Midterm Sample Problems
Chemistry 300
Make sure that you can solve problems like the ones shown here. Please remember, this should
NOT be your only way of studying. Make sure to study all of your notes, homework, quizzes,
and tests. It is important to make your mind active when you study… simply “reviewing notes”
and “looking over problems” DOES NOT WORK. Quiz yourself. Quiz each other. Do practice
problems. Redo old problems under test taking conditions (i.e., with no notes) and analyze how
you did and what you need to work on. Also… this page does not leave space for your work…
so plan on doing these on separate paper or in your notebook.
1. What is the difference between all of the following: pure substance, mixture, element,
compound, homogeneous mixture, heterogeneous mixture, solution, suspension? How
are these terms related to each other?
2. Which words from #1 could be used to classify the following: carbon, aqueous
copper(II)chloride, muddy water, sulfur dioxide, blood, potassium permanganate crystals,
chicken soup, and titanium?
3. What are some ways to separate mixtures? How about pure substances?
4. What is the difference between physical and chemical properties of matter? List a few
examples for each type.
5. What is the difference between a physical change and a chemical change? List a few
examples of each type of change.
6. List some signs to indicate that a chemical change or chemical reaction has taken place.
7. What is density? How can you calculate it? How do you know what units to use?
8. A solid object has a mass of 35.0 g and a volume of 15.00 mL. Find its density.
9. A graduated cylinder is filled with 75.3 mL of water. When a 12.579 g piece of lead is
added, the water level rises to 77.4 mL. What is the volume of the lead? What is the
volume of water displaced? What is the density of the lead?
10. SKIP
11. How do you calculate percent error?
12. What do the metric prefixes micro-, milli-, centi-, kilo-, mean? How would you use them
in conversion factors with a base unit (i.e., meters)?
13. Convert 10.0 km into cm.
14. In an experiment, a student obtained a density for water of 1.05 g/mL. The actual density
of water is 1.00 g/mL. What was the student’s percent error?
15. Describe the contributions of Dalton, Thomson, Rutherford, and Bohr to our “picture” of
the structure of the atom.
16. What is the key difference between Bohr’s idea of atomic structure from the quantum
mechanical model?
17. What does atomic number represent? How do you recognize it on any Periodic Table?
Give an example of an element and its atomic number.
18. What does atomic mass represent? How do you recognize it on any Periodic Table?
Give an example of an element and its atomic mass.
19. How is the atomic mass of an element determined?
20. What does mass number represent? How is it determined? Give an example of an
element and its mass number.
21. What is an isotope? How many protons, neutrons, and electrons are in an isotope of
carbon with a mass number of 14? How about in potassium-40?
22. What is an ion? Why do ions form?
23. What is the difference between a cation and anion? What kind of elements tend to form
cations? How do they do it? What kind of elements tend to form anions? How do they
do it?
24. How many protons, neutrons, and electrons are in an ion of rubidium? How about in an
ion of bromine?
25. What does it mean if something is isoelectronic? What element are the ions in #28
isoelectronic with?
26. What are some of the benefits and dangers of nuclear radiation?
27. What does the concept of half life mean?
28. If isotope Z has a half life of 20 days… if we start out with 100 g of isotope Z, how much
would be left after 80 days? What fraction of the original amount would be remaining at
that point?
29. If we have the same isotope as in #37… this time, we have 15 g remaining after 120 days.
What was the original amount?
30. Describe what happens with electrons to produce the atomic emissions spectrum of an
element.
31. Write out the orbital diagram, full electron configuration, noble gas abbreviated
configuration, number of valence electrons, for the following species: an atom of sodium,
an atom of iron, an atom of iodine, an ion of sodium, and an ion of iodine.
32. In #31, what is the total number of unpaired electrons for each species.
33. In #31, which element is the ion of sodium isoelectronic with? How about the ion of
iodine?
34. What does it mean if something is periodic or has periodicity?
35. What does Periodic Law state?
36. Describe and explain the shielding effect and nuclear charge and their trends on the
Periodic Table.
37. Describe and explain the trends for atomic radius, ionization energy, and
electronegativity, across the periods and down the groups of the Periodic Table.
38. Rank the following elements in order of increasing atomic radius, then decreasing
ionization energy, then increasing electronegativity (i.e., do 3 separate rankings): Ca, K,
Kr, Se, Br.
39. Rank the following elements in order of increasing atomic radius, then decreasing
ionization energy, then increasing electronegativity (i.e., do 3 separate rankings): Mg, Sr,
Be, Ca.
40. Name a couple of properties of metals, nonmetals, and metalloids and give a couple
examples of each.
41. Redo your entire formulas & naming quizzes and test on separate paper with no reference
materials other than a Periodic Table with symbols.
42. Describe the 5 types of reactions. What does each one “look like” in general?
43. Describe how to predict products of reactions once you know what type of reaction it is.
Give an example of each.
44. What is a mole? What is Avogadro’s number? Why is it used?
45. How do you determine the molar mass of a substance? Calculate the molar mass of
calcium, hydrogen, carbon dioxide, and copper(II) sulfate.
46. How do you determine the percent composition of a substance? Determine the percent
composition of sulfur trioxide and calcium phosphate. Determine the percent hydration
in copper(II) sulfate pentahydrate.
47. What is the difference between an empirical and molecular formula? Give an example of
a substance whose empirical and molecular formulas are the same and an example of a
substance whose empirical and molecular formulas are different.
48. Find the empirical and molecular formulas for a compound with the following percent
composition: 54.5 % carbon, 9.1% hydrogen, and 36.4% oxygen. Its molecular mass is
88 g/mol.
49. How do you convert between moles, mass, and particles (which can be molecules or
atoms)? What are the conversion factors for your dimensional analysis?
50. Determine the mass of 5.25 moles of calcium chloride.
51. Determine the number of moles in 25.97 g of water.
52. Determine the number of formula units of calcium chloride in #50.
53. Determine the number of water molecules in #51.
54. Determine the total number of atoms in #53.
55. Write down a general strategy for solving stoichiometry problems. What are the key
things that you need to remember?
56. Write a BALANCED equation for the reaction of hydrogen and nitrogen gases to
produce ammonia (otherwise known as nitrogen trihydride). The remaining questions
will ALL refer back to this equation.
57. What mass of ammonia can be produced if 12.46 g of hydrogen reacts with excess
hydrogen?
58. How many molecules of hydrogen are required to produce 5.00 g of ammonia (if
unlimited nitrogen is available)?
59. What mass of nitrogen is required to react with 12.46 g of hydrogen?