 atm
x x K
dm3•atm
mol•K
M =
M =
54.6 g/mol 23

24. Dalton’s Law of Partial Pressures
The total pressure of a mixture of gases is equal to the sum of the pressures of the individual gases (partial pressures).
PT = P1 + P2 + P3 + P
where PT = total pressure
P1 = partial pressure of gas 1
P2 = partial pressure of gas 2
P3 = partial pressure of gas 3
P4 = partial pressure of gas 4 24

25. Dalton’s Law of Partial Pressures
Applies to a mixture of gases
Very useful correction when collecting gases over water since they inevitably contain some water vapor. 25

26. Sample Problem 7
Henrietta Minkelspurg generates Hydrogen gas and collected it over water.
If the volume of the gas is 250 cm3 and the barometric pressure is torr at 25oC, what is the pressure of the “dry” hydrogen gas at STP?
(PH2O = 23.8 torr at 25oC) 26

27. Sample Problem 8
Henrietta Minkelspurg generated Hydrogen gas and collects it over water. If the volume of the gas is 250 cm3 and the barometric pressure is 765 torr at 25oC, what is the volume of the “dry” oxygen gas at STP?
From the previous calculation the adjusted pressure is torr
P1= PH2 = torr; P2= Std Pressure = 760 torr
T1= 298K; T2= 273K; V1= 250 cm3; V2= ?
(V1P1/T1) = (V2P2/T2) therefore V2= (V1P1T2)/(T1P2)
V2 = (250 cm3)(742.2 torr)(273K)
(298K)(760.torr)
V2 = cm3 27

28. Kinetic Molecular Theory
Matter consists of particles (atoms or molecules) that are in continuous, random, rapid motion
The Volume occupied by the particles has a negligibly small effect on their behavior
Collisions between particles are elastic
Attractive forces between particles have a negligible effect on their behavior
Gases have no fixed volume or shape, but take the volume and shape of the container
The average kinetic energy of the particles is proportional to their Kelvin temperature 28

29. Maxwell-Boltzman Distribution
Molecules are in constant motion
Not all particles have the same energy
The average kinetic energy is related to the temperature
An increase in temperature spreads out the distribution and the mean speed is shifted upward 29

30.  Velocity of a Gas 3RT urms = M The distribution of speeds
of three different gases
at the same temperature
The distribution of speeds
for nitrogen gas molecules
at three different temperatures
urms =
3RT M  30

31. Diffusion
Gas diffusion is the gradual mixing of molecules of one gas with molecules of another by virtue of their kinetic properties.
NH4Cl
NH3
17.0 g/mol
HCl
36.5 g/mol 31

32. DIFFUSION AND EFFUSION
Diffusion is the gradual mixing of molecules of different gases.
Effusion is the movement of molecules through a small hole into an empty container. 32

33. Graham’s Law
Graham’s law governs effusion and diffusion of gas molecules.
KE=1/2 mv2
The rate of effusion is inversely proportional to its molar mass.
Thomas Graham, Professor in Glasgow and London. 33

34. 4 NH3(g) + 5 O2(g)  4 NO(g) + 6 H2O(g)
Sample Problem 9
1 mole of oxygen gas and 2 moles of ammonia are placed in a container and allowed to react at 850oC according to the equation:
4 NH3(g) + 5 O2(g)  4 NO(g) + 6 H2O(g)
Using Graham's Law, what is the ratio of the effusion rates of NH3(g) to O2(g)?
Scheffler 34

35. Sample Problem 10
What is the rate of effusion for H2 if cm3 of CO2 takes 4.55 sec to effuse out of a container?
Rate for CO2 = cm3/4.55 s = cm3/s 35

36. Sample Problem 11
What is the molar mass of gas X if it effuses times as rapidly as N2(g)?
Scheffler 36

37. Ideal Gases v Real Gases
Ideal gases are gases that obey the Kinetic Molecular Theory perfectly.
The gas laws apply to ideal gases, but in reality there is no perfectly ideal gas.
Under normal conditions of temperature and pressure many real gases approximate ideal gases.
Under more extreme conditions more polar gases show deviations from ideal behavior. 37

38. In an Ideal Gas ---
The particles (atoms or molecules) in continuous, random, rapid motion.
The particles collide with no loss of momentum
The volume occupied by the particles is essentially zero when compared to the volume of the container
The particles are neither attracted to each other nor repelled
The average kinetic energy of the particles is proportional to their Kelvin temperature
At normal temperatures and pressures gases closely approximate idea behavior 38

39. Real Gases
For ideal gases the product of pressure and volume is constant. Real gases deviate somewhat as shown by the graph pressure vs. the ratio of observed volume to ideal volume below.
These deviations occur because
Real gases do not actually have zero volume
Polar gas particles do attract if compressed 39

40. van der Waals Equation (P + n2a/V2)(V - nb) = nRT
The van der Waals equation shown below includes corrections added to the universal gas law to account for these deviations from ideal behavior
(P + n2a/V2)(V - nb) = nRT
where a => attractive forces between molecules
b => residual volume or molecules
The van der Waals constants for some elements are shown below
Substance
a (dm6atm mol-2)
b (dm3 mol-1) He
0.0341
CH4
2.25
0.0428
H2O
5.46
0.0305
CO2
3.59
0.0437 40

41. C6H12O6 (s) + 6O2 (g) 6CO2 (g) + 6H2O (l)
Sample Problem 12
What is the volume of CO2 produced at 370 C and 1.00 atm when 5.60 g of glucose are used up in the reaction:
C6H12O6 (s) + 6O2 (g) CO2 (g) + 6H2O (l)
g C6H12O mol C6H12O mol CO V CO2
1 mol C6H12O6
180 g C6H12O6 x
6 mol CO2
1 mol C6H12O6 x
5.60 g C6H12O6
= mol CO2
0.187 mol x x K
dm3•atm
mol•K
1.00 atm =
nRT P
V =
= 4.76 dm3 41