1. Describe the arrangement of the periodic table in terms of chemical reactivity and electron
configurations.
2. Determine the type of bonding involved for the compounds formed between:
a. Lithium and nitrogen
b. Iodine and fluorine
c. Manganese and sulfur
3. What is meant by the term electrolyte? Describe the difference between a strong and a
weak electrolyte.
4. Explain:
a. Why lithium has a smaller atomic radius than potassium.
b. Why oxygen has a lower first ionization energy than fluorine.
5. Describe the difference between metals, non-metals, and metalloids.
6. Write names for the following chemical compounds:
a. Rb2O
b. Ti2O
c. Ag2S
d. Pb(NO3)2 e. Sn3(PO4)2 f.
(NH4)HSO3 g. SiCl4
h. PCl3
i. P4O10
j. BaCO3
k. IF7
l. CuSO 5H O 42  m. HClO3
n. KClO4
o. LiRbSO4
7. Write correct chemical formulae for the following compounds:
a. Sodium oxide
i. Mercury (I) chloride
b. Potassium cyanide
j. Sulfur hexafluoride
c. Copper (II) nitrate
k. Barium hydrogen phosphate
d. Zinc sulfide
l. Hypobromous acid
e. Magnesium fluoride f.
m. Silicon dioxide
Chromium (III) carbonate g.
n. Hydroiodic acid
Chlorine monoxide h.
o. Ammonium phosphite
Strontium carbonate
p. Lead (IV) phosphide
hexahydrate
8. Draw structures for the following organic compounds:
a. 2,2,3 – trimethylhexane
b. 2,3 – dichloro – 3 – ethyl – 2 – methyl – 1 – pentanol
c. 2,5 – dimethyl – 3 – heptene
d. 2 – ethyl – 1,3 – dimethylcyclohexane
e. 1,2 – dichlorobenzene
f. 2 – bromomethylbenzene
g. 4 – chloro – 2 – methyl – 3 – hexanone
h. 4 – chlorobutanal
i. Ethyl butanoate
j. 2,2,3,3 – tetrafluoropropanoic acid
9. What is a mole? Why is it used in chemistry?
10. What is the mass of 500. atoms of Fe?
11. How many moles of molecules are there in the following:
a. 24.5 g of ammonium nitrate.
b. 142.3 g of lead (II) oxide.
c. 1.21 kg of cryolite (Na3AlF6).
12. How many molecules are present in each of the samples in question 11? How many
atoms?
13. Determine the percent composition by mass present in acrylonitrile (C3H3N).
14. Vitamin B12 contains 4.34% cobalt by mass. Assuming that each molecule of Vitamin
B12 contains one atom of cobalt, what is the molar mass of Vitamin B12?
15. A compound contains only carbon, hydrogen, and oxygen. Combustion of 10.68 mg of
the compound yields 16.01 mg of carbon dioxide and 4.37 mg of water. The molar mass
of the compound is 176.1 g/mol. What are the empirical and molecular formulae for the
compound?
16. Only one isotope of this element is found in nature. A single atom of this isotope has a
mass of 23 9.123 10− × g. What is the element?
17. How many molecules are there in 14.3 L of H2 gas?
18. Solutions are made up as follows. What are the concentrations in units of Molarity? a.
43.2 g of calcium nitride dissolved in enough water to make 500. mL of solution. b. 7.53 g
of lithium fluoride in enough water to make 125 mL of solution. c. 214 g of cesium iodide
in enough water to make 600. mL of solution.
19. Write balanced molecular and net ionic equations for the following situations: a.
Magnesium metal is dropped into aqueous hydrochloric acid, producing aqueous
magnesium chloride and hydrogen gas.
b. Solid potassium oxide is mixed with water to form aqueous potassium hydroxide.
c. Aqueous mercury (II) nitrate is mixed with aqueous potassium iodide to produce a
precipitate of mercury (II) iodide.
d. Aqueous solutions of potassium bromide and chlorine are reacted to form aqueous
solutions of bromine and potassium chloride.
20. Classify the following reactions, then predict the products of the reaction, and balance the
reaction equation. Note that not all of the reactions will actually take place. a. Δ HgO( )s
⎯⎯→
b. CuCO ( ) + HNO ( ) 33 s aq ⎯⎯→
c. BaCl (aq) + Na SO (aq) 2 24 ⎯⎯→
d. CaO (s) + SO (g) 2 ⎯⎯→
3
e. () 2 Zn (s) + Cu NO (aq) ⎯⎯→
f. 2 I (aq) + CsBr (aq) ⎯⎯→
g. Al O (s) + H O (l) 23 2 ⎯⎯→
h. Cu (s) + HCl(aq) ⎯⎯→
43 2
i. ()
2 NH CO (aq) + CaCl (aq) ⎯⎯→
j. 4 ()2 HClO (aq) + Mg OH (s) ⎯⎯→
21. Explain the significance of the following concepts:
a. Hund’s Rule
b. Aufbau Principle
c. Pauli Exclusion Principle
22. Write correct electronic configurations for the following elements:
a. Silicon
f. Selenium g.
b. Gallium c.
Aluminum h.
Arsenic d.
Sodium i.
Chromium e.
Krypton j.
Silver
Erbium
23. Write the expected ground state electron configurations for the following: a. The
element with one unpaired 5p electron that would form a covalent compound with
fluorine.
b. The (at present) undiscovered alkaline-earth metal after radium.
c. The noble gas that has electrons occupying 4f orbitals.
d. Element 116.
e. The halogen with electrons in 6p atomic orbitals.
24. Element 117 is currently undiscovered, and will be given a placeholder name of
ununseptium when initially discovered. Predict some of the properties of Uus by
answering the following questions:
a. What will be its electron configuration?
b. What element will it most resemble chemically?
c. What will be the formula of the binary compounds that Uus forms with sodium,
magnesium, nitrogen, and oxygen?
d. What oxyanions would you expect Uus to form?
25. The combustion of propane releases 2061 kJ/mol.
a. Write the balanced chemical reaction
b. How much energy would be released by the combustion of 638 g of propane? c.
What volume of carbon dioxide would be collected at STP by the combustion of
638 g of propane?
26. Mr. Ward analyzes a sulfuric acid sample to determine its concentration. A 10.00 mL
sample of sulfuric acid is titrated with 26.42 mL of a 3.121 M sodium hydroxide solution.
What is the concentration of the sulfuric acid sample?
H SO (aq) + 2 NaOH (aq) Na SO (aq) + 2 H O (l) 24 24 2 ⎯⎯→
27. How many grams of carbon dioxide can be produced by the reaction of 10.0 g of iron (III)
oxide with excess carbon monoxide? Note that this reaction is NOT balanced.
Fe O (s) + CO (g) Fe O (s) + CO 23 34 2 ⎯⎯→
28. Azobenzene (C12H10N2) is an important intermediate in the production of dyes. It is
prepared from the reaction between nitrobenzene (C6H5NO2) and triethylene glycol
(C6H14O6).
6 5 2 6 14 6 12 10 2 6 12 4 2
2C H NO + 4C H O C H N + 4C H O + 4 H O ⎯⎯→
a. What is the theoretical yield of azobenzene if 126 g of nitrobenzene is reacted
with 318 g of triethylene glycol?
b. What was the excess reagent? How much excess reagent remained unreacted?
29. Mr. Hooper observed two Middle School students pouring salt into the sugar bowl. He
was curious to determine how much sodium chloride was in the sugar, so he dissolved a
2.896 g sample of the sugar/salt mixture and precipitated the chlorine with an excess of
silver nitrate. If 1.226 g of silver chloride was isolated, calculate the percentage of
sodium chloride in the sugar sample.
AgNO + NaCl AgCl + NaNO 33 ⎯⎯→
30. Explain the following gas laws in the context of the postulates of the Kinetic Molecular
Theory:
a. Boyle’s Law
b. Charles’ Law
c. Avogadro’s Law
31. A gas occupies a volume of 725 mL at a pressure of 0.970 atm. It is allowed to expand at
a constant temperature until its pressure reaches 0.541 atm. What is its final volume?
32. If 6.9 moles of CO gas are present in a container with a volume of 30.4 L, what is the
pressure of the gas (in atm) if the temperature is 62 °C?
33. At 741 torr and 44 °C, 7.10 g of a gas occupies a volume of 5.40 L. What is the molar
mass of the gas?
34. A solution is made by dissolving 43.5 g of CsOH in 400.0 mL of water. Determine:
a. The molarity of the solution (assume the final volume is 400.0 mL). b. The mass
percentage of CsOH (density of water is 1.0 g/mL).
c. The mole fraction of CsOH.
35. Mr. Hooper prepares a solution by dissolving 0.8214 g of KMnO4 in a 500.0 mL
volumetric flask. He then takes a 2.00 mL sample of that solution and dilutes it to a
volume of 1000.0 mL. Mr. Ward then takes a 10.00 mL sample of the second solution
and dilutes it to 250.0 mL. What is the final concentration (in molarity) of the solution?
What mass of potassium permanganate would be required to directly prepare a solution
with the same concentration as the final solution?