Chemistry 20 Final Exam Review
This review includes the questions from each of the exams you have written. Use these questions to help you study. In
addition to answering these questions you should read over your notes and make sure you understand all of the
concepts we have learned.
Unit One Exam
1. Define the following terms
b) The lab procedure should be read over and
Chemistry, Matter, Proton, Electron, Neutron
2. Explain the Scientific Method.
c) Lab safety and location of safety equipment
3. Lab safety Multiple choice. Circle the best
should be known.
d) After reading the procedure, the
-When you are finished using a Bunsen burner,
observations should be started.
you should:
4. List the easy to observe signs that a chemical
a) Leave it on for the next person to use.
change has occurred.
b) Cover the burner with an inverted beaker
5. Scientific notation
to put out the flame.
a. Write 473000000 in scientific notation.
c) Pull off the hose connecting the burner to
b. Write 0.000000323 in scientific notation.
the gas.
c. Expand 6.832 x 105 from scientific notation.
d) Turn off the gas.
6. SI Prefixes/ Unit Conversions. Use the correct
-Safety glasses:
number of significant figures. Assume all
a) Are lame and should not be worn unless
conversion factors have different numbers of
the teacher is looking
significant figures.
b) Should be worn at all times during a lab
a. Convert 293Gg to grams.
b. Convert 6.51 miles to cm (1 mile =
c) Should be worn when you are looking at
the chemicals or are within arms reach of
c. Convert 342 miles/hr to kilometres/second
7. How many significant figures does
d) Should be worn while performing a lab and
0.0006900900 have?
removed when you are done performing the
8. How many sig figs does the answer to the
following questions have?
All of the following should be done before
a. 1.35 x 2.467
starting a lab, except:
b. 98.3+5.01+23
a) Proper clothing and eye wear should be
c. 505 – 450.25
9. On the following targets, draw results that are precise but inaccurate, accurate but not precise and both precise
and accurate.
10. What are 2 examples of SI Units (Remember SI Units are not the same as SI Prefixes).
11. What is the length of the object below? Use the correct number of decimal places. Assume the ruler is in
12. Identify which chemist proposed the following model by placing the name in the corresponding blank.
13. What do the number of protons determine in an
17. How many Neutrons does Platinum-195 have?
18. Answer the questions below relating to 3216S2- .
14. What do the number of electrons determine in an
a. How many protons does the atom above have?
b. How many neutrons does the atoms above have?
15. What do the number of neutrons determine in an
c. How many electrons does the atom above have?
d. What is the charge of the atom above?
16. How many protons does Germanium have?
e. What is the mass number of the atom above?
19. Fill in the following table:
# of
# of
# of
Mass #
Protons Electrons Neutrons
Chlorine - 37
20. Magnesium has 3 naturally occurring isotopes: 24Mg (23.985042amu) with a percent abundance of 78.99%,
25Mg (24.985837amu) with a percent abundance of 10.00%, and 26Mg (25.982593amu) with a percent
abundance of 11.01%. Calculate the average atomic mass of Magnesium to 3 decimal places.
21. Boron has two naturally occurring isotopes: 10B with a mass of 10.013amu and 11B with a mass of 11.009amu.
Determine the percent abundance of each isotope.
22. What is the name of the chemist who developed
25. List 4 properties of metals.
the periodic table?
26. What is the name for the group 1 elements?
23. List the English name, symbol and Latin name for
27. Which atom has the largest atomic radius,
3 elements whose symbol comes from its Latin
according to the atomic radius trend we learned?
28. What 2 oxides are released from your vehicle
24. List the 7 diatomic elements.
when it burns fuel, and contribute to acid rain.
Unit 2 Exam
1. What are the 4 orbitals found in an atom?
a. a, b, c, d
b. s, p, d, f
c. s, p, g, f
d. s, b, d, g
e. s, c, d, g
3. Which d orbital does the following picture
4. The orbital that holds the least number of
electrons is which orbital?
a. b
b. p
c. d
2. What does the s stand for, when we are talking
about the s orbital?
a. Small
b. Safe
c. Shell
d. Sharp
e. Smelly
d. s
e. g
f. f
5. The orbital referred to in the question above
holds how many electrons?
a. 1
b. 2
c. 6
d. 10
e. 14
f. 18
6. The real orbital that holds the most number of
electrons is which orbital?
a. b
b. p
c. d
d. s
e. c
f. f
7. The orbital referred to in the question above
holds how many electrons?
a. 7
b. 10
c. 14
d. 16
e. 18
f. 20
8. Which of the following is not true about
a. They take part in bonding
b. They like to pair up
c. They have equal attraction to all elements
d. They fill the orbitals in a specific order
e. They can be transferred or shared
9. The order electrons fill orbitals is:
a. 1s 2s 3s 3g 4s 5f 6s 2p 3p 4c
b. 1s 2s 2p 3s 3p 3d 4s 4p 4d
c. 1s 2s 2p 3p 3d 4d 4f 5f 5g 6g
d. 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d
e. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
10. The four quantum numbers are n, l, Ml, Ms. The
quantum number l stands for:
a. Magnetic quantum number
b. Spin quantum number
c. Principal quantum number
d. Heissenburg quantum number
e. Angular quantum number
11. If quantum number n=2, l=1, Ml= -1, 0, +1, and
Ms= + ½ where is the electron described by these
quantum numbers located?
a. 1d orbital
b. 2p orbital
c. -1d orbital, 0d orbital, +1d orbital
d. 2s orbital
e. 2d orbital
12. Valence electrons are in which orbitals?
a. Inner most s and d orbitals
b. Outermost f and g orbitals
c. Innermost s and p orbitals
d. Outermost s, p and d orbitals
e. Outermost s and p orbitals
13. Which of the following is not true?
a. Valence electrons help to determine how an
element reacts.
b. The fewer valence electrons an atom has, the
less stable the atom is.
c. The fewer valence electrons an atom has, the
more likely it is to react
d. The more complete the valence shell is, the
more reactive the atom is.
e. A full valence is having 8 electrons.
14. Sodium has this many valence electrons:
a. 1
b. 2
c. 3
d. 4
e. 5
15. Nitrogen has 5 valence electrons. Where are
these electrons located?
a. 2 in the p orbital, 3 in the d orbital
b. 2 in the g orbital, 3 in the p orbital
c. 2 in the s orbital, 3 in the p orbital
d. 2 in the s orbital, 3 in the d orbital
e. 1 in s orbital, 3 in p orbital, 1 in d orbital.
16. When iodine has a full valance shell, it has a
charge of:
a. 0
b. -1
c. +1
d. -2
e. +2
17. When Magnesium has a full valence shell, it
usually has a charge of:
a. 0
b. -1
c. +1
d. -2
e. +2
18. Which of the following is a chlorine ion?
a. Cl2
b. Cl1c. Cl
d. ICl
e. 2Cl
19. Group 17 elements have how many valence
a. 17
b. 10
c. 7
d. 3
e. -1
20. How are Valence shells filled?
a. Electrons shared or transferred from other
b. Electrons shared or transferred with
electrons from the non-valence shells
c. Electrons gained or lost
d. Electrons are formed or disintegrated
21. Group 15 elements have this charge when their
valence shell is full.
a. -1
b. -3
c. +3
d. +1
e. +5
22. The Lewis dot diagram for Oxygen is:
23. The molecule (NH4)4Fe(CN)6 has how many
a. 11
b. 27
c. 34
d. 33
e. 8
24. The molecule H2SO4 has which kind of bond?
a. Ionic
b. Covalent
c. Both Ionic and Covalent
d. Neither Ionic or Covalent
25. The molecule Mg(C2H3O2)2 has which kind of a
a. Ionic
b. Covalent
c. Both Ionic and Covalent
d. Neither Ionic or Covalent
30. Choose an orbital that is not the s orbital and state
what the letter stands for.
31. Draw the 3 P orbitals and label their axis. The three p
orbitals are the Px, Py and Pz.
32. Write out the electron configuration for Chlorine.
33. The electron configuration [Kr] 5s2 4d10 5p5 if for
which element?
34. What is the electron configuration for Platinum.
35. What is quantum number l for the 68th electron placed
in Platinums electron configuration?
36. What are the four quantum numbers for the 19th
electron placed in vanadiums configuration?
37. Other than the transition metals, which 2 elements
are exceptions to the octet rule?
26. Magnesium and Phosphate combine to form:
a. Mg(PO4)
b. Mg3(PO4)2
c. Mg2(PO4)3
d. Mg2(PO3)3
e. Mg3(PO3)2
27. When Xenon reacts with Bromine, sometimes it
forms the molecule XeBr4. The lewis dot structure
for this molecule is shown below. What is the
molecular geometry of this molecule?
Square Pyramidal
Square Planar
T shaped
28. Which of the following does hydrogen not form
an H-bond with?
a. Nitrogen
b. Oxygen
c. Fluorine
d. Sulphur
e. Hydrogen forms an H-bond with all atoms
29. Which type of reaction is Mg +2H2O -> Mg(OH)2 +
a. Synthesis
b. Decomposition
c. Single Displacement
d. Double Displacement
e. Combustion
38. How many valence electrons does Sulphur have?
Describe which orbitals the valence electrons are
located in.
39. When sulphur has a full valence shell, what charge does
it have?
40. Draw the lewis dot structure for a Phosphorus atom.
41. Describe ionic and covalent bonding.
42. What 2 things are required for bonding to occur?
43. How many atoms are in the molecule Mg(NO3)2?
44. If Copper (II) combines with Fluorine, how many
Copper (II) atoms combine with how many Fluorine
45. Show an example of a Covalent Bond.
46. Draw the lewis dot structure for a hydrogen molecule.
47. Draw the Structural formula for a hydrogen
48. Draw the lewis dot structure for the compound that
is formed between Calcium and Chlorine.
49. What does VSEPR stand for?
50. Use VSEPR to determine the shape/geometry of the
Hydrogen molecule in question 47 and 48 above.
51. Draw the lewis structure for the compound that is
formed between Nitrogen and Iodine. Use VSEPR to
determine the electron pair and molecular
geometry. Draw the 3D structure of this molecule.
52. What is a polymer?
Linus Pauling was an American Chemist who developed the concept that electronegativities can help to determine the
iconicity of a bond (that is how ionic a bond is). Pauling’s scale is what we use to determine whether a bong is ionic or
covalent or polar or non polar. The way the Pauling scale is used is by subtracting the electronegativities of the 2 atoms
involved in a bond. If the difference between the electronegativites is 1.7 or greater, the bond is ionic. If The
electronegativity difference is between 0.4 and 1.7 then the bond is polar covalent and if the electronegativity
difference is less than 0.4 then the bond is non-polar covalent.
53. Use Paulings scale to determine if a bond formed
59. In the reaction below label the products, reactants, and
between Hydrogen and Chlorine is Ionic, Polar
catalyst. Also state what the catalyst is in the reaction.
Covalent, or Non-polar Covalent.
2CH3OH + 3O2
2 CO2 + 4 H2O
54. Draw the Dipole for the bond between Hydrogen and
60. Circle the true statement below:
The reaction that occurs when Alka-Seltzer is placed in
55. What is the definition of Van der Waals Forces?
water is between sodium bicarbonate and citric acid.
56. What is the charge of Lead in the molecule Pb(CrO4)2?
The reaction that occurs when Alka-Seltzer is placed in
57. What type of reaction is the following?
water is between sodium bicarbonate and water
Pb(NO3)2 +2KI -> PbI2 +2KNO3
61. State 5 things you know about either Water Pollution or
58. What are 4 ways to change the rate of a reaction?
Air Pollution
Unit Exam 3
1. Determine the number of atoms in the following
chemical formulas:
a. KNO3
b. (NH4)3PO4
2. Determine the number of Oxygen atoms in Al2(SO4)3.
3. Determine the charge on iron in the compound below:
4. Balance the following equations:
a. ___Li + ___Cl2  ___LiCl
b. ___C3H8 + ___O2 ___CO2 +___H2O
c. ___Pb(NO3)2 + ___NaI  ___NaNO3 + ___PbI2
5. For the following reactions, state whether it is
decomposition, synthesis, single displacement of
double displacement.
10. Name the following compounds:
a. KCl
b. (NH4)3PO4
c. CO2
d. HBr
e. Fe(NO2)2
f. PCl5
g. H2SO3
a. H2 + Cl2  2HCl
b. 2HCl + Zn  H2 + ZnCl2
c. H2CO3H2O + CO2
Write the balanced net ionic equations for the following
a. NaOH(aq) + HCl(aq)  H2O + NaCl(aq)
b. Mg(s) + 2AgNO3(aq)  Ag(s) + Mg(NO3)2(aq)
What is an exothermic reaction?
What is an endothermic reaction?
State whether the following reactions are exothermic or
a. Baking bread
b. N2 + O2 + energy  2NO
11. Write the chemical formula for the following:
a. Hydrofluoric acid
e. Zinc Nitrate
b. Lead (IV) Chromate
f. Diphosphorus pentoxide
c. Boric acid
g. 2- butene
d. Magnesium chloride
h. 3- ethyl hexane
12. Balance the following reactions. You do not need to include states or energy terms.
a. Sulfur trioxide and water yields sulphuric acid
b. Calcium hydroxide and phosphoric acid yield calcium phosphate and water.
13. Balance the following word equation. Include states (s, l, g, or aq) and an energy term:
Zinc metal reacts with hydrochloric acid to produce hydrogen gas and zinc chloride solution. The container gets
14. Predict the products of the reaction below. Be sure so write the balanced equation below.
Aluminum + Iron (III) Oxide 
Unit 4 Exam
1. Complete the diagram below by filling in the boxes with the corresponding words.
2. Calculate the number of moles in a sample of
Carbon Dioxide containing 5.98 x 1016 molecules?
3. If you have a sample of sodium atoms containing
0.00267 moles of sodium atoms, how many atoms
are in the sample?
4. If you have 4 moles of carbon tetrachloride, how
many moles of chlorine atoms are in the sample?
5. What is the molar mass of rhodium?
6. What is the molar mass of Ca(NO3)2?
7. How many moles of Carbon are in a 13.0 g sample?
8. What is the mass of a sample of Barium
hypochlorite that contains 1.03 moles of Barium
9. How many molecules of CF4 are in 2.98g of CF4?
10. What does STP stand for?
11. What is the volume of 6.25 moles of oxygen gas
12. What is the mass of 4.98L of hydrogen gas @ STP?
13. What is the percent composition of CH3COOH?
C= ________ H=________ O=________
14. Determine the empirical formula that corresponds
with 24.7% Potassium, 34.8% Manganese and
40.5% Oxygen.
15. For the equation below, how many moles of Zinc
Chloride would be produced if 6.38 moles of zinc
were reacted with excess hydrochloric acid?
Zn + HCl  ZnCl2 + H2
16. For the equation below, how many litres of carbon
dioxide would be produced if 1.38g of propane gas
were reacted with excess oxygen?
___C3H8 + ___O2  ___CO2 + ___H2O
Answer questions 17-19 according to the reaction below:
5.0g of silver (I) nitrate react with 5.0g barium chloride
___AgNO3 + ___BaCl2  ___AgCl + ___Ba(NO3)2
For the reaction in question 22, which of the reactants would be limiting and which one would be in excess?
What mass of excess reagent would be left over?
What mass of Silver Chloride was produced in the reaction above?
What volume of hydrogen @ STP is produced from the reaction of 50.0g of Mg and the equivalent of 75g of HCl?
Unit 5 Exam
Using the graph on the left answer the following
1. At what temperatures is the substance
(represented in the graph above) a gas?
2. At what temperatures is the substance
(represented in the graph above) a liquid?
3. At what temperatures is the substance
(represented in the graph above) a solid?
4. What is the substance`s melting point?
5. What is the substance`s freezing point?
6. What is the substances boiling point?
7. At 10oC what state is the substance in?
8. For the phase diagram above, label the area of the
graph where the substance is solid, liquid and
9. Explain what a triple point is.
10. For the phase diagram on the left, label the triple
point. At what temperature and pressure does this
11. What is a critical point?
12. For the phase diagram above what is the highest
temperature that a liquid can exist at?
13. What is the normal boiling point of this substance?
14. What is the normal melting point?
15. For the substance graphed above, circle the state
that is more dense. Solid
16. At 1.5atm and 100oC what state would this
substance be in?
17. At 100oC, if you started at 0.5 atm and increased
the pressure to 1.0atm which of the following
changes of state would occur? (circle)
stay the same
18. What is the molarity of a solution that of 4.28L that contains 3.86 moles of hydrochloric acid?
19. What is the molality of a solution of 3.28 moles of NaCl mixed in 672g of water?
20. State whether the following compounds are soluble or insoluble.
a. Barium sulfate
b. Aluminum hydroxide
21. Will a precipitate form in a reaction of silver nitrate with potassium carbonate?
22. Write the net ionic equation for the reaction of sodium hydroxide with calcium nitrate.