Honors Chemistry
Midterm Review
Name __________________________________
Period ______ Date ______/______/______
1. Make the following conversions:
a. 1.9 L to mm3
b. 1.83 x 106 µm3 to ft3
2. Complete the following calculations and report the answer in scientific notation and with the correct
number of significant digits.
a. 23.1 + 4.77 + 125.39 + 3.581
c. 561.0 x 34,908 x 23.0
b. 22.101 - 0.9307
d. 3.4617 x 107 ÷ 5.61 x 10¯4
3. A block of lead has dimensions of 4.50 cm by 5.20 cm by 2.36 in. The block weighs 1591 g.
a. Calculate the density of lead in pounds per liter.
b. A sample of the lead is added to a graduated cylinder containing 45.5 mL of water and the
water level rises to the 49.1 mL mark on the cylinder. Determine the number of protons in the
sample.
4. A .59 kg brass candlestick has an initial temperature of 98.0°C. When 5043 calories of heat are
removed from the candlestick, its temperature lowers to 6.8°C. Determine the specific heat capacity
of brass?
5. Identify as an element, compound, heterogeneous mixture or solution. List one method to separate the
substance.
a. ___________________ sodium carbonate
b.___________________ sugar water
c. ___________________ carbon
d.___________________ rocks and sand
6. Draw the phase change diagram for water. Label the phase changes, states of matter, BP and MP.
a.
b.
c.
d.
In which portions of the graph does more than one state of matter exist?
List two types of phase changes that are endothermic.
List two types of phase changes that are exothermic.
Explain why there is no temperature increase during a phase change.
7. Describe the following types of separation techniques: distillation, filtration and crystallization.
8. Physical or chemical change?
a.___________ burning magnesium
b. __________melting ice
c.___________ powdering a sugar cube
d. __________ iron rusting
9. Calculate the mass of water at 20.0°C needed to lower the temperature of 750.0 g of water at 75.0°C
to body temperature, 37.0oC.
10. Calculate the amount of heat, in kJ, released from the conversion of 75.0 g of steam at 125.0°C to ice
at –15.0°C.
11. Calculate the energy, in kJ, required to melt a tray of ice cubes at 0°C. One ice cube has a mass of
62.0 g and the tray contains 16 ice cubes.
12. Calculate the final temperature if 55.0 g of liquid water at 90.0°C is added to 35.0 g of liquid water at
15.0°C.
13. When 15.4 g of silver at 75.0°C is added to 45.0 g of water in a calorimeter, the final solution reaches
a temperature of 35°C. What was the initial temperature of the water in the calorimeter?
(cAg = 0.235 J/g°C)
14. At constant pressure, a gas occupies 23.5 mL at 125 °C. At what temperature will it occupy 15.6 in3?
15. A gas occupies 1.96 dm3 at STP. At what pressure, in Pascals, will it occupy 895 mL at 64.5 °C?
16. A gas is collected by water displacement at 23°C and the pressure is equilibrated with the air. The
eudiometer reads 45.0 mL and the barometer reads 72.3 cm Hg. Convert volume to STP.
17. A balloon is filled to a volume of 758 mL with 169 grams of steam at 125°C. After a brief period of
cooling, the balloon’s volume is measured to be 745 mL. How much heat energy did the steam lose?
18. Complete the following table:
Element
106
46
# of
Protons
# of
Electrons
# of Neutrons
Mass #
Atomic #
Pd 4 
Xenon – 131
As3-
75
21
52
24
19. The nuclear transmutation of mercry-210 that consists of 2 alpha decays and 1 beta decay.
20. The electron capture of antimony-123 followed by the emission of 2 neutrons.
21. What is the resulting nuclide from the beta decay of radon-226?
22. What is the original isotope of an alpha decay that results in plutonium-244?
23. Gold-191 has a half-life of 12.4 hours. What mass of this isotope would remain after 49.6 hours if
you started with a 7.50 μg sample of pure gold-191?
24. Fluorine-18 is used to scan brains. Its half-life is 1.8295 hours. If only 0.0035 μg of the fluorine-18
sample remained after sitting on a shelf in the hospital for 5.5 hours, how much was originally
present?
25. What is the mass of the electrons, in g, in 4.5 x 1024 atoms of carbon? What is the charge of the
electrons?
26. What is the mass, in lbs, of 3.12 x 1025 atoms of sulfur?
27. Determine the percent composition of aluminum nitrate.
28. Which gas, nitrogen or carbon dioxide, will effuse more rapidly?
29. What is the molar mass of a gaseous element if at room temperature it effuses through a pinhole 2.16
times as rapidly as xenon? Which diatomic element could it be?
30. Complete the following table:
Name
Formula
Oxalic acid
Zinc chloride tetrahydrate
Hypochlorous acid
Xenon hexafluoride
Hydrosilicic acid
Diboron tetrabromide
Silver arsenate
Vanadium (V) dichromate
Aluminum iodite
Calcium carbonate monohydrate
Chromium (III) bicarbonate
Iron (II) thiosulfate
Carbonic acid
Pd(NO2)2
H2CrO4 (aq)
CuCl2·2H2O
NH3
V(ClO4)4
(NH4)3BO2
IF5
FeCl3·6H2O
H3P (aq)
CoCO3
FePO3
Mn(NO3)3
HClO4 (aq)
CoCl2·6H2O
Molecular Formula
Structural Formula
Name
4,6-dimethylnonane
3,3-dimethylhexane
2,3,4-trichloropentane
5-ethyl-6-iodo-2,2-dimethyl3-octene
4-chloro-1-fluoro-2-pentene
CH3
CH3
CH2 CH CH3
CH3 CH2 CH2 C
CH3
CH3
CH3
CH2
CH3
CH3 CH CH CH CH CH CH3
Cl
Br
CH3
CH3
CH2
CH3 CH CH
C
CH CH3
Br
Cl
CH3 CH C
CH
CH3 Br
CH3
CH3 CH2 CH CH C
CH3
CH3
CH CH2 CH3
____
1. What is the SI unit of energy?
a. Joule
b. ° Fahrenheit
c. Kelvin
d. calorie
____
2. Which of the following is a statement of the law of conservation of energy?
a. Energy must be conserved in order to have enough energy for the future.
b. In chemical processes, it is best to conserve as much energy as possible.
c. In any process, energy is neither created nor destroyed.
d. Energy cannot be transformed from one kind of energy into another kind.
____
3. Which of the following techniques cannot be used to separate homogeneous mixtures?
a. distillation
b. crystallization
c. chromatography
d. filtration
____
4. Raising the temperature of a gas in a rigid container will most likely change the
a. number of particles in the container.
c. pressure exerted by the gas in the container.
b. density of the gas in the container.
d. size of the gas particles in the container.
____
5. When a sample of a gas is kept at constant pressure, the volume and Kelvin temperature of the gas are
a. also unchanging.
c. inversely proportional.
b. directly proportional.
d. independent of one another.
____
6. In a mixture of gases, the total pressure of the mixture is equal to the
a. average of the pressures exerted by each gas.
c. sum of the pressures exerted by each gas.
b. pressure exerted by the most abundant gas.
d. pressure exerted by the least abundant gas.
____
7. What does the Sample 1 in the figure illustrate?
a. a gas
b. a solid
c. plasma
d. a liquid
____
8. In which of the samples shown in the figure is the density high?
a. sample 1 only
c. samples 1 & 3
b. sample 2 only
d. samples 2 & 3
____
9. A chemical change occurs when
a. dissolved minerals solidify to form a crystal.
b. ethanol is purified through distillation.
c. salt deposits form from evaporated sea water.
d. a leaf changes color.
____
10. According to the law of definite proportions, any two samples of KCl have
a. the same mass.
c. the same melting point.
b. slightly different molecular structures.
d. the same ratio of elements.
____
11. The nucleus of an atom has all of the following characteristics EXCEPT that it
a. is positively charged.
c. contains nearly all of the atom's mass.
b. is very dense.
d. contains nearly all of the atom's volume.
____
12. An atom is electrically neutral because
a. neutrons balance the protons and electrons.
b. nuclear forces stabilize the charges.
____
c. the numbers of protons and electrons are equal.
d. the numbers of protons and neutrons are equal.
13. The smallest unit of an element that can exist either alone or in combination with other such particles is a(n)
a. electron
b. proton
c. neutron
d. atom
____
14. As the mass number of the isotopes of an element increases, the number of protons
a. decreases.
c. remains the same
b. increases.
d. doubles each time mass # increases
____
15. In determining atomic mass units, the standard is the…
a. C-12 atom
b. C-14 atom
c. H-1 atom
____
d. O-16 atom
16. The average atomic mass of an element is the average of the atomic masses of its
a. naturally occurring isotopes.
c. nonradioactive isotopes.
b. two most abundant isotopes.
d. artificial isotopes