Chapter 12: Gases and Their
Properties
1
Properties of Gases

Gases form homogeneous mixtures
 Gases are compressible
 All gases have low densities



air
water 1.00
iron 7.9
0.0013 g/mL
g/mL
g/mL

Gases expand to fill their containers uniformly
 A gas exerts a pressure
2
Kinetic Molecular Theory





Gases consist of molecular particles moving in
straight lines at any given instant.
Molecules collide with each other and the
container walls without any net loss of energy.
Gas molecules behave independently -attractive/repulsive forces between them are
negligible.
Gas molecules are widely spaced, the actual
volume of molecules is negligible compared to the
space they occupy.
The average kinetic energy of the gas particles is
proportional to the temperature.
3
Kinetic Graph
4
Pressure, Volume, and Temperature
Relationships
A.
B.
C.
D.
E.
Pressure, volume, and temperature units
Boyle's law
Charles' law
Gay-Lussac's law
Avogadro's law
5
A. Pressure, Volume, and
Temperature Relationships
Pressure =
force
unit area
-measured with barometer
or manometer
SI: 1 pascal (Pa) = 1 N/m2
common: 1 atm = 760 mm Hg
= 760 torr
= 14.7 lb/in2 (psi)
= 101.325 kPa
= 1.013 bar
Volume: mL, L
Temperature: K
6
Pressure

Defined as force per unit area
Force mass acceleration
P ressure 

Area
distance
2
(kg)(m/sec )
kg


 kilopascals
2
2
m
msec
7
Atmospheric pressure
8
Barometer
9
Manometers
10
Pressure Units

101.325 kPa
 = 760 mmHg
 =760 torr
 =1 atm
 =30 in Hg
 =14.7 psi
11
B. Boyle’s law
Robert Boyle, 1662: for a sample of gas at constant T, V  1/P
V=
constant
P
or
PV = constant
 P1V1 = P2V2 (at constant T)
12
Boyles Law
13
C. Charles’ law
Jacques Charles, 1787: for a sample of gas at constant P, V  T (K)
V = constant x T

V1
V2
=
T1
T2
or
V
= constant
T
(at constant P)
14
Charles Law
15
D. Gay-Lussac’s law
Joseph Gay-Lussac, ~1800: for a sample of gas at constant V, P  T (K)
P = constant x T

P1
P2
=
T1
T2
or
P
= constant
T
(at constant V)
16
E. Avogadro’s law
Amadeus Avogadro, 1811: at constant T and P, V  n
V = constant x n

V1
V2
=
n1
n2
or
V
= constant
n
(at constant T & P)
17
The Ideal gas laws
Since PV = constant, P/T = constant, V/T = constant, and V/n = constant

PV
= constant, R (universal gas constant)
nT
or PV = nRT
18
STP and molar volume
STP: 0ºC (273 K) and 1.00 atm (760 torr)
molar volume = average volume occupied by one mole of gas at STP
= 22.4 L/mol
PV
(1.00 atm)(22.4 L)
R = nT =
(1.00 mol)(273 K)
= 0.0821 L·atm/mol·K
= 6.24 x 104 mL·torr/mol·K
19
Gas densities and molar mass
m
mw =
n
PV
since n =
RT
mRT
 mw =
PV
m
and since d =
V
dRT
 mw =
P
e.g., If one finds that a 0.108-g sample of gas occupies a volume of
238 mL at 25ºC and 525 torr, what is the molecular weight of the gas?
20
Dalton’s law of partial pressures
1. Dalton’s law
partial pressure, p = pressure exerted by each gas in a mixture of gases
Ptotal 
p
i
i
e.g., If 6.00 g of O2 and 9.00 g of CH4 are placed in a 15.0-L container
at 0º, what is the partial pressure of each gas and the total pressure in
the container.
21
Dalton’s law of partial pressures
2. mole fraction
PT = p1 + p2 + p3…
p1 n1RT / V n1


 X 1 , molefractionof gas 1
PT nT RT / V nT
 p1 = X1PT
e.g., Air is 78 mol % N2 and 22 mol % O2. What is the partial pressure
of each gas if the atmospheric pressure is 713 torr?
22
Dalton’s law of partial pressures
3. collecting gases over water
vapor pressure = pressure exerted by evaporation of a liquid (or
sublimation of a solid)
vacuum
liquid
or solid
pvap
equilibrium
at some T
pvap increases with T
e.g., H2O
T
pvap
0ºC
4.6 torr
25ºC 23.8 torr
100ºC 760 torr (boils!)
23
Dalton’s law of partial pressures
3. collecting gases over water
Mixture of gas and water
vapor.
PT = pgas + pvap(H2O)
24
Mole Fraction
molesof component
Mole fraction(X) 
totalmoles
and
P1  (X1 )(Ptotal )
25