1. Kinetic-Molecular Theory and an Introduction to Pressure & Gas Stoich
Gases
Kinetic-Molecular Theory and an Introduction to Pressure & Gas Stoich

2. The Kinetic-Molecular Theory of Gases
based on five assumptions. . .
1. Gases consist of large numbers of tiny particles that are far apart relative to their size
this means that most of the space occupied by a gas is empty
this is why gases can be compressed

3. 2. Collisions between gas particles and between particles and container walls are elastic.
elastic collisions mean that there is no net loss of kinetic energy
in other words, kinetic energy is never lost, it is just transferred between the colliding particles

4. 3. Gas particles are in constant, rapid, random motion.
the particles are always moving around
4. There are no forces of attraction or repulsion between gas particles
no intermolecular force exists between gas molecules
I would assume that London Dispersion forces are negligible

5. 5. The average kinetic energy of gas particles depends on the temperature of the gas.
temperature is just a measure of the average kinetic energy

6. some additional characteristics of gases. . .
gases will expand to fill any container that you put them in
gases have the ability to flow, just like a liquid
gases have low density
by applying pressure, gases can be compressed to a smaller volume
gases spread out via diffusion due to the positive change in entropy associated with the process
some gases do not behave as ideal gases because the gas molecules are polar and attracted to one another

7. Pressure pressure is defined as the force per unit area on a surface
for example. . .
the air in a balloon exerts a force on the inner walls of the balloon
the SI unit for force is the Newton (N)

8. the atmosphere exerts a pressure on the surface of the earth
this is due to the weight of the gas molecules between the earth and the edge of the atmosphere

9. The air has mass and therefore exerts a force on the surface of the earth due to gravity
As a result, we feel pressure
AIR
would someone on a mountain feel as much pressure as someone at sea level?
Earth
NO, because they don’t have as much air above them
Atmosphere

10. Can anyone guess the average pressure at sea level?
1 atmosphere (atm)
the device used to measure atmospheric pressure is called a barometer
barometers generally consist of a tube filled with mercury
pressure is measured based on the distance the pressure causes the mercury to move

11. Units of Pressure 1. millimeters of mercury (mm Hg)
the average atmospheric pressure at sea level is 760 mm Hg
2. atmospheres of pressure (atm)
1 atmosphere is equal to 760 mm Hg
3. the SI unit for pressure is the pascal (Pa)
one pascal is the same as 1 N/m2
1 atm= x 105 Pa= kPa
4. 1 torr = 1 mm Hg
Inches of mercury
1 atm = in
Pounds per square inch
1 atm = 14.7 lb/in2

12. some equations we will use for gases will require certain units, so you will usually have to convert between various units of pressure
for example. . .
How many millimeters of mercury are in 3.6 atm?
1 atm = 760 mm Hg
3.6 atm X
760 mm Hg
2736 mm Hg =
1 atm

13. How many kPa are in 905 mm Hg?
905 mm Hg X
kPa
121 kPa =
760 mm Hg

14. PV = n R T Review of Gas Stoichiometry
Remember a mole of gas at STP occupies 22.4 L
So, 2 moles would occupy 44.8 L at STP
11.2 L of a gas would have moles at STP
When the conditions are not standard, you must use the ideal gas law
P = pressure in atm
V= volume in L
n = number of mols
R= Gas Constant : atm L/ mol K
T = Temperature in Kelvin
PV = n R T

15. Carbon dioxide decomposes to carbon monoxide and oxygen gas
Carbon dioxide decomposes to carbon monoxide and oxygen gas. How many liters of oxygen gas will be made from 80.0 g of carbon dioxide at 755 mm Hg and 25 oC?
1. Write a balanced equation
2CO2 (g) 2CO (g) + O2 (g)
2. Calculate the moles of CO2
#mol= mass(g)/molar mass
#mol= 80.0 g/44.01 g/mol =
1.818 mol CO2

16. 3. Calculate the moles of O2 made
1 mol O2
1.818 mol CO2 X =
2 mol CO2
0.909 mol O2
4. Rearrange the equation
PV=nRT
V=nRT/P
5. Solve the equation
V=(0.909 mol)(0.0821)(298 K)/( atm)
V= 22.4 L

17. in the previous example, you had to use the ideal gas law because the conditions were not standard
At STP, you could have multiplied the moles of oxygen gas by 22.4 L/mol to find the total volume
using the ideal gas law is always acceptable, just more work than necessary at STP