Chemistry 1-2E
Semester I Study Guide
Name_______________________
Hour_____
Chapter 1
1. Define the following terms.

Matter

Mass

Law of Conservation of Mass
2. Define and give 2 examples of the following:

Pure substance

Element

Compound

Mixture

Homogeneous

Heterogeneous
3. What are the 3 phases of matter?
4. What is a physical change? Give 2 examples.
5. What is a chemical change? Give 2 examples.
6. What is the difference between an endothermic and exothermic reaction?
7. What is a group or family? What do all members of a group have in common?
8. What are the names of the following groups?
Group 1
Group 2
Group 17
Group 18
9. How are group 18 elements different than all other elements?
10. Where are the following found on the periodic table?
Metals
Nonmetals
Metalloids
Transition metals
Inner transition metals
11. List properties of metals, nonmetals and metalloids.
12. Are most elements metals, nonmetals, or metalloids? What state do most elements exist
in at room temperature?
13. What is a period? How many are there on the periodic table?
14. What happens to atoms in chemical reactions?
Chapter 2
15. How are quantitative observations different than qualitative observations? Give 2
examples of each
16. How many significant figures are there in each of the following numbers?
a. 27
b. 2700
c. 2700.
d. 2700.0
e. 2.7 x 10
3
f. 0.0524
g. 0.05240
h. 5.24 x 10
-1
i. 5.240 x 10
-2
17. Give the result of each of the following calculations in the correct number of significant
figures.
a. 124.0 x 2.4 =
2
b. 1.240 x 10 x 2.4 =
c. 83 + 55.89 + 72 =
d. 123.4 + 0.06 + 100.0 =
e. 123.4 - 0.06 + 100.0 =
18. Complete the following chart for SI prefixes:
1J
= ________ mJ
1 kJ
= ________
1J
= ________ cJ
J
19. Convert the following measurements:
a) 1200 cm =________ m
b) 675 mL = ________L
c) 4.8 kg = ________ g
d) .004 g = ________ mg
e) 20 m = ________ mm
20. Distinguish between mass and weight.
21. What is the formula for density?
22. Define weight and explain how it changes as gravity changes.
23. Calculate the density of an object with a mass of 2.34 g and a volume of 4.68 mL.
24. A graduated cylinder has 20 ml (cm3) of water placed in it. An irregularly shaped rock is
then dropped in the graduated cylinder and the volume of the rock and water in the
cylinder now reads 30 ml (cm3). The mass of the rock dropped into the graduated
cylinder is 23 grams.
a. Find the volume of the rock dropped into the graduated cylinder.
b. Find the density of the rock dropped into the graduated cylinder.
25. Explain the difference between accuracy and precision and give an example of each.
26. List the SI base unit for mass, volume, length, time, amount of substance, and
temperature.
Chapter 3
27. What are the charges of the 3 subatomic particles?
28. What is an isotope? Give an example
29. Fill in the chart below for the following neutral atoms.
Element
Atomic #
Mass #
# of protons
# of electrons
# of neutrons
P-32
Ca-40
Br-81
I-125
Au-197
30. How many valence electrons do the following have?
K
Mg
Al
C
P
O
Br
Ar
31. How many valence electrons do most atoms need to become as stable as possible?
32. What part of the atom is involved in compound formation?
33. Why don’t noble gases react with other substances?
34. Define the following: ion, anion, and cation.
35. Summarize Rutherford’s conclusions from his gold foil experiment.
36. Calculate the average atomic mass of magnesium using the following data for three
magnesium isotopes.
Isotope
mass (u)
relative abundance
Mg-24
23.985
78.7%
Mg-25
24.986
10.13%
Mg-26
25.983
11.17%
Chapter 7
37. Write the formula for the following ionic compounds.
a. calcium iodide
__________
b. sodium chloride
__________
c. lead (II) sulfide
__________
d. magnesium fluoride __________
e. iron (III) carbonate __________
f. lithium nitrate __________
g. aluminum hydroxide __________
38. Name the following ionic compounds.
a.
Al2O3
_________________________
b.
K2O
_________________________
c.
FeCl2
_________________________
d.
Na2CO3 __________________________
e.
Mg(NO3)2 _________________________
f.
Cu2SO4 ___________________________
h. Mg3(PO4)2 _________________________
45. Indicate the prefixes used for covalent names for the following:
1 _________
6 _________
2 _________
7 _________
3 _________
8 _________
4 _________
9 _________
5 _________
10 _________
46. Write the formula for the following binary molecular (covalent) compounds.
a. phosphorus trichloride __________
b. carbon dioxide __________
c. phosphorus trichloride __________
d. dinitrogen pentoxide __________
e. silicon dioxide __________
47. Name the following binary molecular substances.
a. SO2 ________________________
b. CO _________________________
c. SO3 ________________________
d. CF4 ________________________
e. IF7 _________________________
48. Write the formula and charge of the following polyatomic ions.
a. ammonium __________
b. nitrite ____________
c. nitrate ____________
d. hydroxide ___________
e. carbonate ____________
f. sulfate ____________
g. phosphate __________
49. Name the following acids
a. HCl __________________
b. HF
__________________
c. H2SO4 ___________________
d. H3PO3 ___________________
50. Write the formula for the following acids.
a. Hydrosulfuric acid
___________
b. Hydrobromic acid _____________
c. Carbonic acid
d. Nitrous acid
______________
_______________
51. What is Avogadro’s constant? What does it mean?
52. Define the term molar mass.
53. Determine the molar mass of the following:
a. CO2
b. H2SO3
c. Mg(NO3)2
54. How many atoms are in 4.67g of zinc?
55. What is the number of moles represented by 25.6 g of CaCl2?
56. What is the mass of 3.45 moles of Al2(SO4)3?
57. What is the mass of 14.2 L of carbon dioxide gas?
58. What is the percent composition of each element in Al2(CO3)3?
59. What’s the empirical formula of a molecule containing 65.5% carbon, 5.5% hydrogen, and
29.0% oxygen?
60. If the molar mass of the compound in problem 61 is 110 grams/mole, what’s the molecular
formula?
61. In an experiment, rubidium chloride ___ hydrate was heated to remove water. The
following data was obtained:
1. Mass of empty crucible
60.286 g
2. Mass of crucible & contents before heating
79.376 g
3. Mass of crucible & contents after heating
70.366 g
4. Calculate:
a) The formula of the hydrate:
____________________________
b) The name of the hydrate:
____________________________
c) The percent water of the compound _________________________
Chapter 8
59. List the five types of signs that indicate a chemical reaction has taken place.
1.
2.
3.
4.
5.
60. Describe what is meant by a balanced chemical equation in your own words.
61. Explain when subscripts and coefficients are used when writing a balanced chemical
equation.
62. Indicate the symbols used in a chemical equation to describe the following:
a. Solid –
b. Liquid –
c. Gas –
d. Aqueous –
63. Given a reaction, be able to balance it and tell whether it is synthesis, decomposition,
single replacement, double replacement or combustion. Balance the examples below and
identify the type of chemical reaction.
a.
Na + Br2  NaBr
b. Al2O3  Al + O2
c.
Mg +
HCl  MgCl2 + H2
d. NaBr + F2  NaF + Br2
e
Al(NO3)3 + H2SO4 
f
C2H6 + O2  CO2 + H2O
Al2(SO4)3 + HNO3
64. Indicate the symbols used in a chemical equation to describe the following:
a. Solid –
b. Liquid –
c. Gas –
d. Aqueous –
65. Predict what would happen to a burning/smoldering wooden splint if inserted into a test
tube containing the following gases:
CO2
O2
H2
Chapter 21
65. Define the following:
a. Radiation –
b. Half-life –
c. Fission –
d. Fusion –
66. How long would an 88 gram sample of Pb-210 (half-life = 3.5 days) have to undergo
radioactive decay before only 6.5 grams of Pb-210 would be remain?
67. Write out the decay reaction to describe the following:
a. Polonium-214 undergoing α decay
b. Astatine-218 undergoing β decay
c. The alpha emission of a radioisotope that would produce Pb-209 as a daughter
product.
d. The β decay of a radioisotope that would produce Fr-230 as a daughter product.