Chemistry Chapter 10 notes
Physical Characteristics of Gases
Kinetic molecular theory of matter
• All matter is composed of tiny particles
which are in constant motion
• This explains observed properties of
matter
Kinetic molecular theory of matter
(KM)
• Ideal gas- an imaginary gas which
perfectly fits all assumptions of the kinetic
molecular theory of matter
• IE: Ideal gas behaves exactly as a gas
should, no deviations
• Kinetic molecular theory of gases based
on 5 assumptions
5 Assumptions
1. Gases consist of large numbers of
particles very far apart from one another
relative to their size
– Most of space occupied by gases is empty
space
– This explains compressibility of gases and
their low density (compared to liquids and
solids)
2. Collision between gas particles/particles
and gas particles/container are elastic
– Elastic collision- no net loss of kinetic
energy.
– KE is transferred, but total KE of 2 particles
remains the same as long as the
temperature is constant
3. Gas particles are in continuous, rapid,
random motion and therefore have kinetic
energy
• Their KE is high enough to overcome any
attractive forces between particles (except
near the temperature of condensation)
4. There are no forces of attraction or
repulsion between gas particles
• When gas particles collide, they
immediately bounce apart
5. Average kinetic energy of gas particles
depends on the temperature of the gas
– For any object KE= ½ m v2
– Where m= mass and v = velocity
– All gases at the same temperature have same
KE, so lighter particles (H) have higher
average speed than heavier particles (O)
KM theory and the nature of gases
• Expansion
– Gases fill any container and take it’s shape
– Gases have no definite shape or volume
• Fluidity
– Gas particles glide past one another
– Behave much like liquids
– Gases and liquids are both considered fluids
KM theory and the nature of gases
• Low density
– Gases typically have about 1/1000 the density
of the same substance in a liquid or solid
state
• Compressibility
– Due to their low density gases can be
compressed dramatically
KM theory and the nature of gases
• Diffusion
– Gases randomly mix with other particles to
even distribution
– Rates of diffusion depend on the speed of
particles, diameter of particles and attractive
forces between particles
– Lighter gases diffuse more rapidly than
heavier ones
KM theory and the nature of gases
• Effusion
– Movement of gas particles through a tiny
opening
– Rates of effusion are directly proportional to
the velocity of the particles
Real gases
• Do not behave completely according to kinetic
molecular theory
• 1873 Van der Waals noted that forces between
particles of gases caused deviation from ideal
gas behavior
• Deviation is most significant at high pressure
and low temperature
• KM theory holds truest in gases with little
attraction between particles (ex. Noble gases)
Pressure
• When describing a gas you must specify
characteristics: Volume, temperature,
number of molecules and pressure
• You’ve got the first 3!
• Pressure is force per unit area on a
surface or Pressure = force/area
Atmospheric pressure
• Pressure exerted by gases of the
atmosphere
• At sea level approximately 10.1 N/cm2
• Barometers are used to measure
atmospheric pressure
• Oldest barometer- mercury column
measurement expressed in mm of Hg
– Normal atmospheric pressure at sea level and
0°C = 760 mm Hg = 1 atmosphere
Pressure!
• SI units for pressure are derived
• 1 Pascale (Pa) = 1 Newton / meter2
• Pressure often expressed in kilopascals
(kPa)
• 1 atmosphere = 1.01325 x 105 Pa
(or 101.325 kPa)
• See table 10-1 on p. 311
STP
• Standard Temperature and Pressure are
needed to compare gas volumes
• STP = 1 atmosphere and 0°C
Gas Laws
Boyles law
• Relates pressure and volume of a gas at
constant temperature
• Pressure and volume are inversely
proportional
• PV = k or V=k1/P
• K is a constant for a given sample of gas
Boyles and changing pressure
• Because k is a constant for a given
sample of gas and we know that the
product of pressure and volume will
always equal k
• P1V1 = k and P2V2 = k we can set P1V1
equal to P2V2
• P1V1= P2V2 and solve for any one of the 4
values
Charles Law
• Relates temperature and volume of gases
at constant pressure
• 1787 Charles found that volume of a gas
changes 1/273 of original volume for each
1°C change in temperature (with a starting
point of 0°C and at constant pressure)
Charles and absolute zero
• Kelvin 0 = -273.15°C
• K= °C + 273.15
• This is useful because it is directly
proportional to gas volume
• Charles law: Volume of a fixed sample of
gas at constant pressure varies directly
with Kelvin temperature
Charles…
• V/T= k or V = kT
• K is a constant based on quantity of gas
and pressure
• Same thing can be done with Charles for
changing volume or temperature as was
done with boyles for changing pressures
• V1/T1 = V2/T2
Gay-Lussacs Law
• Relates pressure and temperature of a
gas at constant volume
• P/T = k or P= kT
• K is a constant depending on quantity and
volume of gas
• P1/T1 = P2/T2 useful when faced with
changing pressures and temperatures
Combined gas law
• Merges three laws just mentioned
• PV/T = k
• k is a constant related to the amount of
gas
• P1V1/T1 = P2V2/T2
• if any one quantity is unchanging one of
the other gas laws can be derived
Daltons combined pressures
• The total pressure of a mixture of gases is
the sum of the individual pressure of each
gas alone
• PT= P1 + P2 + P3…
• This can be used no matter how many
gases are in combination
Law of Combined Pressures
• Is useful when dealing with gases
collected over water
• Gases collected this way are mixed with
water vapor, this exerts water vapor
pressure
• To measure pressure of gas and water vapor in
collection bottle, raise bottle until water level in
and out are same.
• At that point pressure inside bottle =
atmospheric pressure
• Patm = P gas + P H2O
• Obtain atmospheric pressure from barometer in
lab and subtract water vapor pressure at given
temp (from table A8 in book)