1. Molar Volume; Gas Density
AP Chemistry
Unit 4
Gases

2. STP Standard Temperature & Pressure
Common measuring conditions for gases
0˚Celsius or 273 Kelvin
1 atm = kPa =
760 torr = 760 mmHg

3. Molar Volume: All gases @ STP occupy approximately 22.42 L per mole.

4. What mass of helium is required to fill a 1.5 L balloon @ STP.
1.5 L x 1 mole x g =
22.42 L mole
0.27 g
**If STP use Ideal Gas Law
(1 atm) (1.5 L) = n ( L·atm)(273 K)
mol·K
n = moles x 4.00 g/1 mole
= 0.27 grams

5. Densities of Gases n P V = RT
If we divide both sides of the ideal-gas equation by V and by RT, we get n V P RT =

6. Densities of Gases n   = m P RT m V = We know that
moles  molecular mass = mass
n   = m
So multiplying both sides by the molecular mass ( ) gives P RT m V =

7. Densities of Gases P RT m V = d = Note: One only needs to
Mass  volume = density
So, P RT m V =
d =
Note: One only needs to
know the molecular mass, the pressure, and the temperature to calculate the density of a gas.

8. Molecular Mass P d = RT dRT P  = We can manipulate the density
equation to enable us to find
the molecular mass of a gas: P RT
d =
Becomes
dRT P
 =

9. The density of a gas measured at 1. 50 atm, 27˚C and found to be 1
The density of a gas measured at 1.50 atm, 27˚C and found to be 1.95 g/L. Calculate the molar mass of the gas.
dRT P
 =
M = (1.95 g/L) ( )(300 K)
1.50 atm
M = 32 grams

10. Dalton’s Law of Partial Pressures
The total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone.
In other words,
Ptotal = P1 + P2 + P3 + …

11. Mole Fraction χ = Px Ptotal
The ratio of # moles of a given component in a mixture to the total number of moles in the mixture.
Combined with Dalton’s Law is used to find the pressure exerted by the individual gases in a mixture.
Mole fraction can also be calculated with moles or liters.

12. A mixture of 1. 00 g of H2 and 1. 00 g of He is placed in a 1
A mixture of 1.00 g of H2 and 1.00 g of He is placed in a 1.00 L 35˚C. Calculate total pressure and each partial pressure.
1.00 g x 1 mole/2.02 g = moles
1.00 g x 1 mole/4.00 g = mole
total = moles
P(1.00 L) = (0.745 moles)( )(308 K)
P = 18.8 atm
PH2: 0.495/0.745 x 18.8 atm = 12.5 atm
PHe: 0.259/0.745 x 18.8 atm = 6.31 atm

13. Partial Pressures
When one collects a gas over water, there is water vapor mixed in with the gas.
To find only the pressure of the desired (dry) gas, one must subtract the vapor pressure of water (which depends upon temperature only) from the total pressure.