FINAL EXAM REVIEW
Chemistry 12 College – SCH 4C1
UNIT 1:
1. Explain the difference between an inference and an observation.
2. (A) Name the 3 subatomic particles?
(B) What are their charges?
(C) What are their masses?
(D) Where are they found within an atom?
3. What is an isotope?
4. Why are the atomic masses on the periodic table not whole numbers?
5. What is the difference between mass number and atomic number?
6. Draw Bohr diagrams for each of the following. Please take notice of whether the
item is an atom or an ion – it will matter for your diagram.
 Calcium Atom
 Potassium Ion
 Magnesium Atom
 Nitrogen Ion
 Neon Atom
 Beryllium Ion
 Sulfur Atom
 Chlorine Ion
7. (A) What is an ion?
(B) What is the charge of a phosphide ion? A magnesium ion? A fluorine ion?
8. Draw the Lewis symbol for carbon, argon, potassium.
9. (A) How can a flame test be used to identify a metal?
(B) Explain what happens to an electron when a compound is placed in a flame to
produce colour. (Think of your Flame Test Experiment!)
(C) When an electron is in the ground state, is it likely to gove off any light energy?
Explain
10. What column are the halogens in? The noble gases? The alkali metals? The alkaline
earth metals?
11. How does an ionic compound differ from a covalent compound?
12. What is an ionic compound made up of?
13. List the polyatomic ions (don’t forget their charges) that you must know – there are
eight of them!
FINAL EXAM REVIEW
Chemistry 12 College – SCH 4C1
14. Give the formula for the following:
 Sodium chloride
 Potassium hydroxide
 Magnesium carbonate
 Tin (IV) sulfide
 Aluminum sulfate
 Diphosphorus pentoxide
 Nitrogen dioxide







Silicon tetrachloride
Iron (II) oxide
Lithium nitrate
Ammonium oxide
Dinitrogen tetraoxide
Calcium fluoride
Lead (II) bromide
15. Name the following compounds:
PbO2
___________________
CuS
CoI3
___________________
MgSO4
Li2O
___________________
Ca(OH)2
(NH4)3PO4___________________
F2
AlPO4
___________________
N2O3
PbS
___________________
N2S3
CF4
___________________
SF4
Fe2(SO4)3 ___________________
SnCl4
NH3
Mg3(PO4)2
___________________
16. What is a covalent molecule made up of?
17. How does a covalent molecule differ from an ionic molecule? (Think of the
electrons!)
18. What is the octet rule?
19. What does it mean if a molecule is polar?
20. Give an example of a molecule that has a pure covalent bond.
21. (A) If the difference in electronegativity is greater than 1.7 the bond is________.
(B) If the difference in electronegativity is 0 the bond is _______________.
(C) If the difference in electronegativity is between 0.1 and 1.6 the bond is_____.
22. (A) What happens to the electronegativity of compounds as you move across a
period?
(B) What happens to the electronegativity of compounds as you move down a group?
23. Name the five types of reactions discussed in class.
FINAL EXAM REVIEW
Chemistry 12 College – SCH 4C1
24. Are you able to balance equations?
25. Do you know how to do net ionic equations? (in general)
26. Be able to use the solubility table and the activity series!
27. For each of the following reactions, determine what the products of each reaction will
be. When you have predicted the products, balance the equation and use a table of
solubility products to determine which of the products (if any) will precipitate.
Assume all reactions take place in water.
(A)
____ Ca(OH)2 + ____ HF 
(B)
____ Pb(NO3)2 + ____ K2CO3
(C)
____ NaNO3 + ____ H2SO4 
(D)
____ Cu(OH)2 + ____ H3PO4 
(E)
____ AgNO3 + ____ Na2CO3 
(F)
____ Zn + ____ H2CO3 
(G)
____ Pb(OH)2 + ____ Hg2S 
FINAL EXAM REVIEW
Chemistry 12 College – SCH 4C1
UNIT 2: QUANTITATIVE CHEMISTRY
1. (A) What is a mole?
(B) What is Avogadro’s number?
2. What is the percent composition of NH4NO3? Cu(OH)2?
3. Find the atomic mass of the following:
 Copper
 Copper (II) chloride
 Calcium hydroxide
4. Find the number of moles in the following:

30 grams of H3PO4

200 grams of FeCl3

25 grams of HF

564 grams of copper
5. How many molecules are there in 2 moles of H2O?
6. How many molecules are there in 1.5 moles of N2?
7. How many molecules are there in 10 grams of Fe?
8. How many molecules are there in 10 grams of H2O?
STOICHIOMETRY
9. In the following problems, calculate how much of the indicated product is made.
Show all your work.
(A)
LiOH + HBr  LiBr + H2O
If you start with ten grams of lithium hydroxide, how many grams of lithium
bromide will be produced?
(B)
C2H4 + 3 O2  2 CO2 + 2 H2O
If you start with 45 grams of ethylene (C2H4), how many grams of carbon dioxide
will be produced?
(C)
Mg + 2 NaF  MgF2 + 2 Na
If you start with 5.5 grams of lithium chloride, how many grams of calcium
chloride will be produced?
Organic Chemistry Unit Review
Chemistry 12 College – SCH 4C1
10. If calculations show you should produce 12 g of calcium oxide, and after doing the
experiment you produce 9.5 g, what is your percent yield?
11. Consider the following reaction:
CaO (s) + 2 HCl (aq)  CaCl2 (aq) + H2O (l)
(a) How many grams of CaCl2 would be produced from 80 g of HCl?
(b) If 105 g of calcium chloride is formed, what is the percent yield?
PERCENT COMPOSITION & DETERMINING FORMULAS
12. A compound is found to contain 38.72% carbon, 9.72% hydrogen and 51.56%
oxygen. Determine the empirical formula of the compound.
13. What is the empirical formula of a compound containing 26.52% Cr, 24.52% S and
48.96% O?
MEASURING CONCENTRATION
14. What is the concentration of a solution made by dissolving 2 moles of NaCl in 250
mL of water?
15. What is the concentration of a solution made by dissolving 10 g of CaCO3 in 500 mL
of water?
16. If 25 mL of a sodium bromide solution that has a concentration of 1.2 mol/L is
diluted to 100 mL with water, what is the new concentration of the solution?
17. What volume of a solution having a concentration of 2.4 mol/L must a chemist take
in order to prepare 500 mL of a new solution having a concentration of 1.6 mol/L?
Organic Chemistry Unit Review
Chemistry 12 College – SCH 4C1
UNIT 3: Environmental Chemistry
1. What are the three types of contaminants found in water? Give an example of each.
2. What is the purpose of a water softener?
3. Where do we get our drinking water from?
4. What is acid rain?
5. What happened at Walkerton?
6. List three properties of acids.
7. List three properties of bases.
8. What is the range of the pH scale? What values represent acids? Bases? Neutral?
9. Calculate the pH of a solution of HCl that has a concentration of 0.0004 mol/L.
10. Calculate the pH of a solution of NaOH that has a concentration of 3.8 x10-8 mol/L.
11. Calculate the pH of a solution of H3PO4 that has a concentration of 0.00046 mol/L .
12. Calculate the pH of solution of Ca(OH)2 that has a concentration of 5.4 x10-4 mol/L.
13. Calculate the pH of a solution made by dissolving 3.45 g of HF in 1500 mL of water.
14. Calculate the pH of a solution made by dissolving 6.4 g of KOH in 900 mL of water.
Organic Chemistry Unit Review
Chemistry 12 College – SCH 4C1
UNIT 4: ORGANIC CHEMISTRY
Drawing Organic Structures
1. Draw the following hydrocarbons based on their names.
 Methane
 2-Butanol
 Ethyne
 Ethoxypropane
 Butane
 Propoxyhexane
 3-Octene
 3-Chlorohexane
 2-Pentyne
 3-Methylhexane
 1,4–Heptene
 2,3-Dimethylpropane
 2,5–Nonyne
 2.4-Dichloropentane
2. Name the following organic compounds.
(A)
(F)
(B)
(G)
(C)
(H)
(I)
(D)
(E)
(J)
FINAL EXAM REVIEW
Chemistry 12 College – SCH 4C1
Reactions With Organic Compounds
3. Draw the reactants and draw & name the products of each of the following reactions.
Balance the equations if necessary.
REACTION 1
2-Pentene
+

H2O
REACTION 2
2,5-Octene
+
HCl

REACTION 3
2-Butene
+
H2

SHORT ANSWER
4. What is the difference between an alkane, and alkene and an alkyne?
5. What is the rule for assigning a number to a functional group or special bond within
the chain of a hydrocarbon?
6. What functional group is added to a hydrocarbon to make it an alcohol?
7. What is meant by the term “hydrogenation”?
8. What structural feature do all ethers have in common?