Download Chemistry Final Exam Review 2006-2007

Chemistry Final Exam Review
Atomic Theory
1. What is the name of Dalton’s model of the atom?
10.
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?
11.
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?
12.
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
13.
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.
b.
c.
d.
Dense positively charged nucleus, a lot of empty space
Dense negatively charged nucleus, no empty space
Dense positively charged nucleus, no empty space
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= 51
b. e- = 38, p+= 51, n0= 51
c. e- = 51, p+= 38, n0= 51
d. e- = 38, p+= 51, 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
14.
15.
16.
17.
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
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
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
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
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
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
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
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
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
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.
16.
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
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.
20.
Circle B
Circle D
None of these
Circle C
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
a. 1
nature as a diatomic molecule?
b. 3
a. Nitrogen
c. 4
b. Helium
d. 2
c. Hydrogen
11. In the correct Lewis structure for the methane
d. oxygen
molecule, how many unshared electron pairs
2. Ionic compounds generally form:
surround the carbon?
a. Liquids
a. 2
b. Gases
b. 0
c. Crystals
c. 8
d. molecules
d. 4
3. In metallic bonding, the valence electrons of all
12. In nonpolar covalent bonds, valence electrons are
atoms are shared in:
a. Equally shared
a. A nonpolar covalent bond
b. Unequally shared
b. An electron sea
c. Destroyed
c. A polar covalent bond
d. transferred
d. Transferred to metallic ions
13. Which of the following is an acceptable Lewis
4. The metalloids possess properties of metals and
structure for chloromethane (CH3Cl)?
nonmetals and are also known as
a. Semimetals
b. Halogens
c. Gases
d. Liquids
5. The seven elements that occur as diatomic
a.
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
b.
a. Gaseous
b. Nonpolar covalent
c. Ionic
d. Polar covalent
7. When compared to single bonds, double bonds
are generally
c.
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
d.
b. Ionic
14. In a diatomic molecule of an element, the bond
c. Nonpolar covalent
between the atoms must be
d. metallic
a. Nonpolar covalent
9. In the ionic compound magnesium fluoride, what
b. Polar covalent
is the ratio of the two elements necessary so that
c. Metallic
each element obtains its octet from the transfer of
d. ionic
electrons?
15. In polar covalent bonds, valence electrons are
a. 3 magnesium: 1 fluorine
a. transferred
b. 1 magnesium: 1 fluorine
b. unequally shared
c. 2 magnesium: 1 fluorine
c. destroyed
d. 1 magnesium: 2 fluorine
d. equally shared
10. In the correct Lewis structure for water, how
many unshared pairs of electrons will oxygen
16. In ionic bonds, valence electrons are
have?
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?
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.
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.
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?
c. Copper nitrogen
a. Chlorine dioxide
d. Copper (I) nitride
b. Monochlorine dioxide
12. The correct formula when calcium and nitrogen
c. Dichlorine monoxide
bond together would be
d. Chlorine oxide
a. Ca2N3
2. What is the name of SiCl4?
b. Ca3NO3
a. Silicon chloride
c. Ca3(NO3)2
b. Monosilicon tetrachloride
d. Ca3N2
c. Silicon tetroxide
13. The correct formula when nickel (II) and bromine
d. Silicon tetrachloride
combine would be
3. What is the correct formula for nitrogen
a. NiBr
monoxide?
b. Ni2Br
a. NO2
c. NiBr2
b. No
d. NiB2
c. NO
14. The formula that would be made when aluminum
d. N2O
and the phosphate ion bond would be
4. What is the correct formula for the compound
a. AlPO4
tetraphosphorus trisulfide?
b. Al3(PO4)3
a. P3S4
c. Al3(PO4)
b. P5S4
d. Al(PO4)3
c. 4PS2
15. If an element in group 1 was to combine with an
d. P4S3
element in group 7 (17), the resulting compound
5. What is the correct name of CO?
would have a ratio of:
a. Monocarbon monoxide
a. 1 : 1
b. Carbon dioxide
b. 1 : 2
c. Carbon monoxide
c. 2 : 1
d. Carbon oxide
d. 3 : 3
6. What is the correct formula for boron trifluoride?
16. If an element in group 2 was to combine with an
a. B3F
element in group 5 (15), the resulting compound
b. BF3
would have a ratio of:
c. 3BF
a. 2 : 3
d. B3F3
b. 3 : 2
7. What is the correct formula for carbon
c. 2 : 5
tetrabromide?
d. 5 : 2
a. CB4
17. If a molecular formula for some compound is
b. C4Br
X2Y6, a possible empirical formula could be
c. C4Br
a. XY2
d. CBr4
b. X2Y3
8. What is the correct formula for dinitrogen
c. XY3
monoxide?
d. X3Y
a. N2O2
18. What is a possible molecular formula for a
b. N2O
compound with the empirical formula of CH3?
c. NO2
a. C2H3
d. 2NO
b. C3H6
9. The correct name for the compound Fe2S3 is
c. C3H
a. Diiron trisulfide
d. C3H9
b. Iron (III) sulfide
19. Analysis shows that some compound has the
c. Iron sulfide
following percent composition: 40.05% S and
d. Iron (II) sulfide
59.95 % O. What is its empirical formula?
10. The correct formula for chromium (III) oxide is
a. S3O
a. Cr3O3
b. SO3
b. Cr3O2
c. SO
c. Cr2O2
d. S3O6
d. Cr2O3
11. The correct name for the compound Cu3N is
a. Copper (II) nitrate
b. Copper (III) nitride
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
Define these terms:
Single replacement reaction
Combination reaction
Decomposition reaction
Precipitate
Reactants
Products
Coefficient
Chemical equation
Balanced chemical equation
CHEMICAL REACTIONS
Double replacement reaction
Diatomic molecule
Catalyst
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
Yield sign
Activity series
Law of conservation of mass
Subscript
Combustion reaction
Aqueous
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 and
shape , IMF, particle arrangement, and volume
exerts a pressure of 660.0 mmHg. How many moles
expansion.
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 temperature
units?
of 100. C and a pressure of 3.0 atm is heated to 300. C.
5. Know how to convert pressure units:
What pressure does the gas exert at the higher
a) convert .200 atm to mmHg
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 a
carbon monoxide, oxygen, and nitrogen. If air is made
gas will ____________.
up of 0.6 atm of carbon monoxide, 12.6 atm of
b) As the temperature of a gas decreases, the pressure of
oxygen, what would be the partial pressure of
the gas will ____________.
nitrogen?
c) As the volume of the gas decreases, the pressure of the
c) If a sample of gas occupies 15.9 L at 34 C, what will its
gas will ____________.
volume be at 27 C if the pressure does not change?
Know these terms
Vaporization
Freezing point
Freezing
Endothermic reaction
temperature
THERMOCHEMISTRY/SOLIDS & LIQUIDS
Condensation
Sublimation
Deposition
Exothermic reaction
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?
Evaporation
Melting point
Triple point
Melting
Phase Diagram
Boiling
Heating &Cooling Curve Specific heat capacity
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?
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.
Know these terms
Rate
Collision theory
Transition state
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?