Ch. 1: Matter and Change
1. Explain the difference between the terms: atom, element, compound, and
molecule. Give an example of each.
2. Explain the difference between an intensive and extensive property.
3. What is the difference between a physical and chemical property?
4. Explain the difference between a pure substance and a mixture.
5. Classify each of the following as homogeneous or heterogeneous mixtures,
compound, or element:
a. silver alloy
b. nitrogen dioxide
c. sugar dissolved in water
d. lead
e. mud
6. Classify each of the following as physical or chemical change:
a. boiling water
b. mold growing on bread
d. melting iron
e. decomposing hydrogen peroxide
c. pressurization of gas
7. Explain the difference between a physical and chemical change.
8. Describe the differences between gas, liquid, and solid states in terms of particle
movement and whether they have constant volume and shape.
9. List the phase changes that are endothermic and the ones that are exothermic.
10. What is the difference between decanting and filtration?
11. What is the difference between a group and a period?
12. List some properties of nonmetals, noble gases, metals, and metalloids.
13. Do elements of the same group or period have more similar characteristics?
14. Why does gold have a symbol of Au instead of Go or Gd?
15. Why are the lanthanide and actinide series separate from the rest of the table?
16. Identify the period number, group number, symbol, and category (metal,
nonmetal, metalloid, noble gas) for each of the following elements.
Ch. 2: Measurements and Calculations
1. What is the difference between qualitative and quantitative data?
2. Complete the following conversions, showing every step.
a. 958mgkg
b. 0.57μLcL
c. 10.5cm3L
d. 0.079cm2km2
e. 8.79 Mm3 m3
f.148 kg/mL g/L
3. Explain the difference between weight and mass.
4. Complete the chart below, filling in SI units, quantities, and abbreviations for
each.
5. Does a liquid have a higher or lower density than a solid normally? What is an
exception?
6. A substance has a density of 1.58 g/cm3. Would is float in water? What would be
the mass of a sample whose volume is 71.0 cm3?
7. How many significant figures does each of the numbers have?
a. 8.430
b. 7900
c. 0.00872
d. 60450
e. 0.030050
8. What is the difference between accuracy and precision?
9. What is the difference between directly and inversely proportional? Draw graphs
of each.
10. Perform the following calculations.
a. 561.0m+2.73m+0.00687m
b. 2.3x103 mL + 2.1x102 L + 8.7 L
11. Round the following numbers to 2 significant figures.
a. 0.0087950
c. 8.7536
b. 2003
d. 875.5
12. Put the following numbers in scientific notation:
a. 1200
c. 250,000
b. 0.00853
d. 0.000092
13. Explain how the uses of percent error and difference are different.
14. Convert 28 C to Fahrenheit and Kelvin.
Ch. 3: Atoms: The Building Blocks of Matter
1. Who are Democritus and Aristotle?
2. What is the law of conservation of mass?
3. What are the two parts to Dalton’s Atomic Theory that have been found to be
incorrect? Explain why they are incorrect.
4. Describe what the law of definite composition and law of multiple proportions
say. What kind of compounds do they each apply to?
5. Describe the two main parts to an atom. Which has a greater mass and a greater
volume? Draw a diagram to show where they are.
6. What are subatomic particles? What are the three types? How do their amounts
relate?
7. What holds the nucleus together?
8. Describe the cathode-ray tube experiment. What was it used to discover?
9. What did Thomson and Milliken do?
10. Describe the gold foil experiment. What was this used to discover?
11. Explain the difference between mass number and atomic number.
12. How is the average atomic mass calculated for an atom?
13. What is the mass in grams of 8.64 mol of oxygen atoms?
14. A sample of sulfur contains exactly 8.4 x 1016 atoms. How many moles is that?
15. A sodium chunk contains 2.8x1010 atoms. How many grams of sodium is that?
Ch. 4: Arrangement of Electrons in Atoms
1. Describe the photoelectric effect. What did it help prove?
2. Draw the ground state configuration for the electrons in nitrogen. What would an
example of an excited state configuration be?
3. What type of movement of an electron requires the absorption and the emission
of energy?
4. Describe the relationship between energy, wavelength, and frequency.
5. The wavelength of light is 780 nm.
a. Calculate the frequency in waves per second.
b. Calculate the energy of that light.
6. What is a line-emission spectrum? Why is it an identifying feature of a type of
atom?
7. What was wrong with the Bohr model?
8. How did Einstein and Planck contribute to quantum mechanics?
9. Describe what each of the four quantum numbers tells you about an electron.
10. Draw the shapes of the s and p orbitals.
11. Determine the four quantum numbers for the last electron put in each of the
following atoms.
a. Al
b. Mn
c. Sr
d. Xe
12. What does each of the following tell us about the arrangement of electrons:
a. Heisenberg Uncertainty Principle
b. Hund’s Rules
c. Aufbau Principle
d. Pauli’s Exclusion Principle
13. Describe the relationship between electrons, energy levels, sublevels and
orbitals.
14. Write electron configuration in orbital notation, electron configuration notation,
and noble gas notation for the following elements:
a. Si
b. Be
c. Co
d. Te
Ch. 5: Periodic Law
1. What did Mendeleev and Moseley each do? Who is considered more important?
Why?
2. What categories of elements were discovered last? Why?
3. Describe the arrangement of our modern periodic table. How does it differ from
the early periodic tables?
4. Differentiate between alkali metals, alkaline earth metals, and transition metals.
5. Differentiate between noble gases and halogens.
6. How are hydrogen and helium different from the rest of the elements in their
individual groups? Why are they placed where they are in the periodic table?
7. The term “main group elements” or “representative elements” refers to
_________________________________________________.
8. List some characteristics of metalloids. What else are these elements called?
9. Moving from top to bottom down a group:
a. Electronegativity becomes______________________.
b. Electron affinity becomes _______________________________.
c. Ionization energy becomes ______________________________.
d. Atomic radii become ____________________________.
10. Name the element described below:
a. Element in Period2 with smallest electronegativity
b. Alkali metal with most negative electron affinity.
c. Halogen with lowest ionization energy
d. Element in Period 5 with largest atomic radius
11. How does the second ionization energy of an atom compare to its first? Why?
12. Which atom- Li, Na, or K - has the largest atomic radius? Explain why.
(explain why the trend occurs, do not just state it)
13. Which atom- B, C, or O - has the largest ionization energy? Explain why.
(explain why the trend occurs, do not just state it)
Ch. 6: Chemical Bonding
1. Why do atoms form chemical bonds?
2. Explain the basic difference between a covalent and an ionic bond.
3. What is the difference between a polar and a nonpolar bond? Which occurs
between two atoms with large differences in electronegativity?
4. Why is electronegativity so important for determining bond type?
5. How does the difference in electronegavitites point to bond type?
(see pg 151 for values)
Atoms
Difference
Classification
Mg and O
C and F
B and H
6. What determines or influences the distance two bonded atoms stay from each
other?
7. What is a diatomic molecule? List the elements that naturally exist as a diatomic
molecule.
8. Describe the octet rule.
9. Which elements do not need 8 electrons?
10. Which elements can have more than 8 electrons? Why?
11. What is the purpose in drawing resonance structures for a molecule?
12. Draw the Dot Notation for the following atoms.
N
C
F
13.
OF2
Draw the Lewis Structures for the following molecules.
CF4
PF5
14. Explain the relationship between energy, bond length, and bond strength?
15. How is a covalent network different from a normal covalently bonded molecule?
16. Draw Lewis Structures for the following polyatomic ions.
NH4+
PO4-3
17. Explain the VSEPR theory. What does it help us to do?
18. Determine the geometry of the following molecules.
19. Describe the structure that ionic compounds form.
20. What is the difference between bond energy and lattice energy? What does it tell
you about that compound?
21. Explain the “sea of electron” model of metallic bonding. What makes metals
good conductors? Explain how metals are malleable.
Ch. 7: Chemical Formulas and Compounds
1. What is the difference between an ionic and a molecular compound?
2. How can you identify an acid by its formula?
3. Identify the following compounds as acids (binary or ternary), ionic compounds
or molecular compounds.
a. MgBr2
b. XeCl4
c. H3PO4
d. HI
e. FeSO4
f. N2O5
4. Name the following compounds. If there is more than one way to name it, show
both.
a. MgBr2
b. XeCl4
c. H3PO4
d. HI
e. FeSO4
f. N2O5
5. Determine the oxidation state of each element in the following compounds
a. MgBr2
b. C2H4
c. H3PO4
d. HI
e. CaSO4
f. N2O5
6. Determine the number of atoms total in one molecule or unit in the following
compounds
a. Fe3(PO4)2
b. Cr(NO2)3
c. H3PO4
d. Ca(ClO2)2
e. FeSO4
f. (NH4)3N
7. What is the difference between an atom and an ion?
8. When do you need a Roman numeral between the two parts of a compound’s
name? What does it mean or stand for? What do you call the naming system?
9. Which type of compound always has a simplified formula? 11. Explain what the
difference between a cation and an anion is.
10. What type of naming uses numerical prefixes?
11. When do you “cross over”? What does it do?
12. What is the difference between the molar mass and the formulas mass? How are
they calculated?
13. What is the difference between an empirical and a molecular formula?
14. The empirical formula for a compound is AB2. Give 3 possible molecular
formulas for this compound.
15. What do you use to convert between moles and grams? What about moles and
molecules?
16. There are 4.3 x 1011 molecules of CO in a sample. 2
a.
Find the number of moles in the sample.
b.
Find the number of grams in the sample.
c.
Find the number of oxygen atoms in the sample.
17. A sample of PF5 contains76g.
a. Find the moles of the sample.
b. Find the number of molecules in the sample.
18. Find the percent composition of Fe(NO3)2.
19. Find the percent of chlorine in Co(ClO2)2
20. The percent composition of a compound is: 24.3% Fe, 33.9% Cr, and 41.8% O.
Find the empirical formula.
21. The mass composition of a sample for a compound is 9.65 g of Pb and 0.99 g of
O. Determine the empirical formula.
22. The empirical formula for a compound is C2H2O. If the actual molar mass is 254
g/mol, what is the molecular formula?