Honors Chemistry
Name: ___________________________________ Date: _______________ Mods: ______________
10.6 – Gas Mixture & Partial Pressures
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Gas Mixtures:
o
Gases mix _________________________
o
Gases molecules hardly interact with one another so each molecule behaves as if it were
______________ in the container.
Dalton’s Law of Partial Pressures: The _________________ pressure of a mixture of gases is
equal to the ____________ of the pressures that each gas would exert if it were present alone.
Equation:
[P1, P2, etc. are known as the ____________________ pressures of each gas in the mixture]
o
In the same container, each gas within the mixture has the same ________________________
and ____________________, so pressure will only be affected by the ___________________
of each gas.
P1 = n1 RT
+
V
P2 = n2 RT
etc…
V
To find Pt (total pressure) you could solve for
each partial pressure and sum them or…
…since R, T, and V are the same for every gas in a mixture, Pt can be found by the following:
Equation:
[nt is total moles of all gases in the mixture]
o
Mole Fractions: Again, since only the # of moles will effect each gas’s partial pressure, we can
make a ratio of # of moles of each gas in the mixture to the total # of moles of gas (this is
known as the mole fraction of a gas and is represented by a capital X)
X1 = n1
nt

mole fraction of gas #1 =
moles of gas #1
total moles of all gases in mixture
… the mole fraction gives us another way to solve for the partial pressure of a gas if the total
pressure of the mixture and # of moles are known
Equation:
Example Problem: 6.00 g O2 and 9.00 g CH4 are placed inside a 15.0 L vessel at 0ºC. What is
the partial pressure of each gas in the mixture? What is the total pressure inside the container?
Note: this problem can be solved 2 different ways…
Way #1 – using PV = nRT, solve each gas for the partial pressures (PO2 & PCH4)
and then sum the partial pressure to get the total pressure (Pt)
Way #2 – find the total number of moles in the gas mixture (nt) and use that in
PV= nRT to solve for pressure (if you do this, P will be the total pressure (Pt)
then use the mole fractions of each gas (XO2 and XCH4) and the total pressure
with the equation P1 = X1Pt to find the partial pressures (PO2 & PCH4)
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Gases Collected over Water:
o
In many labs, a gas is collected over water (see picture below)
o
Since a body of water always has some water vapor above it, the gas being collected forms a
mixture with H2O (g) …
o
Using Dalton’s Law of Partial Pressures…
Equation:
[ Note: the values of PH2O are listed in Appendix B for various temperatures ]
Example Problem: A sample of KClO3 is partially decomposed, producing oxygen gas which is then
collected over water. The volume of gas collected is 0.250 L at 26ºC and 765 torr total pressure.
2 KClO3 (s)  2 KCl (s) + 3 O2 (g)
a) How many moles of O2 are collected?
b) How many grams of KClO3 were decomposed?
Dalton’s Law of Partial Pressures WS
1)
If I place 3 moles of N2 and 4 moles of O2 in a 35 L container at a temperature of 25ºC, what will
the pressure of the resulting mixture of gases be? What would be the partial pressure of each
gas?
2)
What’s the partial pressure of carbon dioxide in a container that holds 5.3 moles of carbon
dioxide, 3.1 moles of nitrogen, and 1.2 mole of hydrogen and has a total pressure of 1.05 atm?
3)
In a mixture of neon and argon gases, if there are 0.31 mols of argon and 0.67 mols of neon in a
24.9 L container at 33ºC, what is the total pressure of the mixture? What would be the partial
pressure of each gas?
Gases Collected over Water WS
1) A 450 mL sample of hydrogen gas is collect over water at 12oC. The total pressure of the
hydrogen and water vapor mixture is 78.5 kPa. What is the partial pressure of the dry hydrogen
gas? What mass of the hydrogen gas was collected?
2)
0.888 L of oxygen gas is collected over water at a temperature of 25 oC. The total pressure of the
gases is 55.8 kPa. What is the partial pressure of the dry gas?
3)
H2 (g) was collected over water at 25°C and the mixture of gases had a total pressure of 755.0
mm Hg. If 90.0 L of hydrogen gas were collected, how many grams of CaH2 must have been
used?
CaH2 (s) + 2 H2O (l)  Ca(OH)2 (aq) + 2 H2 (g)