South Pasadena • Honors Chemistry
Name
6 • Thermochemistry
Period
6.2
PROBLEMS
1. Write a chemical equation for the following
changes (include “heat”). Indicate whether each is
endothermic or exothermic.
(a) The melting of ice.
Heat + H2O (s)  H2O (ℓ)
endothermic
(b) The condensation of steam.
H2O (g)  H2O (ℓ) + heat
exothermic
endothermic
(d) The freezing of liquid water.
H2O (ℓ)  H2O (s) + heat
exothermic
2. Answer each question. Explain briefly.
(a) Blocks A and B are made of different
substances, but have the same mass. It takes
3000 J to melt a block A, and 4500 J to melt
block B. Which has a greater enthalpy of
fusion, ∆Hfus (in J/g)? Explain without using
math.
Block B has greater ∆Hfus since it is more
difficult (requires more energy) to melt.
(b) The enthalpies of vaporization for ethanol and
water are 4.36 cal/mol and 18.0 cal/mol,
respectively. If 2000 calories of heat are added
to a sample of each substance, which substance
will produce more moles of vapor? Explain
briefly.
CHANGES IN STATE
4. How many grams of methanol (CH3OH) vapor can
be condensed if it releases 1674 J of energy?
∆Hvap = 35.4 kJ/mol
q = ‒1674 J
∆H = ‒35.4 kJ/mol
n=
(c) The vaporization of liquid water.
Heat + H2O (ℓ)  H2O (g)
–
Date
n=?
q
‒1674 J  1 kJ 
=
= 0.0473 mol
∆H ‒35.4 kJ/mol 1000 J
m = 0.0473 mol 
32.0 g
 1 mol  = 1.51 g
5. How much energy is released to vaporize the 30.0
gram sample of acetone, C3H6O?
∆Hvap = 31.3 kJ/mol
1 mol 
n = 30.0 g 
58.08 g = 0.517 mol
∆H = +31.3 kJ/mol
q=?
q = n ∆H = (0.517 mol)(+31.3 kJ/mol) = 16.2 kJ
Questions 6-13: For water,
Cice = 2.10 J/g·°C, Cwater = 4.18 J/g·°C,
Csteam = 2.08 J/g·°C, ∆Hfus = 6.01 kJ/mol,
∆Hvap = 40.68 kJ/mol.
6. Find the value for q when 15.0 g water freezes.
∆Hfus = ‒ 6.01 kJ/mol
1
mol

n = 15.0 g 
18.02 g = 0.833 mol
q = n ∆H = (0.833 mol)(‒6.01 kJ/mol) = ‒5.00 kJ
q=?
7. How many grams of water are converted to steam
when 15,000 J of heat is absorbed?
Ethanol will produce more vapor since it has
a lower ∆Hvap and is easier to vaporize.
q = +15,000 J
n=?
3. What is the enthalpy of fusion of aluminum, in
∆H = +40.68 kJ/mol
kJ/mol, if 5970 J of energy is required to melt a
q
15000 J  1 kJ 
n=
=
= 0.369 mol
15.0 gram sample?
∆H 40.68 kJ/mol 1000 J
q = +5970 J
18.02 g
m = 0.369 mol 
= 6.64 g
1 mol 

1 mol 

n = 15.0 g
26.98 g = 0.556 mol ∆H = ?
q
5970 J
=
n 0.556 mol
= 10700 J/mol or 10.7 kJ/mol
∆H =
8. How much energy is released when 50.0 g of water
vapor condenses into liquid?
12. Find the heat required when 1.50 g ice at 0°C is
completely vaporized at 100°C.
1 mol 
n = 50.0 g 
18.02 g = 2.78 mol
∆H = ‒ 40.68 kJ/mol
q1 = n ∆H = (1.50 g) 
q = n ∆H = (2.78 mol)(‒40.68 kJ/mol) = ‒113 kJ
q2 = m C ∆T
q=?
9. How much energy is absorbed when 1.80 moles of
ice melts?
q=?
n = 1.80 mol
∆H = + 6.01 kJ/mol
q = n ∆H = (1.80 mol)(+6.01 kJ/mol)
= 10.8 kJ
10. Find the energy required to heat up 35.0 g of water
at 80°C to steam at 115°C.
q1 = m C ∆T = (35.0 g)(4.18 J/g·°C)
1 kJ 
(100°C ‒ 80°C)
1000 J = 2.93 kJ
q2 = n ∆H = 35.0 g 
1 mol 
18.02 g (6.01 kJ/mol)
= 0.500 kJ
1 kJ 
= (1.50 g)(4.18 J/g·°C)(100°C ‒ 0°C) 
1000 J
= 0.627 kJ
q3 = n ∆H = (1.50 g) 
1 mol 
18.02 g (40.68 kJ/mol)
= 3.39 kJ
q = q1 + q2 + q3
= (0.500 kJ) + (0.627 kJ) + (3.39 kJ)
= 4.52 kJ
13. Find the heat released for 2.50 g steam at 100°C to
condense to water at 20°C.
q1 = n ∆H = (2.50 g) (‒40.68 kJ/mol) = ‒5.64 kJ
1 mol 
18.02 g (40.68 kJ/mol)
q2 = m C ∆T
= 79.0 kJ
q3 = m C ∆T = (35.0 g)(2.08 J/g·°C)
1 kJ 
(115°C ‒ 100°C)
1000 J = 1.09 kJ
= ‒0.836 kJ
q = q1 + q2 + q3 = 2.93 kJ + 79.0 kJ + 1.09 kJ
= 83.0 kJ
11. Find the energy released when 80.0 g water at 0°C
is cooled to ice at ‒20°C.
q1 = n ∆H = (80.0 g) 
1 mol 
18.02 g (‒6.01 kJ/mol)
= ‒26.7 kJ
q2 = m C ∆T
1 kJ 
= (80.0 g)(2.01 J/g·°C)(‒20°C ‒ 0°C) 
1000 J
= ‒3.22 kJ
q = q1 + q2 = ‒26.7 kJ + (‒3.22 kJ) = ‒29.9 kJ
= (2.50 g)(4.18 J/g·°C)(20°C ‒ 100°C) 
1 kJ 
1000 J
q = q1 + q2 = (‒5.64 kJ) + (‒0.836 kJ) = ‒6.48 kJ