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SEMESTER 1
FINAL EXAM REVIEW
Unit I – Chemical Foundations
I. Good things to know:
 chemistry
 scientific method: hypothesis, data, theory, variables
 matter  substances: elements, compounds; mixtures: homogeneous, heterogeneous
 graphing: x-axis, independent variable; y-axis, dependent variable
 graphing: directly proportional, inversely proportional
 commonly used lab equipment and how to use properly
 lab safety, hazard symbols
 significant figures  their importance in measurement
 scientific notation: base, power  where are the sig figs?
 accuracy vs. precision
 SI system: what are the base units? what units are appropriate in what situations?
 density; intensive properties
II. Problems
_______________________ 1) Which of the following would be considered proper laboratory procedure?
a. determining the odor of a substance by gently wafting the vapors
if you know the substance is not harmful
b. weighing a crucible while it is still hot
c. measuring a liquid in a graduated cylinder by taking a reading from
the top of the liquid
d. rinsing a pipet with the test solution before performing a titration
e. using a pipet bulb to draw liquid into a pipet
f. pulling your goggles up briefly to rub your eyes
g. measuring the volume of a liquid in a beaker
h. diluting sulfuric acid by measuring out the desired amount of water,
and then slowly pouring the concentrated acid to the water
2). How many sig figs in the following numbers?
_______ a) 0.003
3)
________ b) 10.0
51348
.
 10 30
6.012  10 6.938  10 
34
5) 687 mg = ? kg
3

________ c) 550
4)
6) 80.10 km = ? um
_________ d) 310.01
13,4300.0235

40.017240.01
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7) 4.01 x 10 J = ? kJ
8) You are in Paris and want to buy some peaches for lunch. The sign on the fruit stand says that peaches
cost 4.00 euros per kilogram. Assuming that there are 1.14 euros to the dollar, calculate the cost
of a pound of peaches in dollars.
9) If the speed of light is approximately 3.00x 10 8 m/s, what is the approximate speed in miles/hr?
SEMESTER 1
FINAL EXAM REVIEW
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10) If you have a graduated cylinder which initially contains 15.7 mL of water, and when a 65.64 g object
is placed inside, shows a new volume of 22.6 mL, what is the object’s density?
11) The approximate mass of the earth is 5.94 x 10 21 metric tons. If the circumference of the Earth at the
equator is approximately 25,000 miles, what is the approximate density of the earth in g/mL ?
Vsphere = 4/3 πr3
C = 2πr
1 metric ton = 1000 kg
Unit II – Atomic Theory & The Periodic Table
1. Explain the Photoelectric effect as it pertains to quantam theory: Why is it that purple light will expel
electrons from a piece of metal, but red light will not? Why is it that as the intensity of the light increases,
the expelled electrons do not become more energetic, there are simply more of them expelled?
2. Briefly describe the concept of DeBroglie’s Equation. How does the wavelength vary with mass?
3. Briefly describe Bohr’s Atomic Theory for the hydrogen atom. What phenomenon gave rise to his
theory? How do the colors of the bright-line spectra for hydrogen relate to the Energy levels in a hydrogen
atom?
4. Explain why adding an electron to a fluorine atom is an exothermic process.
5. Explain why adding an electron to a neon atom is an endothermic process.
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SEMESTER 1
FINAL EXAM REVIEW
6. Explain the trend in atomic radius from Li to Na to K.
7. Explain the trend in atomic radius from Na to Mg to Al.
8. Explain the trend in first ionization energy from Na to Mg to Al.
9. Explain why the 2nd ionization energy for sodium is so much greater than that for magnesium.
10. What would the first two quantum numbers be for an electron in a 4f subshell? How many possible
quantum number sets are there for a 4f subshell? Give two examples. How many electrons can
be held in the entire 4th energy level?
11. Fill in the table below
Element
Family
Name
# of Valence
Electrons
Potassium
Magnesium
Cadmium
Dysprosium
Chlorine
Neon
XX
XX
Physical State
(metal, nonmetal, metalloid)
Electron Configuration
Paramagnetic
or Diamagnetic?
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SEMESTER 1
FINAL EXAM REVIEW
_______ 12. Which of the following could be the quantum numbers for the valence electron in a ground
state sodium atom?
a. (3,0,0,½) b. (3,1,1,½) c. (4,0,0,½)
d. (4,1,1,½)
e. (4,2,1,½)
_______ 13. Which of the following is an impossible set of quantum numbers? Explain your answer.
a. (3,0,0,½) b. (3,1,1,-½)
c. (3,1,-2,½)
d. (3,2,-1,½)
e. (3,2,2,½)
14. Which fundamental atomic theory is violated by the following list of quantum numbers representing
silicon’s 4 valence electrons?
(3,0,0,-½), (3,0,0,-½), (3,1,1,-½), (3,1,1,-½)
15. Which fundamental atomic theory is violated by the following list of quantum numbers representing
silicon’s 4 valence electrons?
(3,0,0,-½), (3,0,0,+½), (3,1,-1,-½), (3,1,-1,+½)
16. Which of the following atoms or ions is larger:
______ a. sulfur or sulfide
______ b. sodium atoms or sodium ions
______ c. Na or K
______ d. S or Cl
______ e. Na+ or Mg+2
______ f. F-1 or Cl-1
17. Explain why a fluorine atom gains in size when it accepts an electron to become a fluoride ion.
18. Give the number of protons, neutrons, and electrons in the following:
a)
131
54
Xe
b)
56
26
Fe 3
c)
127
53
I 1
______ protons
______ protons
______ protons
______ neutrons
______ neutrons
______ neutrons
______ electrons
______ electrons
______ electrons
SEMESTER 1
FINAL EXAM REVIEW
Unit III – Nomenclature & Chemical Equations
1) Give the name or formula for the following:
_____________ a. sulfuric acid
__________________________________ k. NH3
_____________ b. calcium hydroxide
__________________________________ l. NiS
_____________ c. xenon hexafluoride
__________________________________ m. (NH 4)2CO3
_____________ d. iron(II) dichromate
__________________________________ n. AlCl3
_____________ e. hydrocyanic acid
__________________________________ o. Cr(OH)3
_____________ f. rubidium oxide
__________________________________ p. CCl4
_____________ g. sodium phosphite
__________________________________ q. Fe2(C2O4)3
_____________ h. tin(IV) phosphate
__________________________________ r. HC2H3O2
_____________ i. carbon monoxide
__________________________________ s. NaH
_____________ j. barium sulfate
octohydrate
__________________________________ t. MgF2
2. Write the formulas to show the reactants and products for any FIVE of the laboratory situations
described below. In all cases, a reaction occurs. Assume that all solutions are aqueous unless otherwise
indicated. Represent substances in solution as ions if the substances are extensively ionized. Omit
formulas for any ions or molecules that are unchanged by the reaction. You need not balance the
equations.
a.
b.
c.
d.
e.
f.
g.
h.
A bar of nickel metal is placed in a solution of copper(II) sulfate.
Solid sodium hydride is added to water.
Propanone (C3H6O) is burned in air.
A solution of lead(II) nitrate is added to a solution of potassium sulfate.
A solution of ammonia is mixed with a solution of acetic acid.
Sulfur trioxide gas is bubble into water.
Excess concentrated potassium cyanide solution is added to a solution of nickel chloride.
Solid sodium acetate is added to 1.0M hydrobromic acid.
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SEMESTER 1
FINAL EXAM REVIEW
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Unit IV – Stoichiometry
I. Good things to know
 stoichiometry, chemical formula
 ionic compound  formula unit
 hydrate, anhydrous
 molecular compound  molecule
 mole, Avagadro’s Number
 formula mass, molar mass, molecular mass
 Standard Temperature and Pressure (STP)
 percent yield
 limiting reactant, excess reactant
 Law of Definite Proportions, Law of Multiple Proportions
 empirical formula, molecular formula
 mass spectrometer
 titration, standard solution, unknown solution
II. Problems
1) Calculate the number of moles of HF molecules in:
a) 0.385 g HF
b) 3.02 x 1024 molecules of HF
c) 50.0 mL of a 0.600M HF solution
d) 1.50L of HF at 0.989 atm of pressure
and 23.1oC
2) Answer the following questions using the following equation:
____ HF(g) + ____ SiO2(s)  ____ SiF4(s) + ____ H2O(l)
a) balance the equation
b) How many grams of silicon tetrafluoride will you get if you react 5.00g of silicon dioxide with
excess HF?
c) How many grams of silicon tetrafluoride will you get if you react 5.00g of silicon dioxide with
5.00g of HF? How many grams of the excess reactant will be left over?
d) If you actually produce 5.92g of silicon tetrafluoride in problem (c), what is the percent yield?
SEMESTER 1
FINAL EXAM REVIEW
3) A sample of 0.6760g of an unknown compound containing barium ions (Ba +2) is dissolved in water
and treated with an excess of Na2SO4. If the mass of the barium sulfate precipitate formed is
0.4105g, what is the percent by mass of barium in the original unknown compound?
4) A hydrocarbon was found to be 20% hydrogen by weight. If one mole of the hydrocarbon has a
mass of 30 grams, what is its molecular formula?
5) If you add 20.0 mL of 0.100 M iron(III) nitrate to 20.0 mL of 0.100 sodium hydroxide, how many
grams of precipitate will be formed? Write the net ionic equation for the reaction.
6) If it requires 25.0 L of a 0.500M KI solution to precipitate all of the lead(II) ions out of a 100.0
mL sample, what is the concentration of the lead ions?
7) What is the concentration of an unknown H2SO4 solution if it requires 156.3 mL of 1.50M NaOH
standard to titrate a 100.0 mL sample of the unknown?
8) In your lab, you titrated hydrogen peroxide with potassium permanganate:
6 H+ + 2 MnO4- + 5 H2O2  5 O2 + 2 Mn+2 + 8 H2O
If 36.44 mL of a 0.01652 M KMnO4 solution is required to completely oxidize 25.00 mL of
a H2O2 solution, calculate the molarity of the peroxide solution.
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SEMESTER 1
FINAL EXAM REVIEW
Unit V – Bonding
I. Good things to know
 ionic bonding  Coulomb’s Law
 covalent bonds  lone pair, shared pair
 stable octet
 polar vs. nonpolar, dipole moment
 sigma bond, pi bond
 formal charge, bond length  resonance
 hybrid orbitals, VSEPR model
 organic chemistry
 4 allotropes of carbon
 saturated and unsaturated hydrocarbons
 functional groups – hydroxyl, ethers, aldehydes, ketones, organic acids, esters, amines
 isomers, polymers
 aromatic compounds
II. Problems
1) Which of the following ionic compounds would you expect to have the higher melting point? Explain.
a.
NaCl vs. KBr
2) Put the following molecules in order of their C – O bond length. Explain.
CH3OH
CO2
CO3-2
5) For each of the following, give the Lewis Structure, type of hybrid orbitals used by the molecule, shape,
polarity, bond angle, and number of sigma and pi bonds:
a. PCl3
d. BrF5
c. NO3-1
b. CO2
e. PF5
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SEMESTER 1
FINAL EXAM REVIEW
6) Without using the table of bond energies, which of the following molecules would you expect to
require the most energy to break its C – C bonds? Explain your answer.
C2H6
C2H4
C2H2
7) Give the IUPAC name for the following:
CH3
CH3
CH3
CH
CH
2
CH
CH3
___________________________________________ a.
Br
CH3
___________________________________________ b.
CH2
CH
CH2
CH2
CH3
NH2
___________________________________________ c. Br
Br
Unit VI – Phases & Gas Laws
I. Good Stuff to Know
 definitions of solid, liquid, gas, fluid, condensed state
 Pressure - relationship between force and area
 barometer - atmospheric pressure, air pressure, barometric pressure
 standard temperature and pressure (STP): 1 atm = 76θ Torr = 76θ mmHg = 101.3 kPa and 273K
 manometer
 vaporization – boiling (how does it relate to vapor pressure?), evaporation
 amorphous solid- example
 crystal lattice – crystals
o
 density of solids, liquids, gases - water (most dense at 4 C)
 intermolecular forces: dipole-dipole interactions, London Dispersion Forces, hydrogen bonding, ionic
crystals, network solid. metallic crystals
 Phase Diagrams – melting, freezing, boiling, condensation, sublimation, deposition
 Phase Diagrams - triple point, critical temperature, unique properties of water
 substances that exist as gases
 Avogadro’s Law
 Kinetic Theory  use it to explain Boyle’s Law, Charles’ Law, Avogadro’s Law, Dalton’s Law,
compressibility of gases
 root-mean-square speed, mean free path
 diffusion vs. effusion
 deviations from ideal behavior, vanderWaal’s equation
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SEMESTER 1
FINAL EXAM REVIEW
 II. Multiple Choice
For questions 1-4, select from the following answers
a. metallic bonding
b. network covalent bonding
d. ionic bonding
e. London Dispersion Forces
c. hydrogen bonding
______ 1. Nonpolar substances such as methane (CH4) demonstrate this type of bonding.
______ 2. This kind of bonding is exhibited by diamond and quartz, and explains their hardness and
extremely high melting points.
______ 3. This type of bonding is only exhibited by ammonia, water, and hydrogen fluoride, and results
in these substances having unusually high melting and boiling points.
______ 4. This type of bonding results in solids that are poor conductors of heat and electricity, but which,
when melted, are good conductors of electricity.
Questions 5-8 refer to the following phase diagram:
______ 5. At this point the substance represented by the phase diagram will be solely in the solid phase at
equilibrium.
______ 6. This point represents the boiling point of the substance.
______ 7. At this point, the substance represented by the phase diagram could be undergoing sublimation.
______ 8. At this point the substance represented by the phase diagram will be solely in the liquid phase
at equilibrium.
______ 9. Which of the following lists of species is in order of increasing boiling point?
a. H2 , N2 , NH3 b. N2, NH3, H2 c. NH3, H2, N2 d. NH3, N2, H2 e. H2, NH3, N2
III. Essay/Problems
1. Explain why the boiling point of argon is -186oC, but the boiling point of neon is -246oC.
2) What is the pressure of the gas in the enclosed container?
P = 1.27 atm
243 mm
SEMESTER 1
FINAL EXAM REVIEW
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3) Consider three identical flasks filled with different gases:
Flask A: CO at 760 Torr and 0oC
Flask B: N2 at 250 Torr and 0oC
Flask C: H2 at 100 Torr and 0oC
______a. In which flask will the molecules have the greatest average kinetic energy?
______ b. In which flask will the molecules have the greatest average velocity?
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4) A gas occupies a volume of 34.2 mL at a temperature of 15.0 C and a pressure of 800.0 Torr. What
will the volume of this gas be at STP?
5) Find the formula mass of a gas which diffuses at a rate 1.16 times faster than that of sulfur dioxide gas.
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6) 40.0 mL of helium gas is collected over water at 20.0 C. If this gas exerts a pressure of 790.0 Torr,
what would the volume of the dry gas be at STP? (Remember this is a mixture of two gases: water vapor
o
and helium)
(Pwater = 17.54 mmHg at 20.0 C)
o
7) What is the density of oxygen gas collected at 21.0 C and 103.5 kPa?
8) If you collected 0.506 g of a gas which you know to be composed of 30.4% nitrogen and 69.6% oxygen,
o
and it occupied a volume of 0.134 L at a temperature of 2θ C, and a pressure of 0.986 atm, what would the
formula mass of the gas be? Give the molecular formula of the gas as well.