PCHS Honors Chemistry Review
UNIT 1: Lab Safety/Measurement/Matter (Ch 1,2)
1.
Explain the difference between qualitative and quantitative measurement.
Qualitative Measurement- measurement of observations and physical properties
Quantitative Measurement- measurement of quantity that requires units
2. Define accuracy and precision.
Accuracy- closeness of a measurement to a true or known value
Precision- reproducibility of a measurement
3. Solve the following problems and express them in scientific notation
a. (5 x 10-5) + (2x 10-5) 7x 1025
c. (4 x 102) x (1 x 108) 4x1010
8
8
8
b. (7 x 10 ) – (4 x 10 ) 3 x 10
d. (8 x 104) ÷ (4 x 101) 2x103
4. Convert 1.5 kg to g 1500g
5. Beyoncé is 34.2 years old. How many seconds old is she? 1078531200 seconds
6. Define and give examples of 3 physical properties.
Physical Property- property that can be observed or measured without changing the composition of matter
Physical state, density, color, etc.
7. Define and give examples of 3 chemical properties
Chemical Property- property that is observed during a chemical reaction
Reactivity, flammability, odor production, etc.
8. Define and distinguish between pure substances.
Material made up of only one type of particle that cannot be physically separated: elements and compounds
9. Define and distinguish between mixtures.
Combinations of pure substances that can be physically separated
10. Define filtration, distillation, and chromatography.
Filtration- Separation using physical filters
Distillation- separation using differences in boiling point
Chromatography- separation based on light absorption
11. A square block is measured and found to be 3cm across. Its mass is 17 grams. How dense is this block?
17g/9cm3 = 1.88g/cm3
UNIT 2: Nuclear Chemistry (Ch. 4, 6, 24)
12. Explain the contribution of each scientist to our knowledge of the structure of the atom
a. John Dalton Published 1st atomic theory (no experiment)
b. J.J. Thomson Plum Pudding Model
c. Ernest Rutherford Gold Foil Experiment- atoms have positive center in lots of empty space
d. Niels Bohr Electrons are in energy levels around nucleus
e. James Chadwick Discovered neutron (no experiment)
f. Werner Shrodinger Discovered electron orbitals (electron cloud) w/ fancy math
13. Name each subatomic particle and identify each by their location in the atom and charge.
Proton
inside nucleus
1 amu
+
Electron
outside nucleus
0 amu
Neutron
inside nucleus
1 amu
0
14. How is one element different from another element?
Different number of protons
15. How many protons are in Argon? 18
16. How many electrons are in a neutral atom of zinc? 30
17. Which element has atoms with 56 protons? Ba
18. How many neutrons are in U-238? 238-92 = 146
19. Write the nuclear equations for each of the following nuclear reactions
238
𝑈→
a.
Alpha decay of U-238
b.
Alpha and gamma decay of Po-210
c.
Beta decay of Rn-222
92
222
86
𝑅𝑛 →
0
−1
4
𝐻𝑒 +
2
210
84
𝑒+
234
90
4
𝑃𝑜 →
222
87
𝐹𝑟
2
𝑇ℎ
𝐻𝑒 +
0
0
𝛾+
206
82
𝑃𝑏
PCHS Honors Chemistry Review
d.
Beta and gamma decay of Cm-245
245
96
𝐶𝑚 →
0
−1
0
245
0
97
𝑒+ 𝛾+
𝐵𝑘
20. What is an isotope?
An atom with a different number of neutrons
21. Why is the mass of each element on the periodic table a decimal number?
Average mass of all the isotopes of that element
22. Neon consists of three common isotopes Ne-20, Ne-21, Ne-22. Which isotope of neon is most abundant?
Ne-20; closest to the average 20.18
23. Element X consists of three isotopes, 65.43% X-45, 30.51% X-44, and 4.06% X-46. Determine the average atomic mass of the
element X.
45(.6543) + 44(.3051) + 46(.0406) = 44.74 amu
24. Determine the percent abundancy of each isotope in a sample of element Z with an averate atomic mass of 34.68. The
isotopes are Z-33 and Z-35.
34.68= 33(x) + 35(1-x)
= 33x + 35-35x
-.32=/2x
Z-33: 16%
Z-35: 84%
25. Compare and Contrast nuclear fission and nuclear fusion.
Nuclear Fission: Large unstable nuclei break apart into smaller nuclei
Nuclear Fusion: Small nuclei fuse together to form new elements
26. Bandages can be sterilized by exposure to gamma radiation from cobalt-60, which has a half-life of 5.27 yr. How much of a
10.0 mg sample will be left after 1 half life? 5 half-lives?
1 half life: 5 mg 5 half-lives: .3125mg
UNIT 3: Electrons and the Periodic Table (Ch.5)
27.
28.
29.
30.
31.
32.
33.
How is an ion different from a neutral atom? Different number of electrons (charged)
How is a cation different from an anion? Cation has a positive charge, anion has a negative charge
Identify the number of electrons in a P-3 ion. 18 eWhat is the charge on lead with 78 electrons? +4
Why do atoms form into ions? To become more stable
What electron configuration do most ions form? Noble gas configuration
How is the structure of the periodic table related to electron configurations?
s
d
p
34. Why does the 4s orbital fill before the 3d orbital? 4s has less energy
35. Write the electron configurations for Fe and Ag+1.
Fe: 1s22s22p63s23p64s23d6
Ag+1: 1s22s22p63s23p64s23d104p65s24d8
36. What causes the different colors of light seen during a flame test?
e- FALLING in energy levels
37. What color is characteristic of the flame test for: Li, K, Cu, Ca, Ba, Sr?
Li: hot pink/magneta
K: purple/lavender
Cu: blue green
Sr: Red orange
Ca: orange
Na: Orange
Ba: Yellow
38. Explain how Hund’s Rule and the Pauli Exclusionary Principle relate to electron configuration.
Hund’s Rule: e- fill lowest energy levels first; Pauli Exclusionary Principle: 1 e- per sublevel first, then 2e- if they have
opposite spins
39. Identify and explain the following trends as seen on the periodic table:
a. Atomic radius size of atom. Decreases
c. Electronegativity ability to hold onto/pull a
across, increases down
shared electron. Increase across, decrease
b. Ionization energy the energy needed to
down
remove an electron. Increases across,
d. Metal vs. nonmetal properties
decreases down
PCHS Honors Chemistry Review
Metals: react w/ acid; shiny, malleable, lose e-,
conduct electricity, form positive ions.
Nonmetals: do not react w/ acid, not shiny, brittle,
gain e-, do not conduct electricity, form negative
ions
40. Where are metalloids found on the periodic table? Stair step line
UNIT 4: Bonding and Molecular Shapes (Ch. 7-8)
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
What is an ionic bond? Give an example of an ionic compound transfer of e- between a metal and a nonmetal; NaCl
List properties of ionic compounds. Strong, form crystal lattice structure, high melting and boiling points, not brittle
What is a covalent bond? Give an example of a molecular compound shared electrons between 2 nonmetals; CO2
List properties of molecular compounds Weak, low melting points, brittle
What determines if a covalent bond is polar or nonpolar? Polar: unequal sharing of electrons; Nonpolar: equal sharing of
electrons (can determine by using electronegativity values)
Explain the intermolecular forces: dipole-dipole, H-bonding, and London Forces (Induced Dipole)
Dipole-Dipole: found in polar molecules; movement of eHydrogen Ponding: found in polar molecules; when H bonds to N,O,F
London Disperson Forces: between nonpolar molecules; caused by random movement of unshared eWhat types of compounds will exhibit each of these forces? See above
Explain why H2O has a much higher boiling point than H2S. H2O has stronger intermolecular forces with H bonds
What is VSEPR Theory and how does it help us to predict the shapes of molecules? Valence Shell Electron Pair Repulsion:
valence electrons repel so the molecule will adjust its shape so that the electrons are as far apart as possible
Draw electron dot structures for the following molecules. Identify the shape of the molecule
a. NH3 Trigonal Pyramidal
b.
c.
HF Linear
d.
H2S bent
f.
g.
h.
i.
Aluminum hydroxide Al(OH)3
Lead(IV) oxide PbO2
Carbon tetrachloride CCl4
Carbon dioxide CO2
b.
3,4,5-triethyloctane
SO42- Tetrahedral
UNIT 5: Nomenclature and Chemical Reactions (Ch. 7-8,21)
51. Name the following compounds:
a. HNO3 nitric acid
b. Pb(C2H3O2)2 lead (II) acetate
c. NO2 Nitrogen Dioxide
d. CaCl2 calcium chloride
e. P2S5 diphosphorus pentasulfide
52. Draw the following compounds:
a. 2,3-dimethyl-5-propyldecane
53. Name the following compound:
PCHS Honors Chemistry Review
54. Use the following balanced equation to answer the following:
P4(s) + 3O2(g)  2P2O3(s)
a. Name the products phosphorus
b. Name the reactants oxygen
c. What total number of atoms is in the products? 10
d. What are the number of atoms are in the reactants? 10
55. Define the Definite Proportions. A change in subscripts in a compound causes a change in the compound
56. Define the Law of Multiple Proportions. The same elements can combine in different ratios to form new compounds
57. Balance the following Chemical Equations
a. 2C4H10 + 13O2  8CO2 + 10H2O
b. 1Cu + 4HNO3  1Cu(NO3)2 + 2NO2 + 2H2O
c. 2Na + 2H2O  2NaOH + H2
58. Identify the type of reactions below. Write and balance the equations.
a. Lead (II) nitrate + sodium phosphate  sodium nitrate + lead (II) phosphate
DR: 3Pb(NO3)2 + 2 Na3PO4 6NaNO3 + Pb3(PO4)2
b. Sodium + oxygen  sodium oxide
Synthesis: 2Na + O2  Na2O
c. Aluminum hydroxide  aluminum oxide + water
Decomp: 2Al(OH)3  Al2O3 + 3H2O
d. Tricarbon hexahydride + oxygen  carbon dioxide + water
Combustion: 2C3H6 + O2  6CO2 + 6H2O
UNIT 6: The Mole and Stoichiometry (Ch. 10, 11)
59. Determine the molar mass of C3H7O3N 105.0g
60. Determine the mass in grams of 1 atom of zinc.
1 atom x
65.4𝑔
6.022𝑥1023 𝑎𝑡𝑜𝑚𝑠
= 1.09 x 10-22
61. How many molecules are in a 65.4 g sample of carbon dioxide?
65.4g CO2 x
6.022𝑥1023 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠
44.0𝑔 𝐶𝑂2
= 8.95 x 1023 molecules
62. What number of moles is 8.6 x 1024molecules of sulfur?
8.6 x 1024molecules x
1 𝑚𝑜𝑙𝑒
6.022 𝑥 1023𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠
= 14 moles
63. What number of aluminum atoms is in a 65.8 g sample of aluminum oxide?
65.8gAl2O3 x
6.022 𝑥 1023 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠 𝐴𝑙2𝑂3
𝑔𝐴𝑙2𝑂3
x
2 𝑎𝑡𝑜𝑚𝑠 𝐴𝑙
1 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒
= 7.77 x 1023 atoms
64. Determine the empirical formula of a compound that contains 25.44% copper, 12.84% sulfur, 4.036% hydroge, and 57.76%
oxygen.
DOES NOT WORK (WRONG NUMBERS)
65. What is the percent of oxygen in sodium sulfate?
Na2SO4
64.0
142.1
= 45.0% 𝑂
66. Determine the molecular formula of a compound having an empirical formula of CH2 and a molecular weight of 84?
Molecular/empirical = 6 C6H12
Consider the following reaction:
C3H8 + 5O2  3CO2 + 4H2O
67. Determine the mass of water formed by complete reaction of 68.5g of C3H8
68.5 g C3H8 x
4(18.0𝑔)𝐻2𝑂
1(44.0𝑔)𝐶3𝐻8
= 112g H2O
68. What number of moles of carbon dioxide will form from reaction of 46.2 moles of oxygen?
46.2 moles O2 x
3 𝑚𝑜𝑙𝑒𝑠 𝐶𝑂2
5 𝑚𝑜𝑙𝑒𝑠 𝑂2
= 27.7 moles CO2
69. If 685.2g C3H8 react and 10005g of water is formed, what is the percent yield of the reaction?
685.2g C3H8 x
4(18.0𝑔)𝐻2𝑂
1(44.0𝑔)𝐶3𝐻8
= 1,120g H2O
70. If 500g of C3H8 is mixed with 500g of oxygen,
PCHS Honors Chemistry Review
a.
b.
c.
What mass of carbon dioxide is formed? 413 g CO2
Which reactant is in excess? C3H8
How much is left over? 362g
UNIT 7: Thermodynamics (Ch.15)
71. Describe the arrangement of particles in the 3 states of matter in terms of their kinetic energy
Solid- low KE, atoms held tightly together
Liquid- Medium KE, atoms touching but flowing past each other
Gas- High KE, Atoms pinging around like crazy
72. How many J would it take to raise the temperature of 200 g of water from 5 oC to 85oC?
Q = (200 g)(4.18)(80 oC) = 66880 J
73. What is the specific heat capacity of a substance if 750J caused 100g of it to go from 90 oC to 135oC?
750 J = (100 g)(x)(45 oC) = 0.167 J/gC
74. What is the Latent heat of Fusion of a substance? Heat of Vaporization?
Heat of Fusion: Energy needed to melt or freeze
Heat of Vaporization: Energy needed to boil or condense
75. What do you call each of the following changes of state?
a. Solid to liquid melting
d. Solid to gas sublimation
b. Liquid to gas vaporization
e. Liquid to solid freezing
c. Gas to solid deposition
f. Gas to liquid condensation
76. A piece of ice is placed on the sidewalk on a hot summer day. Describe what happens to energy content of the piece of ice,
does it gain or lose energy? What happens to the air around the ice, does it gain or lose energy?
The ice gains energy from the surrounding air and increases its energy (endothermic). The air loses energy to the
ice and decreases its energy (exothermic)
77. 100 grams of hot water at 80 oC is combined with 100 grams of cool water at 20 oC. What is the final temperature of the
combined water?
Qlost = - Qgain mc(Tf – Ti) = - mc(Tf – Ti)(100 g)(4.18)(x- 80 oC) = - (100 g)(4.18)(x- 20 oC) Tf = 50 oC
78. Describe each of the reactions below as endothermic or exothermic.
a. Exploding fireworks exothermic
c. Lava cooling exothermic
b. Melting snow endothermic
d. Water evaporating endothermic
79. What is a calorimeter used for? Measure energy lost or gained by a system
80. Define Enthalpy. Give an example of three ways a change in enthalpy can be shown.
Enthalpy: measure of a change in heat of a system
Thermochemical Equation, Energy Diagrams, H
81. Compare and Contrast Spontaneous and Nonspontaneous reactions.
Spontaneous Reactions- Occur without outside intervention, proceed in one direction, nonreversible
Nonspontaneous Reactions- Require outside force to cause a change
82. Define Entropy and relate it to a change of state/temperature.
Entropy- degree of randomness or chaos in a system
Solid- Low Temp, Low Entropy
Liquid- Higher Temp, Medium Entropy
Gas- High Temp, High Entropy
83. Summarize each of the laws of Thermodynamics.
1st Law: Energy in the universe is constant, it can neither be created nor destroyed, only change forms
2nd Law: Entropy in the Universe is decreasing
3rd Law: At absolute zero, the entropy of a pure substance is 0
UNIT 8: Acids/Bases/Solutions/Kinetics (Ch.16-18)
84. Identify the following conjugates:
a. Conjugate base of H2SO4 SO4
b. Conjugate base of NH4+ NH3
c.
d.
Conjugate acid of H2O H3O+
Conjugate acid of HAsO42- H2AsO4-
PCHS Honors Chemistry Review
85. Label the conjugate acid-base pairs in the reaction below:
HCO3- + OH- (aq)CO32- (aq) + H2O(l)
Acid
Base
C.B
C. A
86. What are two products of an Arrhenius neutralization reaction? Salt and water
87. Compare and contrast properties of acids and bases
Acids- Sour, pH<7, turn blue litmus red, corrosive, produces H gas in contact w/ metal
Base- bitter pH>7, turn red litmus blue, corrosive, soap, slippery
88. What mass of potassium nitrate will dissolve in 100g of water at 60oC?
95g
89. How many more or less grams of potassium chloride will dissolve in 100g of water at 60 oC than grams of potassium nitrate
in the above question? 55g less
90. What mass of potassium chloride will dissolve in 50g of water at 70oC? 25g
91. What is the effect of an increase in temperature on the solubility of a solid? Explain your answer.
Increase in temperature = increase in solubility
92. What is the effect of an increase in temperature on the solubility of a gas?
Increase in temperature = decrease in solubility
93. List 4 things that affect the rate of solution of a solute.
Temperature, surface area, stirring, amount of solute already dissolved
94. Determine the molarity of sodium chloride in a solution that contains 10.5g NaCl dissolved in 750 mL of total solution.
.35M
95. What volume of 0.35 M CoCl2 is needed to obtain 5.35 moles CoCl2?
.065L
96. Describe the changes in the rate of reaction based on the following changes in the system:
a. Increased temperature increased rate
d. Pressure on a gas increased rate
b. Increased Concentration decreased rate
e. Adding Catalysts increased rate
c. Decreased particle size increased rate
97. What does it mean for a reaction to be in equilibrium?
Forward and reverse rate is equal
98. Use LeChatelier’s Principle to predict the shift in equilibrium for the following:
3H2(g) + N2(g) + heat 2NH3(g)
a. When more NH3 is added shifts towards reactants
b. When N2 is removed shifts towards reactants
c. When H2 is added shifts towards products
d. When pressure is increased shifts towards products
e. When temperature is decreased shifts towards products