PRACTICE EXAM for Unit 8 (Reaction Dynamics)
Text references: Chapter 17, sections 2-4, and Chapter 18, section 1. Useful information:
1 kJ  1000 J
q  mc² T
cH 2 O  4.184 J g C
PV  nRT
P1V1 P2V2

n1T1 n2T2
Standard molar volume = 22.4 L mol
R  0.08206
atm L
mol K
1 atm  760 mmHg
1L  1000mL
1. a. What do we call an instrument for capturing and measuring heat energy?
b. Define specific heat, including the symbol used to represent it.
2. a. When 127.68 g of brass at 163 °C are added to 150.00 g H2O at 22.4 °C, the final temperature of the
brass and water is 32.6 °C. Calculate the value of the specific heat, c, for brass.
b. When 225 g of stainless steel at 183.0 °C are added to 125 g H2O at 23.2 °C, the final temperature
of the steel and water is 51.5 °C. Calculate the value of the specific heat, c, for stainless steel.
3. a. What is enthalpy (∆H)?
b. What is the sign of ∆H for an exothermic reaction? Where is the energy written in the
thermochemical equation for an exothermic process? Where does the released energy come from?
Would an exothermic reaction feel hot or cold to us?
c. What is the sign of ∆H for an endothermic reaction? Where is the energy written in the
thermochemical equation for an endothermic process? Where does the energy go? Would an
exothermic reaction feel hot or cold to us?
4. a. A piece of Mg weighing 0.043 g is placed in an insulated cup containing 75.3 g HCl solution. As
the Mg dissolves, the temperature of the HCl solution increases from 22.0 °C to 24.6°C. Calculate
∆H for this reaction, in kJ/mol and classify it as endothermic or exothermic.
b. When 8.3 g KI are added to 50.0 g water, the temperature changes from 18.4 °C to 14.2 °C as the
salt dissolves. Calculate ∆H (in kJ/mol) for the reaction and classify it as endothermic or
exothermic.
5. Use this balanced thermochemical equation to answer the following questions:
2 C4H10 (g) + 13 O2 (g)  8 CO2 (g) + 10 H2O (l) + 5754 kJ
a. Is the reaction exothermic or endothermic?
b. How much energy, in kJ, is produced when 1.325 g C4H10 burn?
c. How many L of CO2 gas, measured at STP, will form when 2500 kJ of energy are produced?
d. How many grams of O2 are needed to produce 850 kJ of energy?
6. a. What is Hess' Law?
b. Reactions (1) and (2) combine to give reaction (3):
(1) ½ N2 (g) + ½ O2 (g)  NO (g)
∆H = +90.25 kJ
(2) NO (g) + ½ O2 (g)  NO2 (g)
∆H = –57.07 kJ
(3)
½ N2 (g) + O2 (g)  NO2 (g)
∆H = ?
i. What chemical formula cancels out when you combine reactions (1) & (2)?
ii. What is ∆H for reaction (3)?
7. a. Use ∆H°f values from your table to calculate ∆H° for this reaction:
2 H2S (g) + SO2 (g)  3 S (s) + 2 H2O (g)
b. Use ∆H°f values from your table to calculate ∆H° for this reaction:
4 CO (g) + 8 H2 (g)  3 CH4 (g) + CO2 (g) + 2 H2O (l)
* * * * * * * * * * * * * * * * * * * * * * * *
8. On a blank periodic table, fill in the period numbers, group numbers, and staircase; label the alkali
metal, alkaline earth, halogen, and noble gas families; indicate the metals and nonmetals,
semimetal/metalloid and transition metal regions; add the trends in atomic size and ionization energy.
9. a. What is the charge on a cation? How does its size compare to its parent atom?
b. What is the charge on an anion? How does its size compare to its parent atom?
10. Indicate the number of valence electrons in each atom, and draw its dot structure.
a. Br
b. Rb
c. As
11. a. Draw Lewis dot structures to show the formation of the salt NaBr.
b. Which of these drawings best represents the bonding in NaBr?
a
b
c
d
e
12. Draw Lewis dot structures for each molecular species, including the number of valence electrons.
a. CH2Br2
b. COS
c. O3
d. ClO1–
Which of the drawings above best represents a molecular substance?
13. Balance each reaction in lowest terms.
a. FeS2 (s) + O2 (g)  SO2 (g) + Fe2O3 (s)
b. NO2 (g) + H2 (g)  NH3 (g) + H2O (l)
c. Na (s) + H2O (l)  NaOH (aq) + H2 (g)
14. Do these mole conversions and stoichiometry calculations WITHOUT your calculator. [Hint: set
each one up, round molar masses to whole numbers, then look for simple math.]
a. What is the molar mass of (NH4)2SO4?
b. What is the mass of 0.25 mol Ne?
c. How many moles are in 0.64 g S?
d. What is the volume of 0.25 mol O2 gas at STP?
e. How many moles of CO2 are in 11.2 L CO2 at STP?
f. How many grams of H2 form when 0.52 g Cr react?
2 Cr (s) + 6 H1+ (aq)  2 Cr3+ (aq) + 3 H2 (g)
g. How many moles of H2O are produced when 0.15 mol CH4 burn?
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l)
h. What is the molarity of a solution that contains 0.050 mol sugar in 25.0 mL solution?
15. Do these gas law problems. (DO use your calculator!)
a. 325 mL of air at room pressure (765 mm Hg) are compressed with a piston to a volume of 42 mL.
What is the pressure of the compressed air? The temperature and amount of air remain constant.
b. A He weather balloon has a volume of 18 L at the earth's surface, where the temperature is 25 °C
and the pressure is 750 mm Hg. What is the volume of the balloon when it reaches an altitude
where the temperature is –22 °C and the pressure is 375 mm Hg? The amount of He is constant.
c. What is the volume of 0.085 mol hydrogen gas at 737 mm Hg and 21 °C?
d. A sample of nitrogen gas in a 275 mL container at 0.82 atm pressure and 26 °C is transferred to
larger container with volume 750 mL . The temperature in the larger container is now 17 °C.
What is the pressure of the gas in the larger container? The amount of nitrogen is constant.
16. Explain each of these in molecular terms:
a. gas pressure
b. gas temperature
c. why you will eventually smell a gas that was released at the other end of the room
d. why gases can be compressed but liquids and solids cannot
17. a. Draw a heating curve for water (i.e., a graph with T on y and energy added on x) from –20 °C to
+120 °C. Label the curve to show solid, liquid, gas, melting, freezing, and boiling.
b. When the water is melting or boiling, is it absorbing energy, releasing energy, or staying at the
same energy?
c. When the water is freezing, is it absorbing energy, releasing energy, or staying at the same energy?
d. How will the graph change if you dissolve something like salt or sugar in the water?
Answers to Practice exam for Unit 8 (Reaction Dynamics)
1. a. calorimeter
b. energy to change the temperature of one gram of substance by 1°C (symbol c)
2. a. 0.385 J/g °C
b. 0.500 J/g °C
3. a. energy stored in chemical bonds
b. ∆H negative for exothermic; energy written as product (on right), energy comes from bonds.
Exothermic reaction feels hotter as heat comes out of the bonds.
c. ∆H positive for endothermic; energy written as reactant (on left), energy goes into bonds.
Endothermic reaction feels colder as heat is drawn into bonds from the surroundings.
4. a. – 460 kJ/mol (exothermic)
b. +20.5 kJ/mol (endothermic)
5 a. exothermic
c. 78 L
b. 65.59 kJ
d. 61 g
6. a. You can add together thermochemical equations (including energies) to get new thermochemical
equations.
b. i. NO (g) ii. +33.18 kJ
7. a. – 145.5 kJ
b. ∆H° = –747.5 kJ
* * * * * * * * * * * * * * * * * * * * * * * *
8 & 9. See textbook pg 162, 170-172.
13. a. 4, 11, 8, 2
b. 2, 7, 2, 4
14. a. 132 g/mol
b. 5 g
15. a. 7.9 atm
b. 31 L
10–12. Check my drawings in class
c. 2, 2, 2, 1
c. 0.02 mol
c. 2.1 L
d. 5.6 L
e. 0.5 mol
f. 0.03 g
g. 0.3 mol
h. 2 M
d. 0.29 atm
16. a. Gas P is caused by collisions of gas molecules with container wall.
b. Gas T is proportional to molecular kinetic energy, which is the product of mass x speed2; as T
increases, speed increases.
c. Molecules are in constant random motion, so they eventually distribute throughout entire container,
no matter how large it is.
d. Gas particles are very far apart (gases are mostly empty space), so the particles can be squeezed
together easily. Liquids and solids have no significant space between the molecules, so they can’t
be squeezed together any more.
temperature (°C)
17. The curve should look like the one at right, with melting/freezing
(BC) at 0 °C and boiling (DE) at 100 °C. Melting and boiling are
endothermic (absorbing energy, even though temperature is not
changing). Freezing is exothermic (releasing energy, even though
temperature is not changing). A solute like salt or sugar would
cause BC to be at a lower temperature and DE to be at a higher
temperature.
F
D
E
B
C
A
energy added 