Download Semester 1 Final Exam Study Guide

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
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
i. 5.240 x 10
-1
-2
17. Give the result of each of the following calculations in the correct number of significant figures.
2
a. 124.0 x 2.4 =
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:
1g
= ________ mg
1g
= ________ cg
1 kg
= ________
g
19. Convert the following measurements:
a) 1200 cm =________ m
675 mL = ________L
b) 4.8 kg = ________ g
0.004 g = ________ mg
c) 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 4
37. What is a wavelength?
38. Which color on the electromagnetic spectrum have the longest and shortest wavelength?
39. How does an atom move from the ground state to the excited state?
40. Fill in the following chart:
Main
energy
level
Sublevels
in main
energy
levels
Number of
Number of
Number of
Number of
orbitals/sublevel orbitals/main electrons/sublevel electrons/main
level
level
41. Write the electron configuration for:
a. Phosphorus
b. Arsenic
42. Write the noble gas shortcut electron configuration for calcium.
Chapter 5
43. Whose idea was it to arrange the elements in the periodic table according to their chemical and physical
properties?
44. Why did Mendeleev leave empty spaces in his periodic table?
45. What scientists work led to a periodic table based on increasing atomic number?
46. Write the name of the following groups of elements:
Group 1 ______________________
Group 2 _______________________
Group 17 ______________________
Group 18 ______________________
47.
Write the name of the elements with atomic numbers from:
58 to 71 ___________________
90 to 103 ___________________
48.
Define the following terms as they relate to the periodic table:
Group
Period
50.
Indicate the number of valence (outermost) the following elements have:
K
C
Cl
Mg
P
Ar
Al
S
52. What is the main characteristic of noble gases?
53. Which are more reactive, alkali metals or alkaline-earth metals? Why?
54. Define the term electronegativity. Which elements have a high and low electronegativity? What is the
general trend on the periodic table?
55. Define the term ionization energy. Which elements have a high and low ionization energy? What is the
general trend on the periodic table?
56.
How does atomic radius change as you move across a row and down a group on the periodic table?
57. Define the terms anion and cation. Give an example of each.
Chapter 6
58. Which electrons are involved in the formation of a chemical bond
59. What is the difference between a polar covalent and a nonpolar covalent bond?
60. What is the difference between an ionic and a covalent bond?
61.
What subatomic particles are shared in a covalent bond?
62.
What is a molecule?
63.
What does the octet rule state?
64. What do you need to know before you draw a Lewis structure (dot diagram).
65. After you finish drawing a Lewis structure, what should you double check?
66. Draw Lewis structures and use VSEPR theory to predict the shapes (with bond angles) for the following:
NH3
OH-
CH4
H2O
67. What is a hybrid orbital? What type of hybrid orbitals would CCl4 have?
68. What term is used to describe the movement of electrons involved in a metallic bond?
69. Define the following terms:
Ductile
Malleable
Chapter 7
70. 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 __________
lithium nitrate __________
g. aluminum hydroxide __________
71. Name the following ionic compounds.
a.
Al2O3
_________________________
b.
K2O
_________________________
c.
FeCl2
_________________________
d.
Na2CO3 __________________________
e.
Mg(NO3)2 _________________________
f.
Cu2SO4 ___________________________
h. Mg3(PO4)2 _________________________
72. Indicate the prefixes used for covalent names for the following:
1 _________
6 _________
2 _________
7 _________
3 _________
8 _________
4 _________
9 _________
5 _________
10 _________
73. Write the formula for the following binary molecular (covalent) compounds.
a. phosphorus trichloride __________
carbon dioxide __________
c. phosphorus tetraoxide __________
dinitrogen pentoxide __________
e. silicon dioxide __________
74. Name the following binary molecular substances.
a. SO2 ________________________
CO _________________________
b. SO3 ________________________
CF4 ________________________
c. IF7 _________________________
75. Write the formula and charge of the following polyatomic ions.
a. ammonium __________
nitrite ____________
b. nitrate ____________
hydroxide ___________
c. carbonate ____________
sulfate ____________
d. phosphate __________
76. Name the following acids
a. HCl __________________
HF
__________________
b. H2SO4 ___________________
H3PO3 ___________________
77. Write the formula for the following acids.
a. Hydrosulfuric acid
b. Carbonic acid
___________ Hydrobromic acid _____________
______________ Nitrous acid
_______________
78. What is Avogadro’s constant? What does it mean?
79. Define the term molar mass.
80. Determine the molar mass of the following:
a. CO2
b. H2SO3
c. Mg(NO3)2
81. How many atoms are in 4.67g of zinc?
82. What is the number of moles represented by 25.6 g of CaCl2?
83. What is the mass of 3.45 moles of Al2(SO4)3?
84. What is the mass of 14.2 L of carbon dioxide gas?
85. What is the percent composition of each element in Al2(CO3)3?
86. What’s the empirical formula of a molecule containing 65.5% carbon, 5.5% hydrogen, and 29.0% oxygen?
87. If the molar mass of the compound in problem 61 is 110 grams/mole, what’s the molecular formula?
88. 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 _________________________