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. PROVE you can do it. 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?
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11. A solid object has a mass of 35.0 g and a volume of 15.00 mL. Find its density.
12. 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?
13. How do you calculate percent error?
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15. What do the metric prefixes micro-, milli-, centi-, kilo-, and mega- mean? How
would you use them in conversion factors with a base unit (i.e., meters)?
16. Convert 10.0 km into cm.
17. 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?
18. Describe the contributions of Dalton, Thomson, Rutherford, and Bohr to our
“picture” of the structure of the atom.
19. What is the key difference between Bohr’s idea of atomic structure from the
quantum mechanical model?
20. Compare the mass, charge and location of protons, neutrons, and electrons.
21. What does atomic number represent? How do you recognize it on any Periodic
Table? Give an example of an element and its atomic number.
22. What does atomic mass represent? How do you recognize it on any Periodic
Table? Give an example of an element and its atomic mass.
23. How is the atomic mass of an element determined?
24. What does mass number represent? How is it determined? Give an example of
an element and its mass number.
25. 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?
26. What is an ion? Why do ions form?
27. 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?
28. How many protons, neutrons, and electrons are in an ion of rubidium? How about
in an ion of bromine?
29. What does it mean if something is isoelectronic? What element are the ions in
#28 isoelectronic with?
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37. N/A
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39. Describe what happens with electrons to produce the atomic emissions spectrum
of an element.
40. Write out the orbital diagram, full electron configuration, noble gas abbreviated
configuration, number of valence electrons, and Lewis structure 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.
41. In #40, what is the total number of unpaired electrons for each species.
42. In #40, which element is the ion of sodium isoelectronic with? How about the
ion of iodine?
43. What does it mean if something is periodic or has periodicity?
44. What does Periodic Law state?
45. Describe and explain the shielding effect and nuclear charge and their trends on
the Periodic Table.
46. Describe and explain the trends for atomic radius, ionization energy, and
electronegativity, across the periods and down the groups of the Periodic Table.
47. 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.
48. 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.
49. Name a couple of properties of metals, nonmetals, and metalloids and give a
couple examples of each.
50. Redo your entire formulas & naming quizzes and test on separate paper with no
reference materials other than a Periodic Table with symbols.
51. Describe the 5 types of reactions.
52. What does each one “look like” in general (use A and B, etc)?
53. Describe how to predict products of reactions once you know what type of
reaction it is. Give an example of each.
54. What is a mole? What is Avogadro’s number? Why is it used?
55. How do you determine the molar mass (i.e. gram formula mass) of a substance?
Calculate the molar mass of calcium, hydrogen, carbon dioxide, and copper(II)
sulfate.
56. How do you determine the percent composition of a substance? Determine the
percent composition of sulfur trioxide (SO3) and calcium phosphate. Determine
the percent hydration (water) in copper(II) sulfate pentahydrate.
57. 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.
58. 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.
59. How do you convert between moles, mass, and particles (which can be molecules
or atoms)? What are the conversion factors for your dimensional analysis?
60. Determine the mass of 5.25 moles of calcium chloride.
61. Determine the number of moles in 25.97 g of water.
62. Determine the number of formula units of calcium chloride in #50.
63. Determine the number of water molecules in #51.
64. Determine the total number of atoms in #53.
65. Determine the number of ions in #62
66. Write down a general strategy for solving stoichiometry problems. What are the
key things that you need to remember?
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68. Write a BALANCED equation for the reaction of hydrogen and nitrogen gases to
produce ammonia (NH3). The remaining questions will ALL refer back to this
equation.
69. What mass of ammonia can be produced if 12.46 g of hydrogen reacts with excess
hydrogen?
70. How many molecules of hydrogen are required to produce 5.00 g of ammonia (if
unlimited nitrogen is available)?
71. What mass of nitrogen is required to react with 12.46 g of hydrogen?
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