Station 1 – LT 4.1
Lesson
4.1 – Pressure and the Gas Laws
How is pressure created?
Questions 1.
2. Convert the following pressures:
a. 1.45 atm = ________ torr
b. 780. mm Hg = _______ atm
c. 987 torr = ______ barr
d. 9 barr = ______ atm
3. Why does atmospheric pressure decrease as you go higher up in
the earth’s atmosphere?
Station 2 – LT 4.2
Lesson
4.1 – Pressure and the Gas Laws
A 40.0 L tank of ammonia has a pressure of 8.00 atm. Calculate
Questions 1.
the volume of the ammonia if its pressure is changed to 12.0 atm
while its temperature remains constant.
2. At 27.00 °C a gas has a volume of 6.00 L. What will the volume
be at 150.0 °C?
3. A gas has a pressure of 0.370 atm at 50.0 °C. What is the pressure
at standard temperature?
4. 2.00 liters of hydrogen, originally at 25.0 °C and 0.98 atm, are
heated until a volume of 20.0 liters and a pressure of 3.50
atmospheres is reached. What is the new temperature?
5. You have 42 L of N2 gas, how many moles of N2 do you have?
Station 3 – LT 4.3
Lesson
4.1 – Pressure and the Gas Laws
A sample of gas (1.90 mol) is in a flask at 21.0 °C and 697.0
Questions 1.mm
Hg. The flask is now opened and more gas is added to the
flask. The new pressure is 795.0 mm Hg and the temperature is
now 26.0 °C. How many moles of gas are now in the flask’ if the
volume of the flask does not change?
2. What is the density of carbon dioxide at 298 K and 1.0 atm?
3. Find the number of grams of CO2 that exert a pressure of 785
torr at a volume of 32.5 L and a temperature of 32 C.
4. Calculate the density of H2S at STP.
5. At what temperature does 16.3 g of nitrogen gas have a pressure
of 1.25 atm in a 25.0 L tank?
6. What mass of CO2 is needed to fill an 80.0 L tank to a pressure
of 150.0 atm at 27.0°C?
Station 4 – LT 4.4
Lesson
4.2 – Partial Pressures
1. If I place 3 moles of N2 and 4 moles of O2 in a 35 L container at a temperature
Questions of 25 oC, what will the pressure
of the resulting mixture of gases be?
2. Two flasks are connected with a stopcock. The first flask has a volume of 5
liters and contains nitrogen gas at a pressure of 0.75 atm. The second flask has a
volume of 8 L and contains oxygen gas at a pressure of 1.25 atm. When the
stopcock between the flasks is opened and the gases are free to mix, what will the
pressure be in the resulting mixture?
3. What’s the partial pressure of carbon dioxide in a container that holds 5 moles
of carbon dioxide, 3 moles of nitrogen, and 1 mole of hydrogen and has a total
pressure of 1.05 atm?
4. The volume of hydrogen collected over water is 453 mL at 18 oC and 780. mm
Hg. What is the volume of dry hydrogen at STP?
5. Calculate the mass of 400. mL of carbon dioxide collected over water at 30. oC
and 749 mm Hg.
Station 5 – LT 4.5
Lesson
4.3 – Kinetic Molecular Theory
1. You have two gases at 25 oC. One of the gases is N2 and the other is H2. How
Questions do these gases compare in terms of:
a. Average kinetic energy?
b. Root mean square speed?
c. Rate of effusion?
d. Density?
2. You double the temperature of a gas, by how much does the average kinetic
energy of this gas change? What about if you cut the temperature in half?
3. You have a sample of gas at 25 oC and another sample of this gas at 100 oC.
Draw a molecular speed distribution plot for these two gases.
4. You have a sample of He in one flask and you have a sample of O2 in another
flask. Both flasks have a temperature of 25 oC. Draw a molecular speed
distribution plot for these two gases.
Station 6 – LT 4.6
Lesson
4.4 – Diffusion, Effusion, and Real
Gases
1. What does Graham’s Law of effusion/diffusion explain about the rate of two
Questions gases effusion/diffusion rate with respect to their molecular weight?
2. How does kinetic molecular theory explain the process of diffusion?
3. What is the rate of effusion for a gas that has a molar mass twice that of a gas
that effuses at a rate of 3.62 mol/min?
4. Calculate the rate of effusion of NO2 compared to SO2 at the same temperature
and pressure.
Station 7 – LT 4.7
Lesson
4.4 – Diffusion, Effusion, and Real
Gases
Questions
1. Under what conditions does a gas deviate from ideal behavior? Why does
this deviation occur?
2. How does the Van der Waals equation of state correct for these deviations
from ideality?
3. Calculate the pressure exerted by 0.3000 mol of helium in a 0.2000 L
container at -25 °C for both ideal and real behavior. aHe = 0.0341 atm·L2/mol2,
bHe = 0.0237 L·mol
4. 12.0 moles of methane gas are in a 5.00 L tank at 24.0 oC. Calculate the
difference in pressure between methane and an ideal gas under these conditions.
The van der Waals constants for methane are a = 2.300 L2atm mol-2 and b =
0.0430 L mol-1
Station 8 – LT 4.8 and 4.9
Lesson
4.5 – Intermolecular Forces
1. Compare and contrast the three types of intermolecular forces that are found in
Questions molecules (dispersion, dipole – dipole, and hydrogen bonding)
2. Explain why ethyl alcohol (C2H5OH) has a higher boiling point (78.40 C) than
methyl alcohol (CH3OH; 64.70 C).
3. Rank the following by from lowest to highest anticipated boiling point: C 2H4,
CH4, Ne, H3COCH3.
4. Motor oil largely consists of molecules that consist of long chains of carbon
atoms with hydrogen atoms attached to them. Using your knowledge of
intermolecular forces, why wouldn’t it be better to use a compound like glycerol.
The formula of glycerol is CHOH(CH2OH)2.
Station 9 – LT 4.10
Lesson
4.6 – Phase Diagrams
1. Draw the phase diagram for a substance that has a normal melting point 0 oC, a
Questions normal boiling point of 100 oC, a triple point of 4.58 torr and 0.00098 oC and a
critical point at 218 atm and 374 oC.
2. Draw the phase diagram for a substance that has a triple point at 6 atm and -15
o
C, a critical point at 300 atm and 200 oC, a melting point at 30 atm and 0 oC, and
a boiling point at 30 atm and 50 oC. The solid is more dense than the liquid.
3.What is the critical temperature of compound X? ___________
4. If you were to have a bottle containing compound X in your closet, what phase
would it most likely be in? Note: Atmospheric pressure is 1 atm.
5. At what temperature and pressure will all three phases coexist?
6. If I have a bottle of compound X at a pressure of 45 atm and temperature of
1000 C, what will happen if I raise the temperature to 400 0 C? Describe all the
phases that it would exist in.
7. Why can’t compound X be boiled at a temperature of 200 0 C?
8. If I wanted to, could I drink compound X? Note: Atmospheric pressure is 1 atm.
Station 10 – 4.11
Lesson
4.7 – Solution Concentrations
1. What is the molarity of a solution prepared by dissolving 37.94 g of potassium
Questions hydroxide in some water and then diluting the solution to a volume of 500.00 mL
2. Determine the molality of a solution containing 81.3 g of ethylene glycol,
HOCH2CH2OH, dissolved in 166 g of water.
3. Determine the molality of a solution of 560 g of acetone, CH3COCH3, in 620 g
of water
4. You have a solution that is 16.2% sodium thiosulfate, Na 2S2O3. What mass of
sodium thiosulfate is in 80.0 g of solution?
5. A gas mixture contains 45.6g of carbon monoxide and 899 g of carbon dioxide.
What is the mole fraction of carbon monoxide?
6. A mixture of CO2 and Kr weighs 35.0 g and exerts a pressure of 0.708 atm in its
container. Since Kr is expensive, you wish to recover it from the mixture. After
the CO2 is completely removed by absorption with NaOH(s) the pressure in the
container is 0.250 atm. How many grams of CO2 and how many grams of Kr were
initially present?
Station 11 – 4.12, 4.13, 4.14
Lesson
4.6, 4.7, and 4.8 – Colligative
Properties
Questions
1. Which of the following aqueous solutions will have the highest boiling point?
a. 0.1 m Sodium Chloride
b. 0.1 m Potassium Fluoride
c. 0.1 m Magnesium Chloride
d. 0.1 m Sugar (Sucrose) Solution
e. 0.1 m Potassium Nitrate
2. What is the boiling point of a solution that contains 30.0 grams of sucrose
(C12H22O11) dissolved in 1000. grams of water? The Kb for water is 0.51 oC
kg/mol.
3. What is the freezing point of a solution that contains 20.0 grams of phenol
(C6H5OH) dissolved in 250 grams of water? The Kf for water is 1.86 oC kg/mol
4. 40.0 grams of an organic compound dissolved in 1.00 kg of water has a boiling
point of 100.52 oC. What is the molecular weight of the solute?
5. At 20 oC, a nearly ideal solution is prepared by mixing 0.400 mol of methanol
and 0.600 mol of ethanol. The vapor pressures of the pure liquids at 20 oC are: Po
methanol = 88.7 mmHg;
Po ethanol = 44.5 mmHg
(a) Use Raoult's Law to calculate the equilibrium vapor pressure of the solution
(b) Use Dalton's Law of Partial Pressure to calculate the composition of the vapor
above the solution.
Station 12 – 4.12, 4.13, 4.14
Lesson
4.6, 4.7, and 4.8 – Colligative
Properties
Questions
1. Answer the following questions, which refer to the 100 mL samples of aqueous
solutions at 25˚C in the stoppered flasks shown above.
(a) Which solution has the lowest electrical conductivity? Explain.
(b) Which solution has the lowest freezing point? Explain.
(c) Above which solution is the pressure of the water vapor greatest? Explain.
2. If 105 g of I2 is dissolved in 325 g of CCl4 at 65 oC. Given that the pure vapor
pressure of CCl4 at this temperature is 531 mm Hg, what is the vapor pressure of
the solution?
3. A 0.7549 g sample of CxHyOz burns to make 1.9061 g CO2 and 0.3370 g H2O.
Also, in a different experiment, a 0.5246 gram sample was dissolved in 10 g of
lauric acid and the freezing point was lowered by 1.68 oC.
a. What is the empirical formula of the substance?
b. If Kf of lauric acid is 3.90 oC/m, calculate the molar mass of the
substance.
c. What is the molecular formula of the substance?
4. A substance is found to be 65.6% C and 9.4% H and 25% O. In a different
experiment, 1.57 g of the compound were dissolved in 16.08 g of camphor and it
was observed that the freezing point of camphor was 15.2 oC lower than pure
camphor.
a. What is the empirical formula of the compound?
b. If Kf of camphor is 40 oC/m, calculate the molar mass of the substance.
c. What is the molecular formula of the substance?