Chemistry Final Exam Review
Atomic Theory
1. What is the name of Dalton’s model of the atom?
a. Plum pudding model
b. Quantum mechanical model
c. Nuclear model
d. Solid sphere model
2. What is the name of J.J. Thompson’s model of the atom?
a. Plum pudding model
b. Quantum mechanical model
c. Nuclear model
d. Solid sphere model
3. What is the name of Schrodinger’s model of the atom?
a. Plum pudding model
b. Quantum mechanical model
c. Nuclear model
d. Solid sphere model
4. What is the name of Rutherford’s model of the atom?
a. Plum pudding model
b. Quantum mechanical model
c. Nuclear model
d. Solid sphere model
5. Which of the following are results of Rutherford’s model of the atom?
a. Dense positively charged nucleus, a lot of empty space
b. Dense negatively charged nucleus, no empty space
c. Dense positively charged nucleus, no empty space
d. Dense negatively charged nucleus, a lot of empty space
6. Which of the following shows the order of correct charge, mass, and location of a proton in an atom?
a. +1, 0, outside the nucleus
b. +1, 1, outside the nucleus
c. +1, 1, inside the nucleus
d. +1, 0, inside the nucleus
7. Which of the following shows the order of correct charge, mass, and location of an electron in an atom?
a. -1, 0, outside the nucleus
b. -1, 1, outside the nucleus
c. -1, 1, inside the nucleus
d. -1, 0, inside the nucleus
8. How many electrons, protons, and neutrons does a neutral atom of strontium – 89 have?
a. e- = 38, p+= 38, n0= 41
b. e- = 38, p+= 41, n0= 41
c. e- = 41, p+= 38, n0= 41
d. e- = 38, p+= 41, n0= 38
9. How many electrons, protons, and neutrons does a neutral atom of fluorine – 19 have?
a. e- = 19, p+= 19, n0= 10
b. e- = 19, p+= 9, n0= 9
c. e- = 9, p+= 9, n0= 10
d. e- = 10, p+= 10, n0= 9
10. How many electrons, protons, and neutrons does an ion of 9 Be have?
a. e- = 2, p+= 4, n0= 9
b. e- = 2, p+= 4, n0= 5
c. e- = 4, p+= 4, n0= 5
d. e- = 4, p+= 2, n0= 5
11. How many electrons, protons, and neutrons does an ion of 32 P have?
a. e- = 15, p+= 15, n0= 17
b. e- = 15, p+= 17, n0= 15
c. e- = 18, p+= 15, n0= 16
d. e- = 18, p+= 15, n0= 17
12. When an atom of gallium forms a gallium ion, its charge becomes ______ because it _______
electron(s).
a. -3, loses
b. +3, gains
c. +3, loses
d. -3, gains
13. When an atom of bromine forms a bromine ion, its charge becomes _______ because it _______
electron(s).
a. -1, gain
b. -1, lose
c. +1, gain
d. +1, lose
14. Which of the following is not a property of a metal?
a. Conducts electricity in the solid state
b. High melting/boiling point
c. Malleable and ductile
d. Conducts electricity when dissolved in water
15. Which of the following is not a property of a nonmetal?
a. Most solids at room temperature
b. Low melting/boiling point
c. Brittle solids at room temperature
d. Right of the metalloid (semi-metal) “staircase” line
16. Which of the following is a correctly balanced nuclear equation?
a. 23892U  23490Th + 42He
b. 146C  147N + 42He
c. 146C  147N + 0-1 Mg
d. 23892U  23490Th + 42β
17. Which of the following represents alpha decay of radon – 222?
a. 22286Rn  21884Po + 42He
b. 86222Rn  82220Th + 42He
c. 22688Ra  22286Rn + 42He
d. 22286Rn  22287Fr + 0-1β
MOLES
1. What is the mass of 2 moles of propane gas, C3H8?
a. 11 grams
b. 44 grams
c. 22 grams
d. 88 grams
2. How many moles of propane gas, C3H8, are contained in 11 grams of C3H8?
a. 11 × 1023 moles
b. 4 moles
c. 1.5 × 1023 moles
d. 0.25 moles
3. What is the mass in grams of 3 moles of water molecules, H2O?
a. 54 grams
b. 0.166 grams
c. 6 grams
d. 21 grams
4. How many moles of water molecules, H2O, are present in a 27 gram sample of water?
a. 9 × 1023 moles
b. 1.5 moles
c. 2 moles
d. 2/3 mole
5. What is the mass of 10 moles of ammonia, NH3?
a. 1.7 grams
b. 27 grams
c. 170 grams
d. 0.587 grams
6. How many moles of methane, CH4, are in 80 grams of methane?
a. 6.022 × 1080 moles
b. 5 moles
c. 80 × 1023 moles
d. 0.201 moles
7. How many molecules are contained in 3 moles of water, H2O?
a. 6 molecules
b. 54 molecules
c. 1.8 × 1024 molecules
d. 3 × 1023 molecules
8. A sample of carbon dioxide gas (CO2) contains 6.022× 1023 molecules. How many moles of carbon
dioxide does this represent?
a. 1 mole
b. 440 moles
c. 44 moles
d. 10 moles
9. How many molecules of ethane gas, C2H6, are in 15 grams of the compound?
a. 0.5 moles
b. 2 moles
c. 3 × 1023 moles
d. 45 moles
10. What is the mass, in grams, of 3 × 1023 atoms of helium?
a. 2 grams
b. 1.2 × 1024 grams
c. 3 × 1023 grams
d. 8 grams
11. Approximately how many atoms of carbon are present in a 120 gram sample of carbon?
a. 10 atoms
b. 6 × 1022 atoms
c. 6 × 1024
d. 1440 atoms
1 mole of a compound/element = _________________________ particles
1 mole of a compound/element = _________________________ in grams
Molar mass of a compound = 6.022× 1023 particles
ELECTRON CONFIGURATIONS & LIGHT
1. What element has the noble gas configuration [Ne]3s23p1? __________________
2. What element has the electron configuration notation 1s22s22p63s1? ___________
3. Which of the following is the correct noble-gas notation for the element strontium?
a. [Kr]5s1
b. [Xe]5s2
c. [Kr]6s2
d. [Kr]5s2
4.
The above orbital notation is used to represent which element?
a. Boron
b. Sulfur
c. Oxygen
d. fluorine
5.
The above orbital notation is used to represent which element?
a. Phosphorus
b. Arsenic
c. Nitrogen
d. Silicon
6. Which of the following is the correct configuration notation for the element titanium (Ti)?
a. 1s22s22p63s23p64s23d2
b. 1s22s22p63s23p63d24s2
c. 1s22s22p63s23p64s24d2
d. 1s22s22p63s23p64s21d2
7. Which of the following is the correct electron configuration notation for the element nitrogen, N?
a. 1s22s2
b. 1s22s22p6
c. 1s22s22p3
d. 1s22s21p3
8. Which of the following is the correct electron configuration for an aluminum ion?
a. 1s22s22p63s23p1
b. 1s22s22p63s23p6
c. 1s22s22p6
d. 1s22s22p63s2
9. Which of the following electron dot notations is correct for the element phosphorus when it is in the ground state?
a. I
b. II
c. V
d. III
10. Which of the following electron dot notations is correct for the element oxygen when it is in the ground state?
11.
12.
13.
14.
15.
a. I
b. II
c. III
d. V
Which of the following elements has the same number of valence electrons as the element sodium?
a. Ar
b. Cs
c. Ca
d. Mg
Which of the following elements has the same number of valence electrons as the element selenium?
a. Fe
b. K
c. P
d. O
Which of the following elements will have similar physical and chemical properties as lithium?
a. Rubidium
b. Carbon
c. Nitrogen
d. neon
Which of the following elements will have similar physical and chemical properties as Iodine?
a. Te
b. Xe
c. F
d. Po
When an electron gets excited and goes up an energy level, it has ____________ energy and when an electron
goes down an energy level, it has ____________ energy.
a. Absorbed, released
b. Released, absorbed
c. Ground, excited
d. Excited, absorbed
16. When an electron transitions from n = 3 to n = 4, it has ____________.
a. Absorbed energy
b. Released energy
c. Excited state
d. Ground state
17. When an electron transitions from n = 6 to n = 2, it has ___________.
a. Absorbed energy
b. Released energy
c. Ultraviolet radiation
d. Infrared radiation
18. What color light is produced when an electron transitions from the n = 6 to n = 2 energy level?
a. Blue
b. Red
c. No color
d. violet
19.
Given the representation of a chlorine atom, which circle might represent an atom of bromine?
a.
b.
c.
d.
Circle B
Circle D
None of these
Circle C
20.
Given the representation of a chlorine atom, which circle might represent an atom of sulfur?
a.
b.
c.
d.
None of these
Circle B
Circle D
Circle C
21. As one moves from left to right ( → ) within a period across the periodic table, the electronegativity of
the elements encountered tends to:
a. stay the same
b. increase
c. decrease
22. The elements with the largest atomic radii are found in the:
a.
b.
c.
d.
lower right-hand corner of the periodic table
lower left-hand corner of the periodic table
upper right-hand corner of the periodic table
upper left-hand corner of the periodic table
23. Of the following elements, which one would have the largest radius?
a.
b.
c.
d.
Cesium (Cs, atomic #55)
Potassium (K, atomic #19)
Hydrogen (H, atomic #1)
Sodium (Na, atomic #11)
24. The energy required to remove an electron from an atom is known as:
a.
b.
c.
d.
radioactivity
electron affinity
ionization energy
electronegativity
25. Of the following elements, which one would have the largest ionization energy?
a.
b.
c.
d.
Cesium (Cs, atomic #55)
Hydrogen (H, atomic #1)
Sodium (Na, atomic #11)
Potassium (K, atomic #19)
BONDING AND INTERMOLECULAR FORCES
1. Which of the following gases does not exist in nature as a diatomic molecule?
a. Nitrogen
b. Helium
c. Hydrogen
d. oxygen
2. Ionic compounds generally form:
a. Liquids
b. Gases
c. Crystals
d. molecules
3. In metallic bonding, the valence electrons of all atoms are shared in:
a. A nonpolar covalent bond
b. An electron sea
c. A polar covalent bond
d. Transferred to metallic ions
4. The metalloids possess properties of metals and nonmetals and are also known as
a. Semimetals
b. Halogens
c. Gases
d. Liquids
5. The seven elements that occur as diatomic elements are
a. H2,N2,O2,He2,Ne2,C2,Na2
b. H2,N2,O2,He2,Ne2,Cl2,Br2
c. H2,N2,O2,F2,I2,Cl2,Br2
d. Fe2,Rn2,O2,He2,Ne2,C2,Br2
6. The bond between sodium and oxygen is expected to be
a. Gaseous
b. Nonpolar covalent
c. Ionic
d. Polar covalent
7. When compared to single bonds, double bonds are generally
a. Shorter and stronger
b. Longer and stronger
c. Longer and weaker
d. Shorter and weaker
8. The bond between lithium and fluorine is
a. Polar covalent
b. Ionic
c. Nonpolar covalent
d. metallic
9. In the ionic compound magnesium fluoride, what is the ratio of the two elements necessary so that each
element obtains its octet from the transfer of electrons?
a. 3 magnesium: 1 fluorine
b. 1 magnesium: 1 fluorine
c. 2 magnesium: 1 fluorine
d. 1 magnesium: 2 fluorine
10. In the correct Lewis structure for water, how many unshared pairs of electrons will oxygen have?
a. 1
b. 3
c. 4
d. 2
11. In the correct Lewis structure for the methane molecule, how many unshared electron pairs surround the
carbon?
a. 2
b. 0
c. 8
d. 4
12. In nonpolar covalent bonds, valence electrons are
a. Equally shared
b. Unequally shared
c. Destroyed
d. transferred
13. Which of the following is an acceptable Lewis structure for chloromethane (CH3Cl)?
a.
b.
c.
d.
14. In a diatomic molecule of an element, the bond between the atoms must be
a. Nonpolar covalent
b. Polar covalent
c. Metallic
d. ionic
15. In polar covalent bonds, valence electrons are
a. transferred
b. unequally shared
c. destroyed
d. equally shared
16. In ionic bonds, valence electrons are
a. Equally shared
b. Transferred
c. Unequally shared
d. destroyed
17. How many atoms are needed to provide the electrons necessary to complete the valence octet of an
oxygen atom?
a. Three sodium atoms
b. Two sodium atoms
c. Four sodium atoms
d. One sodium atom
18. The measure of the attraction that an atom has for electrons involved in chemical bonds is known as
a. Ionization energy
b. Radioactivity
c. Electronegativity
d. Electron affinity
19. Which of the following is the correct Lewis structure for ammonia?
a.
b.
c.
d.
20. In drawing Lewis structures, a single line (single bond) between two elements represents
a. An octet of electrons
b. An unshared pair of electrons
c. A shared electron
d. A shared pair of electrons
21. Which of the following is a correct Lewis structure for hydrogen cyanide, HCN?
a.
b.
c.
d.
22. Which of the following is the correct Lewis structure for formaldehyde, CH2O
a.
b.
c.
d.
23. Which of the following is the correct Lewis structure for phosphorus tribromide?
a.
b.
c.
d.
24. Which of the diatomic elements has a double bond between its atoms?
a. Fluorine
b. Nitrogen
c. Oxygen
d. Hydrogen
NOMENCLATURE, EMPIRICAL AND MOLECULAR FORMULAS
1. What is the correct name for ClO2?
a. Chlorine dioxide
b. Monochlorine dioxide
c. Dichlorine monoxide
d. Chlorine oxide
2. What is the name of SiCl4?
a. Silicon chloride
b. Monosilicon tetrachloride
c. Silicon tetroxide
d. Silicon tetrachloride
3. What is the correct formula for nitrogen monoxide?
a. NO2
b. No
c. NO
d. N2O
4. What is the correct formula for the compound tetraphosphorus trisulfide?
a. P3S4
b. P5S4
c. 4PS2
d. P4S3
5. What is the correct name of CO?
a. Monocarbon monoxide
b. Carbon dioxide
c. Carbon dioxide
d. Carbon oxide
6. What is the correct formula for boron trifluoride?
a. B3F
b. BF3
c. 3BF
d. B3F3
7. What is the correct formula for carbon tetrabromide?
a. CB4
b. C4Br
c. C4Br
d. CBr4
8. What is the correct formula for dinitrogen monoxide?
a. N2O2
b. N2O
c. NO2
d. 2NO
9. The correct name for the compound Fe2S3 is
a. Diiron trisulfide
b. Iron (III) sulfide
c. Iron sulfide
d. Iron (II) sulfide
10. The correct formula for chromium (III) oxide is
a. Cr3O3
b. Cr3O2
c. Cr2O2
d. Cr2O3
11. The correct name for the compound Cu3N is
a. Copper (II) nitrate
b. Copper (III) nitride
c. Copper nitrogen
d. Copper (I) nitride
12. The correct formula when calcium and nitrogen bond together would be
a. Ca2N3
b. Ca3NO3
c. Ca3(NO3)2
d. Ca3N2
13. The correct formula when nickel (II) and bromine combine would be
a. NiBr
b. Ni2Br
c. NiBr2
d. NiB2
14. The formula that would be made when aluminum and the phosphate ion bond would be
a. AlPO4
b. Al3(PO4)3
c. Al3(PO4)
d. Al(PO4)3
15. If an element in group 1 was to combine with an element in group 7 (17), the resulting compound would
have a ratio of:
a. 1 : 1
b. 1 : 2
c. 2 : 1
d. 3 : 3
16. If an element in group 2 was to combine with an element in group 5 (15), the resulting compound would
have a ratio of:
a. 2 : 3
b. 3 : 2
c. 2 : 5
d. 5 : 2
17. If a molecular formula for some compound is X2Y6, a possible empirical formula could be
a. XY2
b. X2Y3
c. XY3
d. X3Y
18. What is a possible molecular formula for a compound with the empirical formula of CH3?
a. C2H3
b. C3H6
c. C3H
d. C3H9
19. Analysis shows that some compound has the following percent composition: 40.05% S and 59/95 % O.
What is its empirical formula?
a. S3O
b. SO3
c. SO
d. S3O6
20. Propane is a hydrocarbon. it is 81.82 % carbon and 18.18 % hydrogen. What is the empirical formula?
a. CH8
b. C8H3
c. CH3
d. C3H8
21. Analysis shows that some compound has the following percent composition: 48.64% C, 8.16 % H, and
43.20 % O. What is its empirical formula?
a. C1.5H3O
b. C2H3O
c. CHO
d. C3H6O2
22. The hydrocarbon used in manufacture of foam plastics is called styrene. Analysis of styrene indicates
the compound is 92.25 % C and 7.75 % H and has a molar mass of 104 g/mol. What is the molecular
formula for styrene?
a. C8H
b. CH
c. C8H8
d. CH4
23. A colorless liquid composed of 46.68 % N and 53.32 % O has a molecular mass of 60.01 g/mol. What
is the molecular formula?
a. NO
b. N2O
c. N2O2
d. N4O2
24. The empirical formula of a compound is C3H3O. The molecular mass is 110.0 g/mol. What is its
molecular formula?
a. C3H3O
b. CHO3
c. C6H6O
d. C6H6O2
25. The empirical formula of a compound is PNCl2. The molecular mass is 695 g/mol. What is its
molecular formula?
26. Calculate the % composition of sodium sulfate.
27. What is the percent composition of H3PO4?
28. Determine the % by mass (% composition) of sucrose, C12H22O11.
a. 42.10 % C, 6.480 % O, 51.42 % H
b. 58.23 % C, 7.65 % O, 34.12 % H
c. 44.0 % C, 6.75 % H, 53.98 % O
d. 42.10 % C, 6.480 % H, 51.42 % O
CHEMICAL REACTIONS
Define these terms:
Single replacement reaction
Combination reaction
Decomposition reaction
Precipitate
Reactants
Products
Coefficient
Chemical equation
Balanced chemical equation
Double replacement reaction
Diatomic molecule
Catalyst
Yield sign
Activity series
Law of conservation of mass
Subscript
Combustion reaction
Aqueous
1. What are the 5 types of chemical reactions?
2. Know how to identify the different types of chemical
reactions.
Examples:
a) FeCl3 + NaOH → Fe(OH)3 + NaCl
b) Al + O2 → Al2O3
c) C2H2 + O2 → CO2 + H2O
d) Na + H2O → NaOH + H2
e) KClO3 → KCl + O2
3. Know how to balance equations.
Example:
a) __ FeCl3 + __ NaOH → __ Fe(OH)3 + __ NaCl
b) __ Al + __ O2 → __ Al2O3
c) __ C2H2 + __ O2 → __ CO2 + __ H2O
d) __ Na + __ H2O → __ NaOH + __ H2
e) __ KClO3 → __ KCl + __ O2
4. Know how to predict products as in the problems
below.
Example:
a) Al + N2 →
b) H2O →
c) Ca + H2O →
d) Cl2 + NaBr →
e) FeS + HCl →
5. What is an activity series chart? What type of reaction
do you use it for?
a) Using the activity chart, why can sodium replace
hydrogen?
6. What are 5 indicators/observations of a chemical
reaction?
7. List the chemical formulas for the 7 diatomic
molecules.
8. Know how to translate chemical equations and
balance them appropriately.
Example:
a) ammonium chloride reacts with calcium
hydroxide to form calcium chloride and nitrogen
trihydride (ammonia) and water
b) sodium oxide and water yield sodium hydroxide
STOICHIOMETRY
Mole ratio
1. What do the coefficients mean in a chemical equation?
2. Know how to calculate the mole ratio between reactants and products in a chemical formula.
a) What is the mole ratio for calcium and oxygen in 2Ca + O2 → 2CaO
3. Know how to solve mole to mole, mole to mass, mass to mole, mass to mass problems.
a) How many moles of lithium hydroxide are required to react with 20. mol of carbon dioxide? CO2 + 2LiOH →
Li2CO3 + H2O
b) What mass, in grams, of glucose is produces when 3.00 mol of water react with carbon dioxide? 6CO2 + 6H2O →
C6H12O6 + 6O2
c) How many moles of NO are formed when 824 g of ammonia reacts with an excess of oxygen? (balance the
equation first) NH3 + O2 → NO + H2O
d) How many grams of SnF2 are produced from the reaction of 30.0 g HF with Sn? Sn + 2HF → SnF2 + H2
GASES
Ideal gas law
directly proportional
inversely proportional
molar volume
STP
Ideal gas constant
Partial pressure
Dalton’s law
kinetic molecular theory elastic collision
law of
combining volumes
1. Know the 5 assumptions of the kinetic molecular
d) The volume of a sample of oxygen gas is 300.0 ml
theory.
when the pressure is 1.00 atm and the temperature is
2. Know the difference between an ideal gas and a real
27.0 C. At what temperature would the volume
gas.
change to 1.00 L and the pressure change to 0.500
3. Explain a gas based on the following properties:
atm?
density, compressibility, diffusion, effusion, fluidity,
e) A sample of gas at 25.0 C has a volume of 11.0 L
shape , IMF, particle arrangement, and volume
and exerts a pressure of 660.0 mmHg. How many
expansion.
moles of gas are in the sample?
4. Define pressure. What are some common pressure
f) A sample of gas in a closed container at a
units?
temperature of 100. C and a pressure of 3.0 atm is
5. Know how to convert pressure units:
heated to 300. C. What pressure does the gas exert at
a) convert .200 atm to mmHg
the higher temperature?
b) convert 345.8 kPa to atm
8. Use the law of combining volumes, Avogadro’s law,
c) convert 760 mmHg to kPa
and molar volume to solve
6. What is standard temperature and standard pressure?
these problems.
7. Know how to solve problems using Boyle’s law,
a) 3O2 → 2O3 Both gases are measured at the same
Charles law, Gay-Lussac, Combined, Ideal, Density
temperature and pressure. How many liters of O2 are
and Molar mass using the Ideal gas law and Dalton’s
required to make 24 L of O3 ?
law of partial pressure.
b) How many liters of O3 are formed from 12 mol of O2
a) A gas occupies a volume of 200. ml at 100. mmHg.
at STP?
What volume will the gas occupy at 300. mmHg?
11. Know these answers:
b) Air has a total pressure of 20.6 atm and contains
a) As the temperature of a gas decreases, the volume of
carbon monoxide, oxygen, and nitrogen. If air is
a gas will ____________.
made up of 0.6 atm of carbon monoxide, 12.6 atm of
b) As the temperature of a gas decreases, the pressure
oxygen, what would be the partial pressure of
of the gas will ____________.
nitrogen?
c) As the volume of the gas decreases, the pressure of
c) If a sample of gas occupies 15.9 L at 34 C, what will
the gas will ____________.
its volume be at 27 C if the pressure does not
change?
THERMOCHEMISTRY/SOLIDS & LIQUIDS
Know these terms
Vaporization
Condensation
Evaporation
Melting point
Freezing point
Sublimation
Triple point
Melting
Freezing
Deposition
Phase Diagram
Boiling
Endothermic reaction
Exothermic reaction
Heating &Cooling Curve
Specific heat capacity
temperature
heat
1. State the 6 phase changes of state and which ones
work in opposition to each other. i.e. sublimation
and deposition
2. Explain how a solid melts into a liquid using kinetic
energy in your explanation.
3. What 2 temperatures measure the same amount
during a phase change of a liquid pure solvent to a
solid?
4. Know how to read phase diagrams. Sketch a quick
diagram locating the triple point, critical point, the
melting point /freezing point line and the boiling
point/condensation point line. Also label the 3
sections as solid , liquid, and gas.
5. Know how to read a heating and cooling curve.
What do the plateaus tells you? What do the slopes
tell you? Where is the KE of the substance
constant?
6. Sketch an endothermic reaction graph, labeling the
reactants, products, activation energy, activated
complex, and the heat of reaction.
7. What is the sign of an endothermic reaction and
exothermic reaction?
8. Using the specific heat values for water and iron,
which one would have the largest temperature
change if they have the same mass?
9. Know how to calculate the heat released or absorbed
during a physical change.
a.
Calculate the heat absorbed when 15.0 g
of ice melts to liquid. See reference sheet
for Hfus
b. Calculate the heat released when 75.4 g of
vapor condenses into liquid. See
reference sheet for Hvap
10. Know how to calculate the heat released or absorbed
in a chemical
reaction?
Know these terms
Rate
Collision theory
Transition state
a) What is the specific heat of a metal that releases
2500 J of energy. The metal has a mass of 25 g
and had a temperature change of 5C.
b) How much heat is released when iron is dropped
in a beaker of water. The mass of the metal was
43 g and the initial temperature of the metal was
78 C. The water temperature changed from 25
C to 32 C. The specific heat of the metal is
.45J/gC.
c) What is the amount of heat absorbed by water if
23.4 g of water is heated from 34C to 78 C. See
reference sheet for specific heat of water.
KINETICS & EQUILIBRIUM
Activation energy
Activated complex
Catalyst
1. Explain the three criteria of the collision theory.
2. On the pathway below, label the activated complex, activation energy with
catalyst, and activation energy without catalyst
3 What are the five factors that affect the rate of a reaction?
4. Which of the five factors change collision frequency?
5 Which factor changes collision frequency and the energy of the collisions?
6. How does rate change if you increase the concentration of the reactants?
7. How does rate change if you increase the surface area?
8. How does rate change if you decrease the temperature?
9. How does rate change if you add a catalyst?
10. Write the equilibrium expression for the following reaction.
a) H2(g) + Cl2(g) 2HCl(g) + heat
11. In the process of chemical equilibrium, what stays constant at equilibrium?
12. In the process of equilibrium, are the rates equal to each other?
13 Using the reaction above, answer the following questions regarding Le Chatelier’s
principle.
a) Which direction does the reaction shift if temperature increases?
b) Which direction does the reaction shift if hydrogen gas is increased?
c) Which direction does the reaction shift if HCl is removed?
d) Which direction does the reaction shift if the volume is decreased?
e) Which direction doe the reaction shift if temperature is decreased?
14. If K = .00045, what side of the reaction will be favored?
SOLUTIONS
Know these terms:
Solution
solute
soluble
insoluble
miscible
immiscible
electrolytes (strong and weak)
non-electrolytes
solubility
solvent
supersaturated solution aqueous solution
unsaturated solution
saturated solution
Henry’s law
molarity
Know the following:
1. Explain the like dissolves like rule and give an example following the rule.
2. Name 3 factors that increase the rate of dissolution of a substance.
3. Describe solution equilibrium.
4. Name substances that are considered electrolytes and non-electrolytes.
5. What is the effect of temperature and pressure on gas solubility?
6. What is the effect of temperature on the solubility for most ionic solids?
7. Know how to calculate molarity
a) What is the molarity of 4.5 moles of Ba(OH)2 in 10.0 L?
b) A solution has a molarity of 2.8 M and a volume of 250 ml. How many moles of solute are in the solution?
8. Know how to read a solubility graphs.
a. Using the solubility graph from the notes, how much of NaCl can be dissolved at 45C
b. Using the solubility graph from the notes, 50 g of KClO3 is dissolved in 100 g of water at 45C.
Is the solution saturated or unsaturated?
9. Know how to solve dilution problems.
a) How many ml of a 2.0 M NaBr solution are needed to make 200 ml of a 0.50 M solution?
10. Which types of substances produce electrolytes?
11. Which type of substances produce non electrolytes?
ACIDS & BASES
Know these terms:
Arrhenius acid
Arrhenius base
Bronsted-Lowry acid
Bronsted-Lowry base
pH
conjugate acid
Conjugate base
hydroxide ion
hydronium ion
neutralization
titration
equivalence point
1. List some common properties of an acid.
2. List some common properties of a base.
3. Define self-ionization of water.
4. Know how to predict the products and balance neutralization (double replacement) reactions.
a) H2CO3 + Fe(OH)3 →
5. Know how to calculate the pH from hydrogen and hydroxide ion concentrations
a) What is the pH of a [OH-] = 1 x 10-5 M?
b) What is the pH of a [H+] = 1 x 10-5 M?
c) What is the pOH of a [H+] = 1 x 10-1 M?
d) What is the pOH of a [OH-] = 1 x 10-12 M?
6. What is the hydrogen ion concentration of 0.001 M HNO3? What is the [OH-]?
7. What is the hydrogen ion concentration of [OH-] = 3.0 x 10-2 M? What is the pH?
8. What is the pH of a solution if the [H+] = 3.4 x 10-5 M? What is the hydroxide concentration?
9. Determine the pH of a 2.0 x 10-2 M Sr(OH)2?
10. The pH of a solution is measured and determined to be 7.52? What is the hydrogen ion concentration? Is the
solution acidic or basic?
11. Know how to look at an equation and predict Bronsted-Lowery acids and bases and conjugate acids and conjugate
bases.
a) NH4+ + H2O → NH3 + H3O+
What is the base? What is the conjugate base? What is the acid? What is the conjugate acid?
12. What are the products of neutralization?
13. Know how to name acids and bases
a) HF
b) H2SO4
c) NaOH
d) HNO2
e) Fe(OH)2
14. In a titration, how much of .15 M NaOH is needed to neutralize 20 ml of .500M HCl solution? HCl + NaOH 
H2O + NaCl
15. In a titration, what is the molarity of HNO3 if
25 ml of it neutralized 15 ml of .60M Ca(OH)2
2 HNO3 +
Ca(OH)2  2 H2O + Ca(NO3)2
16. What is the difference between end point and equivalence point?