H CH 11 Homework
Stoichiometry:
1. How many liters of ammonia gas (NH3) can be produced by the reaction of 4000 L of
Nitrogen(g) w excess Hydrogen(g)?
2. If 1800 L of acetylene gas (C2H2) is burned, how many liters of CO2(g) are formed?
3. Electrolysis of liquid water produced 500 L of Hydrogen gas. What volume of
Oxygen gas was also produced?
4. If 200 L of carbon monoxide gas react with 400 L of Oxygen gas, how many liters of
carbon dioxide gas will be formed?
5. How many liters of Oxygen gas at .78 atm and 34 C are required to react with 55 L
of Hydrogen gas at the same conditions?
6. When 70 g of H2O2 (s) are decomposed, how many liters of oxygen gas are formed?
2H2O2  2H2O + O2
7. If 40 g of solid Aluminum react with 60 g of hydrochloric acid (aq), how many liters
of Hydrogen gas are produced?
8. How many grams of solid potassium chlorate are required in the preparation of 90 L
of Oxygen gas at 1.20 atm and 45 C?
9. How many grams of solid Sodium Peroxide must react with water to produce 35 L of
Oxygen gas at 5 C and 660 mmHg?
2 Na2O2 + 2 H2O  4 NaOH + O2
10. Calculate the number of liters of Oxygen gas needed to react with 75 g of solid
Aluminum.
11. If 35 g of solid Sulfur react with 80 L of Oxygen gas, how many liters of Sulfur
Dioxide gas would be produced?
12. Oxygen gas was originally produced by the decomposition of solid Mercuric Oxide
(HgO). What volume of O2(g) at 15 C and .98 atm can be produced from 2.36 g of
HgO (s)?
13. If 32 g of solid Sodium are added to 25 L of Chlorine gas at 45 psi and 300 K, how
many grams of solid salt would be produced?
1
Collection Over Water:
1. 11 g of solid potassium chlorate is decomposed and the resulting oxygen gas is
collected over water. What volume of oxygen would be collected at .93 atm and 25
°C?
2. 20 g of solid potassium bromide is mixed with 10 L of chlorine gas at 34 psi and 20
°C. What volume of liquid bromine can be collected over water at these conditions?
3. 16.7 ml of liquid water is electrolyzed. The resulting gases are sent to 2 different
tanks that are initially filled with water. The hydrogen tank is at 30 °C and the
oxygen tank is at 70 °C. Which tank must have the larger volume? Atmospheric
pressure on this day is 710 mmHg.
Density
1. What is the density of Oxygen gas at STP? Nitrogen gas?
2. What is the density of Fluorine gas at STP and at 30 C and 725 Torr?
3. Find the density of Freon 12 (g) (CF2Cl2) at 315 K and .945 atm.
Grahams Law
1. If Hydrogen gas will diffuse at a rate of .3 m/s at certain conditions, how fast will
Oxygen gas diffuse at those conditions?
2. How much faster will Carbon Monoxide gas diffuse as compared to Carbon Dioxide
gas?
3. What is the relative rate of diffusion of Fluorine gas to Chlorine gas?
2
Review:
1.
47 g of solid magnesium oxide is heated and decomposes. How many liters of
oxygen gas are produced at 582.3 mmHg at 300 K? What if the oxygen was
collected over water?
2.
What is the density of carbon monoxide gas at 27 psi and 200 °C?
3.
If you combine 30 g of hydrogen gas with 15 liters of chlorine gas at 744 Torr
and 73°F, how many of mL of liquid HCl (density 1.18 g/mL) would be
produced?
4.
Chlorine gas diffuses at .15 m/s. How fast will Ne gas diffuse at the same
conditions?
3
Temperature Pressure Temperature Pressure
(degrees C) (mmHg)
(degrees C)
(mmHg)
0
4.6
50
92.5
1
4.9
51
97.2
2
5.3
52
102.1
3
5.7
53
107.2
4
6.1
54
112.5
5
6.5
55
118
6
7
56
123.8
7
7.5
57
129.8
8
8
58
136.1
9
8.6
59
142.6
10
9.2
60
149.4
11
9.8
61
156.4
12
10.5
62
163.8
13
11.2
63
171.4
14
12
64
179.3
15
12.8
65
187.5
16
13.6
66
196.1
17
14.5
67
205
18
15.5
68
214.2
19
16.5
69
223.7
20
17.5
70
233.7
21
18.7
71
243.9
22
19.8
72
254.6
23
21.1
73
265.7
24
22.4
74
277.2
4
25
23.8
75
289.1
26
25.2
76
301.4
27
26.7
77
314.1
28
28.3
78
327.3
29
30
79
341
30
31.8
80
355.1
31
33.7
81
369.7
32
35.7
82
384.9
33
37.7
83
400.6
34
39.9
84
416.8
35
42.2
85
433.6
36
44.6
86
450.9
37
47.1
87
468.7
38
49.7
88
487.1
39
52.4
89
506.1
40
55.3
90
525.8
41
58.3
91
546
42
61.5
92
567
43
64.8
93
588.6
44
68.3
94
610.9
45
71.9
95
633.9
46
75.7
96
657.6
47
79.6
97
682.1
48
83.7
98
707.3
49
88
99
733.2
100
760
5
/20
LAB: GAS COLLECTION
1. Obtain the following:
pneumatic trough
#6 rubber stopper with glass bend and rubber tube and thistle tube
250 ml Erlenmeyer flask
5 - 125 ml Erlenmeyer flasks
1 piece of rubber tubing
an unknown amount of NaHCO3
A glass plate and a 50 ml graduated cylinder
2. Place the 2 pieces of the rubber tubing on the pneumatic trough. The piece of tubing on the
glass bend should be connected to the bottom opening in the trough. The other piece of tubing
will be connected to the top hole of the trough and the other end should be placed in the sink.
3. Place the rubber tubing in the 250 ml flask so that the thistle tube is just above the bottom of
the flask.
4. Fill the trough to just below the overflow outlet (top hole).
5. Since we are going to be determining the volume of gas produced by collecting it in the 125
ml flask, we need to calibrate the flask using the graduated cylinder so that we can determine the
exact volume of the flask.
6. Fill your 5 - 125 ml flasks with water and, by placing a glass plate over the top of the flask,
invert the flasks and place them in the trough. Place one of the flasks over the opening in the
bottom of the trough.
7. Add your unknown amount of NaHCO3 to the 250 ml flask (NOT through the thistle tube).
Write the letter of your unknown in the data table.
8. Measure out about 20 ml of 3M acetic acid. Place the lid on the 250 ml flask and add no more
than 5 ml of acetic acid through the thistle tube at a time. Make sure to swirl the acid around to
make sure that it is all reacted. Once all the acid has reacted, add more.
9. Collect the CO2 in the trough. When the 125 ml flask is almost full, you will need to remove
it and place the next flask over the opening in the trough.
10. Once all of the acetic acid has reacted (ie: no more bubbling when you add acid) measure the
volume of the CO2 produced. Also determine the temperature of the water in the trough.
Calculations:
1. Calculate the amount of NaHCO3 that you started with. It was either 1.00g, 1.50g, 2.00g.
Correct the volume of your gas to standard conditions
REMEMBER:
Reactions do not produce 100% yield.
6
/20
Experiment: Determination of the Gas Law
Constant
Purpose
In this experiment you will calculate the gas law constant, R, by collecting a
known quantity of hydrogen gas and measuring the temperature, pressure and
volume of the gas collected.
Introduction
From the ideal gas law, PV = nRT, you can see that it is possible to determine a
value for R if you can isolate a sample of gas for which P, V, T and n are all
known. In this experiment you will accomplish this by collecting hydrogen gas
formed in the reaction of magnesium metal with hydrochloric acid.
When you collect the hydrogen gas you will also measure the temperature,
pressure and volume of the gas collected. From this data a value for R can be
calculated.
Since the hydrogen will be collected over water in a gas collection tube, it will be
saturated with water vapor. According to Dalton’s Law, the total pressure of the
gas mixture is the sum of the partial pressure of H2 plus the partial pressure of
the water vapor. The partial pressure of the water vapor can be looked up in
tables in the CRC, and subtracted from the total pressure to find the pressure of
the H2. The volume of the gas will be measured directly from the gas collection
tube.
The necessary calculations are as follows:

Determine the number of moles of H2 produced from the experimental
mass of Mg used in the reaction.

Record the temperature of the gas.

Measure the volume of gas collected in the tube.

Determine the total pressure of the gas in the tube using a barometer.

Determine the partial pressure of the H2 gas by subtracting the partial
pressure of the water vapor from the total pressure.

Calculate a value for R.
7
Procedure
1. Obtain a gas collection tube and a plastic beaker. You will also need a
clamp for mounting the tube on a ringstand.
2. Get 10 cm copper wire, a 1.00 cm strip of magnesium, and a “00” one hole
stopper.
3. Mass the magnesium.
4. Roll the Mg into a loose ball, and wrap the Mg with the Cu wire to form a
cage around the Mg. Leave at least 2.5 cm length of Cu wire coming off of
the cage.
5. Obtain about 10 mL of 6 M HCl. Add the acid to the gas collection tube.
Then carefully add water to completely fill the tube, being careful to not
disturb the HCl as the water is added. Pouring the water slowly down the
side of the tube will help.
6. Run the length of wire from the copper cage through the one hole stopper,
bend the wire over the top of the stopper, and insert the assemblage into
the gas collection tube so that the cage of Cu holding the Mg is inside the
tube. Make sure that the tube is completely filled with water and that no
air bubbles remain.
7. With one finger covering the hole in the stopper, invert the tube into the
beaker of water. Mount the tube to a ring stand, being careful to not let the
end of the tube come out of the water in the process. As the HCl flows
down the tube it will react with the Mg, generating H2 gas which will rise
to the top of the tube and displace water out of the bottom.
8. When the reaction is complete (you will see a dramatic change in the
amount of hydrogen being given off) record the volume of gas in the gas
collection tube.
9. Record the temperature of the gas collected, assuming its temperature is
the same as that of the water in the beaker and gas collection tube. Also
record the barometric pressure.
10. Repeat the experiment a second time.
8
Analysis and Calculations
1. Assuming that the magnesium reacted completely, calculate the
moles of hydrogen formed from the mass of magnesium you
actually used.
2. Calculate the pressure of the dry H2 from the barometer reading
and the vapor pressure of water. (Look in the CRC index under
vapor pressure, aqueous vapor, below 100°C.)
3. Using the experimental values for pressure, temperature, volume,
and moles of the gas, calculate a value for R. Repeat the
calculations for your second trial. If your two trials agree report an
average.
4. Compute a percent error.
9
/15
LAB: GRAHAM'S LAW
1. Obtain a glass tube and test tube clamp.
2. Clamp the tube to the ring stand so that the tube is level.
3. Obtain a cotton ball and using forceps, dip the cotton ball into the hydrochloric acid.
Do the same with another cotton ball and the unknown.
4. Place the cotton balls in each end of the tube and gently place a rubber stopper in each
end of the tube to close it.
5. Watch for a precipitate to form inside the tube.
6. Measure the distance from each end of the tube to the precipitate and determine which
of the following gasses is your unknown:
NH3
NH4OH
(NH4)2SO4
NH4Br
7. Turn in all calculations and data.
10
NH4C2H3O2