Problem #1: This reaction was carried out:
CaCO3(s) + 2HCl(aq) ---> CaCl2(s) + CO2(g) + H2O(ℓ)
What would be the volume of CO2 (at STP) produced from the complete reaction of 10.0 grams of CaCO2?
1) Determine moles of CaCO3 reacted:
10.0 g / 100.086 g/mol = 0.099914 mol
2) Note that there is a 1:1 molar ratio between CaCO3 used and CO2 produced.
3) Since the problem is using STP, we can use molar volume:
(22.414 L/mol) (0.099914 mol) = 2.24 L (to three sig figs)
4) If the pressure and temperature were not at STP, we would use the ideal gas law to calculate the volume produced.
PV = nRT
(1.00 atm) (V) = (0.099914 mol) (0.08206 L atm / mol K) (273 K)
V = 2.24 L (to three sig figs)
Note that I set up the problem using STP values. If non-STP values were to be used, they are typically given in the body of the question.
Problem #2: 0.84 g of ammonium dichromate is decomposed. Here is the chemical reaction:
(NH4)2Cr2O7(s) ---> N2(g) + 4H2O(g) + Cr2O3(s)
The gases from this reaction are trapped in a 13.6 L flask at 26.0 °C. (a) What is the total pressure of the gas in the flask? (b) What are the partial pressures of N2 and H2O?
1) Determine the moles of ammonium dichromate used:
0.84 g / 252.0622 g/mol = 0.00333251 mol
2) From the balanced equation, each mole of ammonium dichromate decomposed produces 5 moles of gas.
0.00333251 mol x (5 mol gas / 1 mol dichromate) = 0.01666255 mol gas
3) Use the ideal gas law to calculate the pressure:
PV = nRT
(P) (13.6 L) = (0.01666255 mol) (0.08206 L atm/mol K) (299.0 K)
P = 0.030 atm (to two sig figs, based on the 0.84)
4) Of the gases produced, 1/5 is N2 and 4/5 is H2O.
partial pressure N2 ---> (0.030 atm) (0.2) = 0.0060 atm
partial pressure H2O ---> (0.030 atm) (0.8) = 0.024 atm
Problem #3: A 4.90-g sample of solid CoCl2 ⋅ 4H2O was heated such that the water turned to steam and was driven off. Assuming ideal behavior, what volume would that steam occupy at 1.00 atm and 100.0 °C?
1) Determine moles of CoCl2 ⋅ 4H2O in 4.90 g:
4.90 g / 201.8982 g/mol = 0.0242696 mol
2) Determine moles of water produced:
1 mol CoCl2 ⋅ 4H2O will produce 4 mol H2O
0.02427 mol CoCl2 ⋅ 4H2O will produce:
(4) (0.0242696 mol) = 0.0970784 mol H2O
3) Calculate volume at 1.00 atm and 100.0 °C:
PV = nRT
(1.00 atm) (V) = (0.0970784 mol) (0.082057 L atm / mol K) (373 K)
V = 2.97 L (to three sig figs)
Problem #4: If 39.5 mL of H2 are produced at 21.0 °C when the atmospheric pressure is 762.8 mmHg, and the height of the liquid column in the eudiometer is 11.2 cm, what mass of aluminum is used?
1) The pressure of the wet gas in the eudiometer plus the 11.2 cm of water equals the measured atmospheric pressure. We need to obtain the pressure of the dry gas.
11.2 cmH2O = 112 mmH2O
(112 mmH2O) (1.00 g/mL) = (mmHg) (13.6 g/mL)
x = 8.2353 mmHg
At 21.0 °C, the vapor pressure of water is 18.7 mmHg. Found here.
762.8 − 8.2353 − 18.7 = 735.8647 mmHg
2) The next step is to use PV = nRT to determine moles of H2.
(735.8647 mmHg / 760 mmHg/atm) (0.0395 L) = (n) (0.08206 L atm / mol K) (294 K)
n = 0.0015853 mol of H2 (notice that I have carried several guard digits)
3) Let's write a chemical equation. Let us assume the Al reacted with HCl.
2Al + 6HCl ---> 2AlCl3 + 3H2
The key ratio is the Al to H2 molar ratio of 2 to 3
2 x ––– = –––––––––––– 3 0.0015853 mol
y = 0.00105687 mol of Al
4) The last step:
(0.00105687 mol) (26.98 g/mol) = 0.0285 g (to three sig figs)
Problem #5a: A 0.616 gram sample of a metal, M, reacts completely with sulfuric acid according to the reaction:
M(s) + H2SO4(aq) ---> MSO4(aq) + H2(g)
A volume of 239 mL of hydrogen is collected over water; the water level in the collecting vessel is the same as the outside level. Atmospheric pressure is 1.0079 bar and the temperature is 25.0 °C. The vapor pressure of water at 25.0 °C is 0.03167 bar. Calculate the molar mass of the metal.
1) Use Dalton's Law to get the pressure of the dry hydrogen:
Ptotal = PH2 + PH2O
1.0079 − 0.03167 = 0.97623 bar
2) Determine moles of H2 produced:
PV = nRT
(0.97623 bar) (0.239 L) = (n) (0.0831447 L bar / mol K) (298 K)
n = 0.00941671 mol
Note that R has the units L bar / mol K. This is a lesser-used unit for R, but the value associated with that unit can be easily looked up.
3) From the balanced chemical equation, M and H2 are in a 1:1 molar ratio. Therefore:
0.00941671 mole of M was consumed.
4) Calculate the molar mass of M:
0.616 g / 0.00941671 mol = 65.4 g/mol
If the problem had asked to identify the metal, the answer would have been zinc.
Problem #5b: 19.4 grams of metal sulfide has reacted with H2SO4 solution and 4.46 L of H2S gas was collected at STP. The oxidation number of the metal is +2. Determine what element the metal is.
MS + H2SO4 ---> H2S + M2SO4 [balanced as written]
(4.46 L) / (22.414 L/mol) (1 mol MS / 1 mol H2S) = 0.1989828 mol MS
(19.4 g MS) / (0.1989828 mol MS) = 97.495864 g /mol <--- the molar mass of MS
(97.495864 g/mol) − (32.0655 g/mol) = 65.43 g/mol <--- 32.0655 g/mol is the atomic weight of sulfur
65.43 g/mol is (within experimental error) the atomic weight of zinc.
Problem #6: Calculate the volume of nitrogen monoxide gas produced when 8.00 g of ammonia is reacted with 11.0 g of oxygen at 25.0 °C. The density of nitrogen monoxide at 25.0 °C is 1.23 g/L.
4NH3(g) + 5O2(g) ---> 4NO(g) + 6H2O(ℓ)
1) Determine the limiting reagent:
ammonia: 0.46973 mol / 4 mol = 0.1174
oxygen: 0.34376 mol / 5 mol = 0.0687
Oxygen is the limiting reagent.
2) Determine moles of O2 present:
11.0 g / 31.9988 g/mol = 0.343763 mol
3) Determine moles of NO produced:
The molar ratio beween O2 and NO is 5 to 4.
5 0.343763 mol ––– = –––––––––––– 4 x
x = 0.2750104 mol of H2
4) Determine mass of NO produced:
0.2750104 mol times 30.006 g/mol = 8.251962 g
5) Determine volume of NO produced:
8.251962 g divided by 1.23 g/L = 6.61 L
Problem #7: C4H10 combusts. What mass of oxygen is needed to make 3.00 L of water at 0.990 atm and 295 K.
1) The balanced equation:
2C4H10 + 13O2 ---> 8CO2 + 10H2O
2) Moles of water present in the 3.00 L:
PV = nRT
(0.990 atm) (3.00 L) = (n) (0.08206 L atm / mol K) (295 K)
n = 0.122688 mol
3) The molar ratio between O2 and H2O is 13 to 10.
13 x ––– = –––––––––––– 10 0.122688 mol
x = 0.1594944 mol of H2
4) Moles to grams:
(0.1594944 mol) (32.0 g/mol) = 5.10 g (to three sig figs)
Problem #8: A student collected 17.32 mL of H2 over water at 30.0 °C. The water level inside the collection apparatus was 6.60 cm higher than the water level outside. The barometric pressure was 731.0 torr. How many grams of zinc had to react with the HCl solution to produce the H2 that was collected?
1) The pressure inside the collection tube is made up of three things:
1) the H2 gas
2) water vapor
3) the 6.60 cm column of water
The sum of the three above pressures is equal to 731 torr.
2) The 6.60 cm column of water must be converted to mmHg pressure.
We use the densities of water and mercury to do this:
(6.60 cm) (1.00 g/cm3) = (x) (13.534 g/cm3)
x = 0.488 cm of mercury
This equals 4.88 mmHg
731.0 − 4.88 = 726.12 mmHg <--- the pressure of the gas minus the 6.60 cm column of water
3) Now, we need to remove the pressure of the water vapor.
We need the vapor pressure of water at 30.0 °C, a value we look up.
726.12 − 31.8 = 694.32 mmHg <--- that's the pressure of just the H2
4) Now, we use PV = nRT to determine the moles of H2:
(694.32 / 760) (0.01732 L) = (n) (0.08206 L atm / mol K) (303 K)
n = 0.0006364 mol
(694.32 / 760) <--- that converts mmHg to atm
5) Now, some stoichiometry to get the mass of zinc:
Zn + 2HCl ---> ZnCl2 + H2
The molar ratio of Zn to H2 is 1:1, so we now know that 0.0006364 mol of Zn was used.
0.0006364 mol times 65.38 g/mol = 0.0416 g
Problem #9: Hydrogen gas is produced by the reaction of sodium metal with an excess of hydrochloric acid solution. The hydrogen gas was collected by water displacement at 22.0 °C and 127.0 mL was collected with a total pressure of 748.0 torr. The vapor pressure of water at 22.0 °C is 19.8 torr.
(a) What mass of sodium metal was consumed in the reaction?
(b) What is the volume of dry H2 gas at STP?
Solution to (b):
1) The chemical reaction is this:
2Na + 2HCl ---> 2NaCl + H2
In order to determine the amount of sodium that reacted, I must know how many moles of H2 was produced. I do that by solving (b) first.
2) Use Dalton's Law of Partial Pressures to determine the pressure of the dry hydrogen gas:
PH2 + 19.8 torr = 748.0 torr
PH2 = 728.2 torr
3) Use PV = nRT to determine moles of gas:
(728.2 torr / 760.0 torr/atm) (0.1270 L) = (n) (0.08206 L atm / mol K) (295 K)
n = 0.00502675 mol
I used the Ideal Gas Law because I knew I needed moles of hydrogen to get my answer for (a).
4) Continue with the solution for (b) by using the Combined Gas Law to get the volume of gas at STP:
P1 = 728.2 torr P2 = 760.0 torr V1 = 127.0 mL V2 = x T1 = 295 K T2 = 273 K
(728.2 torr) (127.0 mL) (760.0 torr) (x) –––––––––––––––––– = ––––––––––– 295 K 273 K
x = 112.6 mL
Solution to (a):
1) Use 2:1 molar ratio between Na and H2 to determine moles of Na that react:
2 is to 1 as x is to 0.00502675 mol
x = 0.0100535 mol
2) Convert moles to grams:
(0.0100535 mol) (22.99 g/mol) = 0.2311 g
Problem #10: Automobile air bags inflate during a crash or sudden stop by the rapid generation of nitrogen gas from sodium azide according to the reaction:
2NaN3 ---> 2Na + 3N2
How many grams of sodium azide are needed to provide sufficient nitrogen gas to fill a 30.0 cm x 30.0 cm x 25.0 cm bag to a pressure of 1.20 atm at 26.0 °C?
1) Determine the volume of the bag:
30.0 cm x 30.0 cm x 25.0 cm = 22500 cm3 = 22500 mL = 22.5 L
2) Determine moles of gas filling the bag:
PV = nRT
(1.20 atm) (22.5 L) = (n) (0.08206 L atm mol¯1 K¯1) (299 K)
n = 1.100426 mol
3) Determine moles, then mass of sodium azide required:
2 x ––– = –––––––––––– 3 1.100426 mol
x = 0.733617 mol of NaN3
(0.733617 mol) (65.011 g/mol) = 47.7 g
Bonus Problem #1: Given the following reaction: N2H4 + O2 ---> N2 + 2H2O
Determine the kL of water produced at STP when 12.36 kg of hydrazine reacts with sufficient oxygen.
12.36 kg N2H4 1000 g N2H4 1 mol N2H4 2 mol H2O 22.414 L H2O 1 kL H2O –––––––––––– x –––––––––– x –––––––––––– x ––––––––– x ––––––––––– x ––––––––– = 17.29 kL H2O 1 1 kg N2H4 32.0452 g N2H4 1 mol N2H4 1 mol H2O 1000 L H2O
Bonus Problem #2: One of the steps in the production of iron utilizes the following chemical reaction:
3CO(g) + Fe2O3(s) ---> 2Fe(s) + 3CO2(g)
(a) What mass of Fe2O3 would react with 200.0 liters of CO at STP?
(b) What volume of carbon dioxide (CO2) at STP is produced from 100.0 grams of Fe2O3?
(c) What mass of iron (Fe) is produced when 300. mL of CO2 is produced at STP?
Solution to (a):
1) Determine moles of CO:
200.0 L / 22.414 L/mol = 8.923 mol
Note use of molar volume in this problem (and the one above) rather than PV = nRT. Since we are at STP, we can use molar volume. Any other conditions require the use of PV = nRT.
2) Determine moles, then mass, of Fe2O3:
CO and Fe2O3 react in a 3:1 molar ratio.
8.923 mol is to x as 3 is to 1
x = 2.97433 mol
(2.97433 mol) (159.687 g/mol) = 475 g
Solution to (b):
1) Moles of Fe2O3 consumed:
100.0 g / 159.687 g/mol = 0.626225 mol
2) Moles of CO2 produced:
Fe2O3 and CO2 are in a 1:3 molar ratio.
(0.626225 mol) (3) = 1.878675 mol
3) Determine volume of CO2 produced:
PV = nRT
(1.00 atm) (V) = (1.878675 mol) (0.08206 L atm / mol K) (273 K)
V = 42.1 L
Solution to (c):
1) Moles of CO2. Use molar volume:
0.300 L / 22.414 L/mol = 0.0133845 mol
2) Moles of Fe:
Iron and carbon dioxide are in a 2:3 molar ratio
2 is to 3 as x is to 0.0133845
x = 0.008923 mol of Fe
3) Mass of Fe:
(0.008923 mol) (55.845 g/mol) = 0.498 g