AP chem.
Thermochemistry, Solids & Liquids
Name ____________________
(lesson 1 AIM: How do we measure heat transfer using calorimetry?)
6.1
video
1 Calculate the total heat required to bring 18 g H2O
from solid ice at -30oC to steam at 140oC.
Hvap =
Hfus =
cice =
cwater =
csteam =
2259 J/g
334 J/g
CALORIMETRY (q1 + q2 + q3
SPECIFIC HEAT OF A METAL
2 A 150.0 g sample of a metal at 75.0oC is added to 150.0 g H2O at 15.0oC. The temperature of the
water rises to 18.3oC. Calculate the specific heat of the metal, assuming all the heat lost by the
metal is gained by the water.
HEAT OF A REACTION
3 When 1 mol Ba(NO3)2(s) reacts with 1 mol Na2SO4(s) at 25.0oC in 2000 g of water, the
temperature of the mixture rises to 28.1oC. Calculate H for this reaction. (assume that
csolution = 4.18 J/goC)
HEAT OF SOLUTION FORMATION
4 Calcium chloride has a heat of solution listed as -81.3 kJ/mol when dissolved in water. Using this, suppose we
dissolve 10.0 g of calcium chloride in 50.0 mL of water at 25oC. What will be the new temperature of the water?
(assume that csolution = 4.18 J/goC)
o
C
C
oC
o
AP chem.
Thermochemistry, Solids & Liquids
Name ____________________
PROBLEM SET 6.1
5 40 g ice at 0oC is mixed with 250 g water at 90oC.
8 A 100 g sample of a liquid with a specific heat of 2
Determine the final temperature of the mixture,
J/g·oC absorbs 4000 J of heat energy. If the sample
assuming everything is liquid water in the end.
started at 30oC, what is the final temperature? No
phase change took place.
6 How many joules are required to heat 50. g of
o
aluminum (c = 0.9
C) from 68oC to 78oC?
9
its
own heat capacity (C) which is the amount of heat it
must absorb to change its temperature by one
degree. It is defined as C = q/ T. When 1.50 g of
methane burns in a bomb calorimeter (C = 11.3
kJ/oC) the temperature raises by 7.3oC. What is the
heat of combustion, in kJ/g, for CH4 based on this?
7 A 30.0 g sample of a liquid at 280. K is mixed with
50.0 g of the same liquid at 330. K. Calculate the
final temperature of the mixture, assuming no heat
is lost to the surroundings.
10 The heat capacity of a bomb calorimeter was
determined by burning 6.79 g of CH4 ( Hrxn = -802
kJ/mol CH4). The temperature increased by 10.8oC.
What is the heat capacity of the bomb, in kJ/oC?
11 Match the following terms with the definitions below.
endothermic
system
calorimetry
exothermic
heat
positive
energy
negative
Joule
a. The capacity to do work ......................................................................................................... _____________
b. Energy transferred due to temperature differences .............................................................. _____________
c. The experimental measurement of heat flow ........................................................................ _____________
d. The part of the universe under study, which exchanges energy with the surroundings ....... _____________
e. When a system absorbs heat, the sign is _______________ and the process is................... _____________
f. When a system releases heat, the sign is _______________ and the process is ................. _____________
g. The SI unit of heat (equal to 1 N m) is named after a famous physicist and brewer, James . _____________
5 67oC 6 450 J 7 311 K 8 50oC 9 -55 kJ/g 10 31.5 kJ/oC 11 see textbook
6.2
video
AP chem.
Thermochemistry, Solids & Liquids
Name ____________________
(lesson 2 AIM: How do we measure heats of reaction?)
Write the complete thermochemical equation for the combustion of methane:
12 What amount of heat is released when 5.8 g of
methane burns? ( Hcomb for CH4 = -890 kJ/molrxn)
H for a reaction is
the same whether it occurs in one step or several.
14 H2 + ½ O2 H2O
N2O5 + H2O 2 HNO3
½ N2 + (3/2) O2 + ½ H2
2 N2 + 5 O2 2 N2O5
HNO3
H = -285.8 kJ/molrxn
H = -76.6 kJ/molrxn
H = -174.1 kJ/molrxn
H=
13 The thermite reaction (below) uses iron(III) oxide
and aluminum to make molten iron. What mass of iron
is produced if a thermite reaction releases 25 kJ of
heat? 2 Al + Fe2O3 2 Fe + Al2O3 H = -850 kJ/molrxn
The sign changes if the reaction is _________________.
Multiply H by a factor if the ________________ change.
15 C + O2 CO2
H2 + ½ O2 H2O
C6H12O6 + 6O2 6 CO2 + 6H2O
6 C + 6 H2 + 3 O2 C6H12O6
H = -393.51 kJ/molrxn
H = -285.83 kJ/molrxn
H = -2803.02 kJ/molrxn
Germain
H=
Hess
HEATS OF FORMATION show the enthalpy change for the formation of 1 mole of a compound from elements.
C(s) + 2 H2(g) CH4(g)
= ________
H2(g) + ½ O2(g) H2O(g)
= ________
= np
(products) - nr
(reactants)
C(s) + O2(g) CO2(g)
= ________
Use heats of formation to calculate the heat of reaction for each reaction below:
16 CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)
17 liquid methanol methane gas + oxygen gas
BOND ENERGIES
18 Draw structures for CH4 + 2 O2
CO2 + 2 H2O & calculate the H:
H
C
N
O
S
F
Cl
Br
I
C=C
C=N
C=O
C=S
H
436
C
414
347
612
615
715
477
Single bond enthalpy (kJ/mol)
N
O
S
F
389
464
339
565
293
351
259
485
159
222
-272
138
-184
226
285
153
Multiple bond enthalpy (kJ/mol)
N=N
418
N=O
607
O=O
498
S=O
498
Cl
431
331
201
205
255
255
243
Br
368
276
243
201
213
255
218
193
I
297
218
-201
-277
209
180
151
820
890
1075
941
AP chem.
Thermochemistry, Solids & Liquids
PROBLEM SET 6.2
19
H for: 2 B + 3 H2 B2H6 given:
24
2 B + 3/2 O2 B2O3
H = 1273 kJ
H2 + O2
B2H6 + 3 O2 B2O3 + 3 H2O(g)
H = 2035 kJ
H2 + ½ O2 H2O(l)
H = 286 kJ
H2O(l) H2O(g)
H = +44 kJ
Name ____________________
H for the reaction:
H2O2
Bond energy (kJ/mol)
H H 432
H O 459
O O 207
O=O
494
25 (a) At STP, how many L H2(g) will release 550 kJ in:
H2(g) + ½ O2(g) H2O(g) H = -286 kJ/molrxn?
20 Use heats of formation values (below) along with
the reaction below to calculate
of C6H6:
2 C6H6 + 15 O2 12 CO2 + 6 H2O(l) H = -6534 kJ
(b) At STP, how many L CH4(g) will release 550 kJ in:
CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g), H = -891
kJ/molrxn?
Use
values (below) to find Hrxn for each reaction:
21 4 NH3(g) + 7 O2(g) 4 NO2(g) + 6 H2O(l)
26 Given the thermochemical equations below:
2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(l) H = -10900 kJ
H2(g) + ½ O2(g) H2O(g)
= -286 kJ
22 2 Al(s) + Fe2O3(s)
23 2 CH3OH(l) + 3 O2(g)
Al2O3(s) + 2 Fe(s)
How many L of gaseous H2 (at 1.0 atm and 25oC) are
required to furnish the same amount of energy as 80. L
of liquid gasoline, C8H18 (d = 0.74 g/mL)?
2 CO2(g) + 4 H2O(l)
19 +36 kJ 20 45 kJ/mol 21 -1396 kJ 22 -850. kJ 23 -1454 kJ 24 -199 kJ 25 (a) 43 L H2 (b) 14 L CH4 26 240,000 L H2(g)
6.3
AP chem.
Thermochemistry, Solids & Liquids
Name ____________________
(lesson 3 AIM: What causes substances do have different forces of attraction?)
27 DEMO: Consider the following substances. Which has the highest melting point? Which has the lowest?
aluminum metal
glycerol
glass
graphite
wax
olive oil
salt
TYPES OF SOLIDS
network covalent solids
metallic solids
ionic solids
INTERMOLECULAR FORCES (only for molecular compounds!)
1. Dispersion forces
2. Dipole forces
molecular solids
3. H-bonding
28 For each structure below, state its most significant intermolecular force (IMF):
H H
H
O
H
b.p. = 100oC
H C C O H
H H
b.p. = 78oC
H H
H H
H C C O C C H
H H
H H
b.p. = 35oC
29 Draw a structure for each and state the significant IMFs. Pick two and explain the osbserved boiling points.
b.p. = -33oC
77oC
-100oC
61oC
167oC
-79oC
-64oC
65oC
NH3
CCl4
BF3
CHCl3
PCl5
CO2
SF6
CH3OH
video
AP chem.
Thermochemistry, Solids & Liquids
Name ____________________
PROBLEM SET 6.3
30 Identify the most important types of intermolecular / interparticle forces present in each of the following:
HCl
HF
CaCl2
CH4
CO
NaNO3
31 TYPES OF SOLIDS Use the following terms to complete the paragraph below:
electrons
ionic
network
metallic
amorphous
molecular
crystalline
Ice, iodine, and salicylic acid are all covalent solids, and can thus be classified as __________ solids.
NaCl, KNO3, and CuSO4, on the other hand, are __________ solids, and as such have ionic forces of attraction
which are much stronger than any of the intermolecular forces. Copper, iron and tungsten are all examples of
_______ solids, and can conduct heat and electricity due to their mobile ____________. Quartz, diamond and
graphite are common examples of ____________ solids, and have a continuum of covalent bonds. Solids such
as glass and wax with highly disordered structures are called _______________ solids, as opposed to solids such
as diamond, salt and ice, which have highly ordered structures, and are classified as _____________ solids.
32 ALLOYS & SEMICONDUCTORS
interstitial
quartz
SiO2
Use the following terms to complete the paragraph below:
graphite
semiconductor
glass
substitution
diamond
Alloys in which atoms are replaced by a different atom of similar size (such as in brass, sterling silver, &
pewter) are known as ___________________ alloys. Alloys in which additional atoms occupy spaces between
atoms (such as in steel) are known as ____________________ alloys. The hardest known natural substance is
________________. This differs greatly from the more common allotrope of carbon, ________________, which,
unlike diamond, can conduct due to its overlapping pi electrons. Silicon-based solids are primarily composed of
silica, which is _____. The crystalline form of silica is _________, and the amorphous form of silica is ______.
Many silicon compounds are manufactured to contain sites in which a silicon atom is replaced with an atom
having more or less valence electrons (such as an arsenic atom). This creates a ______________________,
which has an intermediate conductivity, in between that of a pure metal conductor, or a perfect insulator.
33 ALLOYS homogeneous mixtures that behave like metals
Steel is an alloy that exhibits both interstitial and substitutional properties. It consists of iron with a small
amount of carbon and chromium to make the steel less likely to rust. Chromium atoms react with oxygen in the
air to form a nonreactive layer of chromium oxide on the surface, preventing further oxidation of underlying
iron. Which of the following diagrams best shows a particle-level view of a surface section and an interior section
of the alloy?
A
B
C
D
(a)
(b)
(c)
Fun fact:
Samurai swords
consist of
different types
of steel (iron
and carbon).
The core is
made of lowcarbon steel,
and the outer
layer is made of
high-carbon
steel. The
carbon gives it
much more
strength than
the iron would
have by itself.
6.4
AP chem.
Thermochemistry, Solids & Liquids
Name ____________________
(lesson 4 AIM: How do forces of attraction affect the physical properties of substances?)
STRONG FORCES OF ATTRACTION?... high mp, high bp, low Pvap, high surf ten!
PHASE DIAGRAMS
Triple point:
Critical point:
VAPOR PRESSURE
Pvap
temp
CALCULATING CHANGES IN BOILING POINT / VAPOR PRESSURE
34 At what temperature, in oC, would water boil at when the pressure is 23.8 torr,
Hv = 43.9 kJ/mol for water?
video
AP chem.
PROBLEM SET 6.4
Thermochemistry, Solids & Liquids
35 In each pair, which substance is expected to
have the higher boiling point?
Name ____________________
37 Draw a structural formula for each of the
following and determine which should have the
highest boiling point. Briefly justify your reasoning.
a. HF or HCl
CH3CH2OH
b. HCl or LiCl
CH3OCH3
c. C5H12 or C6H14
CH3CH2CH3
d. Br2 or F2
36 Answer the following. Briefly justify your reasoning.
a. Which has the higher boiling point:
HBr or Kr
b. Which has the higher freezing point:
NaCl or HF
c. Which has the lower vapor pressure at 25oC:
Br2 or I2
d. Which has the lower boiling point:
CH3CH3 or CH3CH2CH3
e. Which has the higher boiling point:
HCl or HBr
f. Which has the lower vapor pressure at 25oC:
CH3CH2CH3 or CH3COCH3
38 LIQUIDS Use the following terms to complete the paragraph:
viscosity
adhesive
surface
convex
cohesive
capillary
temperature
_____________________, which typically decreases with an increase in ___________________. The rising of a
attraction between molecules and a container, also known as __________________ forces. These differ from
the attractive forces which exist between similar molecules, known as _______________________ forces. A
column of mercury in a glass tube, for example, gives a ___________ meniscus because its cohesive forces for
itself are stronger than its adhesive forces for the glass.
39 PRACTICE AP QUESTION
White gold is a common substitutional alloy of gold (atomic radius =
135 pm) and palladium (atomic radius = 140 pm). Imagine a ring
made of white gold that is 75 mole % gold and 25 mole % palladium.
Draw a particle-level diagram of the alloy consisting of 12 atoms
total. Use empty circles for gold and shaded circles for palladium.
40 PRACTICE AP QUESTION
Consider nonane and 2,3,4-trifluoropentane, shown below, with their respective boiling points.
H H H H H H H H H
H
F
F
F
H
H C C C C C C C C C H
H C C C C C H
H H H H H H H H H
H H H H H
151oC
b.p. =
(a) What type of intermolecular forces are present
in each compound?
b.p. = 89oC
(b) Explain the difference in boiling points.
6.5
AP chem.
1.
2.
3.
4.
6.
9.
10.
From the table below, determine the enthalpy change for the reaction:
2 H2(g) + O2(g) 2 H2O(g)
Bond
energy (kJ/mol)
H H
436
O=O
499
H O
464
a.
0 kJ
b. 485 kJ
c.
-485 kJ
d. 463 kJ
11.
e.
443 kJ
Name ____________________
If 10 g of a liquid at 300 K is heated to 350 K by adding 6 kcal, what is
the specific heat of the liquid?
a.
6 cal/goC
b. 12 cal/goC
c.
60 cal/goC
d. 600 cal/goC
e.
120 cal/goC
Given that the reaction below has a change in enthalpy of -889.1 kJ,
calculate the heat of formation of one mole of CH 4.
CH4 + 2 O2 CO2 + 2 H2O
Hof of H2O = -285.8 kJ/mol
Hof of CO2 = -393.3 kJ/mol
a.
-210.0 kJ/mol
b. -107.5 kJ/mol
c.
-76.0 kJ/mol
d. +76.0 kJ/mol
e.
+210.0 kJ/mol
When 58 g H2O (c = 4.18 J/g°C) is heated from 275 K to 365 K,
a.
absorbs 21,820 J
b. absorbs 377 J
c.
releases 5220 J
d. absorbs 242 J
e.
releases 90 J
Given the equation: C2H4 + 3 O2 2 CO2 + 2 H2O(g), H = -1323 kJ,
what is the value of H if liquid water is made instead of water vapor?
(Note: H2O(g) H2O(l), Hv = -44 kJ)
a.
-1235 kJ
b. -1279 kJ
12. A student uses a coffee-cup calorimeter in an experiment. The cup is
c.
-1323 kJ
weighed, filled halfway with warm water then weighed again. The
d. -1367 kJ
temperature of the water was measured, and some ice cubes from a
e.
-1411 kJ
0oC ice bath were added to the cup. The mixture was gently stirred as
the ice melted, and the lowest temperature reached by the water was
The enthalpy of formation for potassium chloride is given by:
recorded. The cup and its contents were weighed again. The purpose of
a.
K(g) + ½Cl2(g) KCl(g)
weighing the cup and its contents again at the end of the experiment
b. K+(g) + Cl-(g) KCl(s)
c.
d.
e.
5.
Thermochemistry, Solids & Liquids
From the given data, determine the heat of formation of one mole of
carbon dioxide in kJ/mol:
2 C + O2 2 CO
H = -220 kJ
2 CO + O2 2 CO2
H = -560 kJ
a.
(-220) + (-560)
b. ½[(-220) + (-560)]
c.
(-220) (-560)
d. ½[(-220) (-560)]
e.
- (-220) + (-560)
2K(s) + Cl2(g)
K(s) + ½Cl2(g)
K+(g) + Cl-(g)
2KCl(s)
KCl(s)
KCl(g)
For which of these processes is H expected to be negative?
i. temperature increases when CaCl2 dissolves in water
ii. steam condenses to liquid water
iii. water freezes
iv. dry ice sublimes
a.
iv only
b. i, ii and iii
c.
i only
d. ii and iii only
e.
i and ii only
Which has a value of Hof 0?
a.
F2(g)
b. Br2(g)
c.
I2(s)
d. C(s, graphite)
e.
N2(g)
7.
Which equation shows the formation of CH3OH(l) from its elements?
a.
CH3OH(l) + (3/2) O2(g) CO2(g) + 2 H2O(l)
b. CH3OH(l) + (3/2) O2(g) CO2(g) + 2 H2O(g)
c.
2 CH3OH(l) + 3 O2(g) 2 CO2(g) + 4 H2O(l)
d. CH3OH(l) C(s) + 2 H2O(l)
e.
C(s) + 2 H2(g) + ½ O2(g) CH3OH(l)
8.
16 g H2O freezes at 273 K. What amount of heat is involved? ( Hfus =
6.02 kJ/mol for H2O)
a.
-16 J
b. -4368 J
c.
-18,258 J
d. -350 J
e.
-5351 J
a.
b.
c.
d.
e.
mass of ice added
mass of the thermometer
mass of water that evaporated
mass of water that was cooled
true mass of the calorimeter cup
13. Gaseous H2 and F2 combine in the reaction below to form HF with an
enthalpy change of -540 kJ. What is the value of the heat of formation
of one mole of HF(g)?
H2(g) + F2(g) 2 HF(g)
a.
-1080 kJ/mol
b. -540 kJ/mol
c.
-270 kJ/mol
d. +270 kJ/mol
e.
+540 kJ/mol
14. Based on the information below, what is H for the reaction:
CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = ?
C(s) + 2H2(g) CH4(g) H = x
C(s) + O2(g) CO2(g) H = y
H2(g) + ½O2(g) H2O(l) H = z
a.
x+y+z
b. x + y z
c.
z + y 2x
d. 2z + y x
e.
2z + y 2x
15. The densities of certain compounds are greater as liquids than as solids
(e.g. H2O). This means that increasing the pressure will result in
a.
a solid becoming a liquid.
b. a liquid becoming a solid.
c.
lowering of the critical temperature.
d. elevation of the freezing temperature.
e.
lowering of the triple point.
AP chem.
Thermochemistry, Solids & Liquids
16. The liquefied hydrogen halides (HX) have the normal boiling points
given below. The relatively high boiling point of HF can be explained:
HF = +19 oC
HCl = -85 oC
HBr = -67 oC
HI = -35 oC
a.
HF gas is more ideal
b. HF is the strongest acid
c.
HF molecules have a smaller dipole moment
d. HF is much less soluble in water
e.
HF molecules tend to form hydrogen bonds
Name ____________________
22. The boiling points of He, Ne, Ar, Kr, and Xe increase in that order.
What accounts for this?
a.
The London (dispersion) forces increase
b. The hydrogen bonding increases
c.
The dipole-dipole forces increase
d. The chemical reactivity increases
e.
The number of neighbors increases
23. The graph here shows the temperature of a
pure substance as it is heated at a constant
rate. The substance is boiling point at
17. The normal boiling point of the substance in the phase diagram here is 24. Which likely has the highest melting point?
a.
S8
b. I2
c.
SiO2
d. SO2
e.
C6H6
a.
b.
c.
d.
e.
-15 oC
-10 oC
140 oC
Greater than 140 oC
Not determinable from the graph
18. The critical temperature is the
a.
T at which Pvap of the liquid is equal to the external pressure
b. T at which the Pvap of the liquid is equal to 760 mmHg
c.
T at which the solid, liquid, and vapor phases are all in equilibrium
d. T at which liquid and vapor phases are in equilibrium at 1 atm
e.
Lowest T above which a substance cannot be liquefied at any
applied pressure
25. In methane, the forces between neighboring molecules are
a.
ionic bonds
b. covalent bonds
c.
hydrogen bonds
d. ion-dipole forces
e.
London (dispersion) forces
26. Two closed containers of acetone are below, each with same volume of
acetone at the same temperature. The vapor pressure of the acetone is
a.
b.
c.
d.
e.
higher in 1 because the surface area of the liquid is greater
higher in 1 because the volume of the vapor is greater
lower in 1 because the level of the liquid is lower
the same in both because the volume of the liquid is the same
the same in both because the temperature is the same
19. Which of the following is true at the triple point of a pure substance?
a.
The Pvap of the solid phase equals the Pvap of the liquid phase.
b. The T is always 0.01 K lower than the normal melting point.
c.
The liquid and gas phases of the substance always have the same
density and are therefore indistinguishable.
27. Which explains why the boiling point of CCl4 is higher than that of CF4?
d. The solid phase always melts if P increases at constant T.
a.
The C Cl bonds in CCl4 are less polar than the C F bonds in CF4
e.
The liquid phase always vaporizes if P increases at constant T.
b. The C Cl bonds in CCl4 are weaker than the C F bonds in CF4
c.
The mass of the CCl4 molecule is greater than that of the CF4
20. Which of the following actions would be likely to change the boiling
molecule
point of sample of a pure liquid in an open container?
d. The electrons in CCl4 are more polarizable than in CF4
i. Placing it in a smaller container
e.
The bonds in CCl4 are covalent, whereas ionic in CF4
ii. Increasing the number of moles of the liquid in the container
iii. Moving the container and liquid to a higher altitude
28. Which exhibits H-bonding?
a.
i only
a.
CH2F2
b. ii only
b. N2H4
c.
iii only
c.
CH3OCH3
d. ii and iii only
d. C2H4
e.
ii, ii, and iii
e.
CH
2 2
21. The phase diagram for a pure substance is shown below. Which point 29.
on the diagram corresponds to the equilibrium between the solid and
a.
liquid phases at the normal melting point?
b.
c.
d.
e.
Solid silver melts
Solid potassium chloride melts
Solid carbon (graphite) sublimes
Solid iodine sublimes
Glucose dissolves in water
30. Which shows a correct vapor pressure trend?
a.
CH4 < C2H5OH < C2H5OC2H5 < Ne
b. Ne < CH4 < C2H5OC2H5 < C2H5OH
c.
C2H5OH < C2H5OC2H 5 < CH4 < Ne
d. C2H5OC2H5 < C2H5OH < CH4 < Ne
e.
C2H5OC2H5 < C2H5OH < Ne < CH4
ANSWERS
1. B
2. C
3. E
4. D
5. B
6. B
7. E
8. E
9. B
10. C
11. A
12. A
13. C
14. D
15. A
16. E
17. C
18. E
19. A
20. C
21. C
22. A
23. D
24. C
25. E
26. E
27. D
28. B
29. C
30. C
AP Chem
Thermochemistry
Practice Test 06
Name ____________________
FREE RESPONSE – Answer the following in the space provide below. You must show work for full credit.
The boiling points, dipole moments, and polarizabilities of three hydrogen halides are given in the table below.
A)
(5 points) Based on the data in the table, what type of intermolecular force among the molecules HCl, HBr, and HI is able to
account for the trend in boiling points? Justify your answer.
B)
(5 points) Based on the data in the table, a student predicts that the boiling point of HF should be 174 K. The observed boiling
point of HF is 293 K. Explain the failure of the student’s prediction in terms of the types and strengths of the intermolecular
forces that exist among HF molecules.
C)
(5 points) A representation of five molecules of HBr in the liquid state is shown in Box 1 below. In Box 2, draw a representation
of the same five molecules of HBr after complete vaporization has occurred.
MULTIPLE CHOICE (5 points each) – Bubble the letter of the best answer to each question.
1. Which of the following arranges the molecules N2, O2, and
3. Which liquid possesses dipole-dipole intermolecular forces?
F2 in order of their bond enthalpies (aka bond energies),
a. F2(l)
from least to greatest (aka weakest to strongest)?
b. CH4(l)
a. F2 < O2 < N2
c. CF4(l)
b. O2 < N2 < F2
d. CH2F2(l)
c. N2 < O2 < F2
d. N2 < F2 < O2
4. Which process is exothermic?
a. condensation
2. Based on the information in the table below, which liquid,
b. melting
CS2(l) or CCl4(l), has the higher equilibrium vapor pressure
c. sublimation
at 25oC, and why?
d. vaporization
molar mass (g/mol)
boiling point (oC)
5. A sample of a solid binary compound at room temperature
CS2(l)
76
46.5
does not conduct electricity but becomes conductive
CCl4(l)
154
76.7
when dissolved in water. Which type of interactions are
a. CS2(l); it has stronger London dispersion forces
likely found in this substance?
b. CS2(l); it has weaker London dispersion forces
a. Ionic bonds
c. CCl4(l); it has stronger London dispersion forces
b. Metallic bonds
d. CCl4(l); it has weaker London dispersion forces
c. Covalent bonds
d. Hydrogen bonds
AP Chem
6.
7.
8.
9.
Thermochemistry
At room temperature I2(s) is a molecular solid. Which of
the following provides a characteristic of I2(s) with a
correct explanation?
a. It has a high melting point because it has weak
intermolecular forces.
b. It is hard because it forms a three-dimensional
covalent network.
c. It is not a good conductor of electricity because its
valence electrons are localized in bonding and
nonbonding pairs.
d. It is very soluble in water because its molecules are
polar.
Name ____________________
11. Use the thermodynamic information to find the mising ∆H:
½ N2(g) + ½ O2(g) → NO(g)
∆H = +90. kJ/molrxn
½ N2(g) + O2(g) → NO2(g)
∆H = +35 kJ/molrxn
2 NO2(g) → N2O4(g)
∆H = -60 kJ/molrxn
2 NO(g) + O2(g) → N2O4(g)
∆H = ? kJ/molrxn
a. -170. kJ
b. -115 kJ
c. +115 kJ
d. +170. kJ
12. Determine the enthalpy change when 5.00 g of Fe2O3(s)
(MM = 159.70 g/mol) reacts with excess Al(s) in the
reaction: Fe2O3(s) + 2 Al(s) → Al2O3(s) + 2 Fe(l)
5.00
A student mixes a 10.0 mL sample of 1.0 M NaOH(aq) with
a.
-825.2
4 kJ Enthalpy of formation, ∆Hof
159.70
a 10.0 mL sample of 1.0 M HCl(aq) in a polyester container.
5.00
o
b.
-837.8 kJ Fe2O3(s)
-825.5 kJ/mol
The temperature of the solutions before mixing was 20.0 C.
159.70
2 x 5.00
-1675.7 kJ/mol
If the final temperature of the mixture is 26.0oC, what is
c.
-837.8 kJ Al2 O3(s)
159.70
the experimental value of ∆Hrxn? (Assume the solution
Fe(l)
+12.4 kJ/mol
159.70
d.
-825.2
4 kJ
mixture has a specific heat of 4.2 J/(g•K) and a density of
5.00
1.0 g/mL.)
a. -50. kJ/molrxn
13. The critical temperature of water is the…
b. -25 kJ/molrxn
a. temperatrue at which solid, liquid, and gaseous water
4
c. -5.0 x 10 kJ/molrxn
coexist.
d. -5.0 x 102 kJ/molrxn
b. tempearture at which water vapor condenses.
c. maximum temperature at which liquid water can
When water is added to a mixture of Na2O2(s) and S(s), a
exist.
redox reaction occurs according to:
d. minimum temperature at which water vapor can exist.
2 Na2O2 + S + 2 H2O → 4 NaOH + SO2
∆H298 = -610 kJ/molrxn
14. Use bond energies to calculate ∆H for the reaction:
Two trials are run, using excess water. In the first trial, 7.8
H2 + O2 → H2 O2
Bond energy (kJ/mol)
g of Na2O2 (78 g/mol) is mixed with 3.2 g S (32 g/mol). In
a. -521 kJ
H—H 432
the second trial, 7.8 g Na2O2 is mixed with 6.4 g S. The
b. -486 kJ
H—O 459
Na2O2 and S react as completely as possible. Both trials
c. -199 kJ
O—O 207
yield the same amount of SO2. Which of the following
d. +199 kJ
O=O
494
identifies the limiting reactant and the heat released, q, for
the two trials at 298 K?
15. Under certain conditions CO2 melts rather than sublimes.
a. Limiting reactant = S
q = 30. kJ
To which transition in the phase diagram does this change
b. Limiting reactant = S
q = 61 kJ
correspond?
c. Limiting reactant = Na2O2
q = 30. kJ
a. A → B
d. Limiting reactant = Na2O2
q = 61 kJ
b. A → C
c. B → C
Which reaction correctly depicts the enthalpy change of
d. C → B
formation of sodium carbonate?
3
16. Which liquid has the highest vapor pressure at 25oC?
a. 2 Na(s) + C(s) + O2(g) → Na2CO3(s)
2
a. Butane, C4H10
b. Glycerol,
b. Na2O(s) + CO2(g) → Na2CO3(s)
H H H H
C3H5(OH)3
c. 2 Na+(aq) + CO32-(aq) → Na2CO3(s)
d.
H C C C C H
H H H H
2 Na+(aq) + 2 OH-(aq) + CO2(aq) → Na2CO3(s) + H2O(l)
10. The reaction below takes place in a rigid, insulated vessel
that is initially at 600 K:
CH3OH(g) → CO(g) + 2 H2(g); ∆H = +91 kJ/molrxn
Which of the following statements about the bonds in the
reactants and products is most accurate?
a. The sum of the bond enthalpies of the bonds in the
reactant is greater than the sum of the bond
enthalpies of the bonds in the products.
b. The sum of the bond enthalpies of the bonds in the
reactant is less than the sum of the bond enthalpies
of the bonds in the products.
c. The length of the bond between carbon and oxygen
in CH3OH is shorter than the length of the bond
between carbon and oxygen in CO.
d. All the bonds in the reactant and products are polar.
c.
Octane, C8H18
H H H H H H H H
H C C C C C C C C H
H H H H H H H H
H H H
H C C C H
OH OH OH
d.
Propanol, C3H7OH
H H H
H C C C OH
H H H
17. Of the species below, the lowest melting points overall
occur for members of which class of solids?
a. Ionic
c. Metallic
b. Molecular
d. Network covalent
18. (EXTRA CREDIT) An example of an interstitial alloy is
shown here. Compared with the pure metal, how would
you expect the properties of the alloy to vary?
a. The alloy has higher malleability and lower density.
b. The alloy has lower malleability and higher density.
c. The alloy has higher malleability and higher density.
d. The alloy has lower o’malleability and lower density.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
A
B
D
A
A
C
A
C
A
A
A
A
C
C
A
A
B
B