Download Chemistry Standards Review

CHEMISTRY CST REVIEW
Atomic and Molecular Structure
1. What is atomic number?
2. How does increasing the atomic number affect the identity of an atom?
3. What is atomic mass? What two subatomic particles make up the majority of an atom’s mass?
4. Describe the gold foil experiment. What was the important discovery in that experiment?
5. Where are the metals, nonmetals, and semimetals on the periodic table?
6. Identify the groups for alkali metals, alkaline earth metals, halogens, and transition metals?
7. Why do elements in the same group (family) have similar properties?
8. State the definition and trend for ionization energy.
9. State the definition and trend for electronegativity.
10. State the definition and trend for the size of atoms.
11. State the trends for the relative sizes of ions to atoms.
12. How many electrons do the following atoms have available for bonding: I, H, Ca, Al, Sb, Rb, S, and C?
13. What types of ions do metals form? Nonmetals?
14. What will be the most likely ion that the following atoms form: Na, P, Mg, Br, Al, and S?
15. Determine the number of protons, neutrons, and electrons in each atom or ion:
91
122
200
2+
31
3Zr
Sb
Hg
P
40
51
80
15
Chemical Bonds
16. Define the octet rule.
17. What is a covalent bond? What types of elements form covalent bonds? How is the octet rule satisfied?
18. What is the difference between polar and nonpolar covalent bonds?
19. What is an ionic bond? What types of elements form ionic bonds? How is the octet rule satisfied?
20. What is a metallic bond?
21. What type of bond does a molecule have?
22. What type of bond does a salt have?
23. What type of bond has a repeating pattern of positive and negative ions? How are these ions held
together?
24. Draw Lewis dot structures for the following:
a. H2O b. CCl4 c. PH3 d. CO2 e. Br2 f. H2SO4
25. Give a physical description of how the atoms and molecules are arranged in solids, liquids, and gases.
26. Use the information from the previous question to describe the shape and volume for each phase (state)
of matter.
27. List the three phases of matter in order of increasing intermolecular attractions.
28. Why do the atoms and molecules in liquids move in a random pattern relative to one another instead of
being in a solid form?
Conservation of Matter and Stoichiometry
29. Balance the following equations:
a. Al2O3 + Cl2 + C  AlCl3 + CO
b. C4H10 + O2  CO2 + H2O
c. P + O2  P2O5
30. What isotope is used to define one mole?
31. How many particles (atoms or molecules) are in one mole?
32. Find the molar mass for the following chemical formulas:
a. ZnF2
b. Al2(SO4)3
c. NH4IO3
33. How many moles are in 4.37 kg of NaOH?
34. How many atoms are in 2075 g of He?
35. What volume will 92.31 g of CO2 occupy at STP?
36. How many grams of sodium hydroxide are needed to react with 6.23 g of barium bromide? NaOH +
BaBr2  NaBr + Ba(OH)2
37. In the reaction, 2 Mg + O2  2 MgO, if 100.0 g of magnesium reacts with 50.0 g of oxygen, what mass
of product is produced?
Gases and Their Properties
38. What is the kinetic molecular theory?
39. How do gases create pressure, use KMT to support your answer.
40. Explain diffusion, use KMT to support your answer.
41. Is Boyle’s law direct or inverse? Charles’s Law? Gay-Lussac’s Law?
42. If 735 L of a gas is at 3.11 atm and 34oC, what is its temperature at 6.11 atm and 235 L?
43. If 12.2 mL of a gas is at 178oC, what is its volume at 53.0oC?
44. What are the values of STP? What is the meaning behind STP?
45. What is the lowest temperature possible? What is it called?
Acids and Bases
46. What are the major observable properties of acids, bases, and salt solutions?
47. What gas is produced when an acid reacts with a metal?
48. What is the Arrhenius definition for acids and bases?
49. What is the Brønsted-Lowry definition for acids and bases?
50. Describe the dissociation (ionization) of strong acids and bases versus weak acids and bases.
51. List the 6 strong acids and state the rule for strong bases.
52. What are the pH values for acids? Bases?
53. What is more acidic, a solution with a pH of 2 or 5? What is more basic, a solution with a pH of 8 or
13?
54. What is the pH value for a neutral substance?
55. List the color for each indicator in an acid solution and a base solution: phenolphthalein, red litmus, and
blue litmus.
Solutions
56. What is a solute? Solvent?
57. Describe the dissolving process at the molecular level by using the concept of random molecular motion.
58. What three factors affect the dissolving process?
59. When 5.20 g of salt is added to 5000 g of water, what is the concentration in parts per million (ppm)?
60. How many grams of KOH would you need to make 750. mL of solution with a concentration of 5.5% by
mass?
61. If you add 25 g of CaCl2 to 1000. mL of water, what would the concentration of the solution be in
grams/liter?
62. What is the molarity of a solution that contains 78.2 grams of NaCl in 4.25 liters of solution?
Chemical Thermodynamics
63. What happens to atoms or molecules as their temperature is increased?
64. Describe heat flow.
65. What is the difference between an exothermic reaction and an endothermic reaction?
66. Draw an energy diagram for an endothermic and exothermic reaction.
67. Are the following changes of state exothermic or endothermic:
a. An ice cube melting
c. Water vapor condensing on a mirror
b. Dry ice subliming to carbon dioxide d. Water freezing into an ice cube
68. How many calories are needed to raise 450. grams of water from 21.0oC to 85.5oC?
69. How many grams of water can be heated 46.0oC by 34.8 kJ?
Reaction Rates
70. Define reaction rate.
71. What are the units for reaction rate?
72. Describe what happens to the concentration of reactants during a chemical reaction. Products?
73. What happens to the reaction rate when there is an increase in concentration? Temperature? Pressure?
74. What is a catalyst? How does the addition of a catalyst affect the rate of a reaction?
75. Draw an energy diagram for a catalyzed reaction.
Chemical Equilibrium
76. Define equilibrium.
77. What is Le Chatelier’s Principle?
78. Which direction will the following reaction shift if: PCl3(g) + Cl2(g) ↔ PCl5(g) + heat
a. Add PCl3
c. Decrease Pressure
b. Remove Cl2
d. Add heat
79. Which direction will the following reaction shift if: H2(g) + I2(g) + 50.9 kJ ↔ 2 HI(g)
a. Add HI
c. Remove I2
b. Decrease temperature
d. Increase pressure
Organic Chemistry and Biochemistry
80. What are polymers?
81. Describe the bonding characteristics of carbon. How many bonds can carbon form? What types of
covalent bonds can carbon form?
82. What type of bond is found in most large organic molecules?
83. What monomers make up proteins? What monomers make up DNA?
84. What is the chemical structure of an amino acid?
85. How many amino acids are found in the human body?
Nuclear Processes
86. What is the strong nuclear force?
87. What is the equation that allows us to calculate the energy released in nuclear reactions?
88. What are the three most common forms of radioactive decay? Describe each type of decay.
89. How does the nucleus change with each type of decay?
90. Complete and balance the following nuclear equations
14
a.
6C
14

7N
218
b.
84Po
c.
11Na
4

22
88Ra
90Th
+ __?___
10Ne
+ __?__
222

86Rn

-1e
234
e.
2He
22

226
d.
+ ___?___
+ __?__
0
+ __?__
91. Describe the different amounts and kinds of damage in matter produced by the different penetrations of
each type of radioactive decay.
92. How does the energy release in a nuclear reaction compare to the energy release in a chemical reaction.
Investigation and Experimentation
93. What is the purpose of an experimental control?
94. Compare accuracy and precision of a measurement.
95. Why is it necessary to perform multiple trials of a scientific experiment?
96. What type of instrument is best for measuring mass, volume, and length?
97. How is the uncertainty of an instrument determined?
98. State the Atlantic-Pacific Rule for determining significant figures.
99. How many significant figures are in the following measurements?
a. 0.000653 g
c. 8.50x10-9 m
b. 24 000 mL
d. 0.025 060 s
True/False
100.
101.
102.
103.
104.
105.
106.
107.
108.
109.
110.
111.
112.
The nucleus of an atom is much larger than the atom yet contains most of its mass.
Atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by
exchanging electrons to form ionic bonds.
The chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2, and
many large biological molecules are ionic.
Large molecules (polymers), such as proteins, nucleic acids, and starch, are formed by repetitive
combinations of simple subunits.
The bonding characteristics of nitrogen result in the formation of a large variety of structures
ranging from simple hydrocarbons to complex polymers and biological molecules.
The energy release per gram of material is much smaller in nuclear fusion or fission reactions
than in chemical reactions.
Some naturally occurring isotopes are radioactive.
The isotopes formed in chemical reactions are radioactive.
The quantity one mole is set by defining one mole of carbon 12 atoms to have a mass of exactly
12 grams.
Acids are hydrogen-ion-accepting and bases are hydrogen-ion-donating substances.
Energy is absorbed when a material condenses or freezes and is absorbed when a material
evaporates or melts.
The rate of a reaction is the decrease in concentration or reactants or the increase of products
with time.
A catalyst increases the rate of a reaction by raising the activation energy of a reaction.