Chem 1A
Dr. White
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Workbook 5
5-1: Dalton’s Law, KMT, Effusion/Diffusion/Real Gases
1. What is the total pressure and the partial pressure of each gas (in atm) in a mixture of 3.2 g of
O2, 1.6 g of CH4 and 6.4 g of S2 in an 11.2 L container at 273°C?
2. What is the partial pressure of He (in atm) in a mixture of 1.0 g H 2 and 5.0 g He in a 5.0 L
cylinder at 20.0oC?
3. A 1.545 g sample of impure calcium carbide reacts with water to give 561 mL of C2H2 collected
by water displacement at 20°C and 745 mm Hg. The vapor pressure of water is 17.5 mm Hg.
What is the % CaC2?
CaC2(s) + 2 H2O(l) → Ca(OH)2(s) + C2H2(g)
4. A 3.00 g sample of a mixture contains copper and zinc. Zinc reacts with HCl but copper does
not. What is the % Zn if 927 mL of hydrogen gas is collected over water at 740 mm Hg and
20°C. The vapor pressure of water is 17.5 mm Hg.
5. In an apparatus, helium effuses at the rate of 15 mL/min. At what rate will xenon effuse in the
same apparatus?
6. In an effusion apparatus, H2 is found to effuse at the rate of 5.9 mL/s. Another gas in the same
apparatus effuses at the rate of 0.55 mL/s. What is the molar mass of the gas?
7. The van der Waals constants a and b are 4.194 L2 atm mole-2 and 0.05105 L mole-1 for Xe.
Calculate the observed pressure for 1.25 moles of the compound in a 1.000 L flask at 75°C.
8. The van der Waals constants a and b are 4.170 L2 atm mole-2 and 0.03707 L mole-1 for NH3.
Calculate the pressure (atm) of a 2.00 moles sample of NH3 in a 2.95 L flask at 47oC.
9. The Ne atom has 10 times the mass of H2. Which of the following statements is true?
I. At 25°C they both have the same kinetic energy.
II. Ten moles of H2 would have the same volume as 1 mole of Ne.
III. One mole of Ne exerts the same pressure as one mole of H2 at STP.
IV. A H2 molecule effuses 10 times faster than a Ne atom.
10. Real gases approach ideal gas behavior at: (chose one)
a) high pressure and low temperature
b) low pressure and low temperature
c) low pressure and high temperature
d) high pressure and high temperature
5-2: Bond Energies and Calorimetry
1. Use the bond energies in the table in your handout to calculate ΔE for the following reactions
(suggestion: draw Lewis structures first):
(a)
CH4 + Cl2 → CH3Cl + HCl
(b)
CH3CH2OH + 3 O2 → 2 CO2 + 3 H2O
Chem 1A
Dr. White
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2. How much heat is required to raise 335 g of water from 20.0 °C to 95.5 °C?
3. A 36.9 g sample of metal is heated to 100.0 °C, and then added to a calorimeter containing
141.5 g of water at 23.1 °C. The temperature of the water rises to a maximum of 25.2 °C
before cooling back down. Did the water absorb heat or did it release heat? How many joules
of heat was exchanged between the water and the metal?
4. When 1.00 g of solid NH4Cl is dissolved in 25.00 g water contained in a coffee cup calorimeter,
both reagents initially being at 25.0°C, the temperature falls to 22.4°C. Calculate the heat
(enthalpy) of solution of NH4Cl, (a) in J/g and (b) in kJ/mol.
5.
2.53 g of solid NaOH is dissolved in 100.0 g water in a coffee cup calorimeter, all the reagents
initially being at 20.0°C. Calculate the final temperature of the solution obtained, given the
following information:
NaOH(s) → NaOH(aq)
ΔHsoln = - 43.0 kJ
6. The reaction
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
was studied in a coffee cup calorimeter. 100. mL portions of 1.00 M aqueous NaOH and
H2SO4, each at 24.0°C, were mixed. The maximum temperature achieved was 30.6°C.
Neglect the heat capacity of the cup and the thermometer, and assume that the solution of
products has a density of exactly 1 g/mL. Calculate ΔH, the heat (enthalpy) of reaction, in
kJ/mol of Na2SO4 produced.
7. A mass of 1.250 g of benzoic acid (C7H6O2) was completely combusted in a bomb calorimeter.
If the heat capacity of the calorimeter was 10.134 kJ/K and the heat of combustion of benzoic
acid is -3226 kJ/mol, calculate (to three decimal places) the temperature increase that should
have occurred in the apparatus.
8. A common laboratory reaction is the neutralization of an acid with a base. When 50.0 mL of
0.500 M HCl at 25.0°C is added to 50.0 mL of 0.500 M NaOH at 25.0°C in a coffee cup
calorimeter, the temperature of the mixture rises to 28.2°C. What is the heat of reaction per
mole of acid? Assume that the densities of the reactant solutions are both 1.00 g/mL
9. A 5.00 g sample of HNO3 is dissolved in water in a calorimeter whose heat capacity is 5.16
kJ/oK. The temperature increases 0.511oK. Calculate the heat released (kJ) per mole of HNO3
dissolved.
10. When 2.62 g of lactic acid, C3H6O3, is burned in a calorimeter whose heat capacity is 21.7
kJ/oK, the temperature increases by 1.800oK. Calculate the heat released by the combustion of
lactic acid in kJ per mole.
5-3: Thermochemical Stoichiometry & Hess’ Law
1.
Calcium hydroxide, which reacts with carbon dioxide to form calcium carbonate, was used
by the ancient Romans as mortar in stone structures. The reaction for this process is
Ca(OH)2(s) + CO2(g) → CaCO3(s) + H2O(g)
ΔH = -69.1 kJ What is the enthalpy
change if 3.8 mol of calcium carbonate is formed?
2. The highly exothermic thermite reaction, in which aluminum reduces iron(III) oxide to
elemental iron, has been used by railroad repair crews to weld rails together.
2Al(s) + Fe2O3(s) → 2Fe(s) + Al2O3(s) ΔH = -8.5 x 102 kJ. What mass of iron is formed
when 725 kJ of heat are released?
Chem 1A
Dr. White
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3. Triglycerides are the main form in which fats are stored in the body. During periods of
starvation, a person’s fat stores are used for energy. Tristearin (C57H110O6) is a typical
animal fat that is oxidized according to the following thermochemical equation:
2C57H110O6 (s) + 163 O2 (g)  114 CO2 (g) + 110 H2O (l) ΔHrxn = -7.0 x 104 kJ
a. How much heat is released per gram of tristearin oxidized?
b. When 325 L of O2 at 37˚C and 755 torr is used, how many grams of tristearin can be
oxidized?
c. When 325 L of O2 at 37˚C and 755 torr is used, how many kJ of heat are released?
4. The reaction of barium metal with liquid water produces 660.2 kJ of heat for every mole of
barium that reacts.
(a) Write a complete balanced thermochemical equation for this reaction.
(b) Is this reaction endothermic of exothermic?
(c) Calculate the amount of heat associated with 3.65 g of water reacting at constant
pressure. Make sure your answer has the proper sign!
(d) How many grams of barium metal must react to produce 586 kJ of heat?
5. Pure liquid octane (C8H18, d= 0.702g/mL) is used as the fuel in a test of a new automobile
drive train.
a. How much energy is produced (in kJ) when a tank full (20.4gal) is combusted?
ΔHcomb= -5.45x103 kJ/mol)? (Start with a balanced equation for the combustion of one
mole of octane.)
b. The energy delivered at the wheels at 65mph is 5.5x 104 kJ/hr. Assuming that all the
energy is transferred to the wheels, what is the cruising range of the car (in km) on a
full tank?
6.
Calculate the enthalpy change for the reaction NO(g) + O(g) → NO2(g) from the following
data:
NO(g) + O3(g) → NO2(g) + O2(g) ΔH = -198.9 kJ
O3(g) → 1.5O2(g)
ΔH = -142.3 kJ
O2(g) → 2O(g)
ΔH = 495.0 kJ
(Note: O is NOT stable, as you know, and exists for a very short amount of time)
7. Use the thermochemical equations shown below to determine the enthalpy for the reaction:
H2O(l) → H2(g) + 1/2O2(g)
C(s) + O2(g) → CO2(g)
CO2(g) + 2H2O(l) →CH4(g) + 2O2(g)
C(s) + 2H2(g) → CH4(g)
ΔH=-590.2KJ
ΔH=1335.7KJ
ΔH=-112.2KJ
Chem 1A
Dr. White
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8. One problem with using hydrogen as a fuel is producing enough hydrogen efficiently. One
series of reactions being studied has as its net reaction the splitting of liquid water:
H2O (l)  H2 (g) + 1/2 O2 (g)
ΔH˚rxn = 285.8 kJ
This series of reactions involves each of the following steps in some form. Use Hess’s
Law to calculate the missing ΔH˚rxn.
H2 (g) + I2 (g)  2HI (g)
ΔH˚rxn = -10.64 kJ
H2O (l) + 1/2 SO2 (g) +1/2 I2 (g)  1/2 H2SO4 (aq) + HI (g)
2 H2O (l) + 2 SO2 (g) + O2 (g) 2 H2SO4 (aq)
ΔH˚rxn = ?
ΔH˚rxn =-649.82 kJ
(NOTE that this particular system is inefficient as one reaction requires a temperature of 825˚C. The
goal of research in this area is to find a system that requires low enough temperatures that sunlight
can be used as an energy source)
5-4: Heats of Formation & The Born Haber Cycle
1. a. Define, or explain fully what is meant by the standard enthalpy of formation of a substance, ΔH°f .
b. What is the standard state of the element oxygen?
c. Write down in full the formation reaction for liquid ethanol, C2H5OH(l). The equation should
be balanced and should indicate the physical state of each substance.
2. Write balanced chemical reactions for the formation of one mole of each of the following
compounds from its elements in their standard states:
a)
b)
c)
d)
e)
Liquid water
Aqueous strontium nitrate
Solid iron (III) bromide
Solid aluminum oxide
Solid magnesium phosphate
3. Lithium fluoride is formed from lithium and fluorine. Its lattice energy may be calculated from a
Born-Haber cycle using the following experimental data.
i. Li (g) → Li+ (g) + eΔH= +520 kJ
ii. Li (s) → Li (g)
ΔH= +161 kJ
iii. Li(s) + ½ F2 (g) → LiF (s)
ΔH = -617 kJ
iv. F2(g) → 2F(g)
ΔH= +159 kJ
v. F(g) + e → F (g)
ΔH = -328 kJ
a.
b.
c.
d.
e.
Which reaction above refers to the heat of formation for lithium fluoride?
Which reaction above refers to an electron affinity rection?
Which reaction above refers to a bond energy?
Write the reaction that refers to the lattice energy of LiF (s).
Calculate the lattice energy of lithium fluoride.
Chem 1A
Dr. White
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4. Nitric acid, which is among the top 15 chemicals produced in the United States, was first prepared
over 1200 years ago by heating naturally occurring sodium nitrate (called saltpeter) with sulfuric acid
and collecting the vapors produced. Calculate ΔH°rxn for this reaction. ΔH°f [NaNO3(s)] = -467.8
kJ/mol; ΔH°f [NaHSO4(s)] = -1125.5 kJ/mol; ΔH°f [H2SO4(l) = -814.0 kJ/mol; ΔH°f [HNO3(g)] = -135.1
kJ/mol
NaNO3(s) + H2SO4(l) → NaHSO4(s) + HNO3(g)
5. The space shuttle orbiter uses the oxidation of methyl hydrazine by dinitrogen tetroxide for
propulsion. The unbalanced reaction is as follows:
N2H3CH3 (l) +
N2O4 (l) 
H2 (g) + N2 (g) +
CO2 (g)
a)
b)
Balance this equation
Calculate ΔH˚rxn for this reaction using the following info:
Substance
N2O4 (g)
ΔH˚f (kJ/mol)
9.16
N2O4 (l)
-20.0
N2H3CH3 (l)
54.0
CO (g)
-110.5
CO2 (aq)
- 412.9
CO2 (g)
-393.5
6. a. Write a balanced equation for the combustion of benzene, C6H6(l) in oxygen.
b. The standard heat of combustion of benzene is -3271 kJ/mol. Calculate its standard heat of
formation, ΔH°f , given the data:
ΔH°f [CO2(g)] = -394 kJ; ΔH°f [H2O(l)] = -286 kJ
7. Lightweight camp stoves often make use of a mixture of C5 and C6 liquid hydrocarbons (a fuel
called “white gas.”)
a. Write the reaction for the combustion of C5H12 (l) and determine the standard heat of
formation of C5H12 (l) if the standard heat of combustion is -3540 kJ per mole of C5H12.
The ΔH°f [CO2(g)] = -393.5 kJ/mol and The ΔH°f [H2O(g)] = -241.8 kJ/mol.
b. How much heat is produced by the complete combustion of 3.00 L of C5H12 (l) if the density
of C5H12 (l) = 0.625 g/mL?
5-5: Heat, Work, and the First Law of Thermodynamics
1.
What are the two main components of internal energy of a substance? What are the
symbols for internal energy and its two components?
2. A system which undergoes an adiabatic change is one in which no heat is transferred. For
an adiabatic change that does work on its surroundings, indicate if q, w and ΔE for such a
process should be positive, negative, or equal to zero. Explain.
Chem 1A
Dr. White
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3. For each of the following, define the system and the surroundings, and indicate the
direction of heat transfer.
a) Natural gas is burned in a gas furnace in your home.
b) Water drops, sitting on your skin after a dip in the pool, evaporate
4. A system delivers 200. J of pressure-volume work against the surroundings while releasing
300. J of heat energy. What is the change in the internal energy of the system?
5. The work done when a gas is compressed in a cylinder is 199 J. A heat transfer of 270 J
occurs from the surrounding to the gas. Calculate E of the gas in J.
6. One mole of a gas at 25oC expands in volume from 1.0 L to 4.0 L at constant temperature.
What work (J) is done if the gas expands against an external pressure of 3.0 atm?
7. One mole of a gas at 25oC expands in volume from 2.0 L to 6.0 L at constant temperature.
What work is done if the gas expands against an external pressure of 0.75 atm?
8. 0.506 moles of a gas with a molar heat capacity of 5.75 J/mol°C are placed in a 1.50 L
container at 25.0°C. The temperature increases from 25.0°C to 31.0°C, and the container
expands to 2.75 L against a pressure of 1.02 atm. Calculate q, w, ΔE, and ΔH for the
gas.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
5-1: Dalton’s Law, KMT, Effusion/Diffusion/Real Gases
Ptotal = 1.2 atm, partial pressure of each = 0.40 atm
6.0 atm
92.6%
79.8%
2.6 mL/min
2.3 x 102 g/mol
31.6 atm
16.3 atm
I. TRUE II. FALSE
III. TRUE
IV. FALSE
C
1.
2.
3.
4.
2
5-2: Bond Energies & Calorimetry
b. -1267 kJ
a. -1.10 x 10 kJ
1.06 x 105 J
absorbs, 1.24 x 103 J
a. 2.8 x 102 J/g
b. 15 kJ/mol
5. 26.3°C
6. a. -5.5 × 103 J
b. -110 kJ/mol Na2SO4
7. Temperature increase is 3.258 K (or °C)
8. -54 kJ/mol
9. -33.2 kJ/mol
10. -1.34 x 103 kJ/mol
1. -2.6 x 102 kJ
2. 95 g
5-3: Thermochemical Stoichiometry & Hess’ Law
Chem 1A
Dr. White
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3. a. -39 kJ/g
b. 138 g
c. -5.4 x 103 g
4. a. Ba(s) + 2H2O (l) → Ba(OH)2 (aq) + H2 (g) ΔHrxn = -660.2 kJ
b. exothermic
c. -66.9 kJ
d. 122 g
5. a. -2.59 x 106 kJ
b. 4.9 x 103 km
6. -304.1 kJ
7. 428.9 kJ
8. -24.88 kJ
5-4: Heats of Formation & The Born Haber Cycle
1.
a. ΔH°f is the enthalpy change accompanying the formation of one mole of a substance
from its elements, all substances being in their standard states.
b. Pure O2 gas at a pressure of 1 atm and a specified temperature.
c. 2C(graphite) + 3H2(g) + ½O2(g) → C2H5OH(l)
2. a. ½ O2 (g) + H2 (g) → H2O (l))
b. Sr(s) + N2 (g) +3 O2 (g) → Sr(NO3)2 (aq)
c. Fe (s) +3/2 Br2 (l) → FeBr3 (s)
d. 2Al (s) + 3/2 O2 (g) → Al2O3 (s)
e. 3Mg (s) + ½ P4 (s) + 4O2 (g) → Mg3(PO4)2 (s)
3. a. rxn iii b. rxn v
c. rxn iv
d. Li+ (g) + F- (g) → LiF (s) e. -1050 kJ
4. 21.2 kJ
5. a. 2 N2H3CH3 (l) + N2O4 (l) -> 6H2 (g) + 3 N2 (g) +2CO2 (g)
b. -875.6 kJ
6. a. C6H6 (l) + 15/2 O2 (g) → 6CO2 (g) + 3H2O (l)
b. 49 kJ
7. a. 122 kJ/mol
b. 9.20 x 104 kJ
5-5: Heat, Work, and the First Law of Thermodynamics
1. Internal energy (E) is composed Heat (q) and work (w).
2. q = 0 (no heat is exchanged), w = - (energy lost from the system as work), ΔE = - (since ΔE
= q + w and q is 0)
3. a. System – burning of gas
Surroundings – everything else (furnace, home and the entire rest of the universe)
Heat transferred from the system to surroundings (-q)
b. System – water drops
Surroundings – everything else
Heat transferred from the surroundings to system (+q)
(this is why your skin – part of the surroundings – feels cool when sweat or water evaporates)
4. -500. J
5. 469 J
6. -9.1 x 102 J
7. -3.0 x 102 J
8. q = ΔH = 17 J
w = -129 J
ΔE = -112 J
Chem 1A
Dr. White
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