SCH 3U End of Year Review
Units 1and 2: Matter and Chemical Bonding (Chapters 1-4)
1. Define the following terms:
Matter, Mass, Weight, Energy, Kinetic Energy, Potential Energy, Law of Conservation of Mass, Law of Conservation of Energy, Solid, Liquid, Gas, Physical Properties, Chemical Properties, Density, Chemical Reaction,
Quantitative Property, Qualitative Property, Physical Change, Chemical Change, Pure Substance, Element, Compound, Mixture, Solution, Mixture,
Heterogeneous Mixture, Homogeneous Mixture, accuracy, precision, atomic number, mass number, isotope
2. Compare Dalton’s Atomic Theory to the Modern Atomic Theory 3. Complete the Practice Problem #1 on p. 19
4. Draw orbital diagrams and Lewis dot diagrams for the first 20 elements of the periodic table
5. Define ionization energy, electron affinity, and atomic radius, and explain why each varies across the elements of the periodic table.
6. Explain the differences between ionic and covalent compounds in terms of the bonding and their properties.
7. Describe how bonding occurs in metals.
8. Define electronegativity and use it to predict whether a compound is ionic or covalent. (try p. 63 #10)
9. Draw Lewis structures for ionic and covalent compounds. (try p. 63 #4, p. 75 #12)
10. Use the water molecule to explain what a polar covalent bond is.
11. Write the molecular formulae and the names of the compounds formed in each box:
HCO3- ClO- ClO3- OH- CrO42- NO3- NO2- C2O42- MnO4- PO43- SO42- SO3
2-H+
Li+
Na+
Ca2+
Ba2+
Mg2+
Fe2+
Fe3+
Zn2+
Al3+
12. Write the molecular formulae for the following: a. Sodium hypochlorite
b. Ammonium dichromate c. Copper (II) Nitrate d. Calcium phosphate e. Lithium sulfite f. Tin (II) fluoride
g. Ammonium carbonate h. Iron (III) Chloride i. Aluminum Sulfate
13. Write a balanced chemical equation to represent each of the following chemical equations. Classify each reaction as a synthesis, decomposition, single displacement, or double displacement reaction.
a. aluminum + oxygen aluminum oxide
b. aluminum + hydrogen sulfate aluminum sulfate + hydrogen
c. potassium bromide + aluminum nitrate potassium nitrate + aluminum bromide
d. calcium + aluminum chloride calcium chloride + aluminum e. nitrogen + hydrogen ammonia
f. butane (C4H4) + oxygen carbon dioxide + water
g. sodium + iodine sodium iodide
h. potassium hydroxide potassium oxide + water
i. magnesium + water magnesium hydroxide + hydrogen
j. sodium carbonate + calcium hydroxide sodium hydroxide + calcium carbonate
k. calcium oxide + water calcium hydroxide
l. potassium carbonate + barium chloride potassium chloride + barium carbonate
Unit 3: Chemical Quantities (Chapters 5-6-7)
1. A student has worked out the details for an experiment that will use 0.115 mol of Ca3(PO4)2 as a starting material. How many grams of Ca3(PO4)2 should be
weighed out?
1. How many grams of sodium carbonate, Na2CO3 , must be taken to obtain
0.125 mol of Na2CO3?
2. A sample of 45.8 g of H2SO4 contains how many moles of H2SO4?
3. If an analysis of a compound shows that for every 3.75 mol of Cl there are 11.25 mol of O, what is the empirical formula?
4. A sample of carbon reacts with a 9.2-gram sample of hydrogen to form a compound that has a mass of 120.0 grams. What is the empirical formula of this compound?
5. 18.48 g of carbon and 58.52 g of fluorine are obtained from the analysis of a molecular compound. Determine the empirical formula of the compound. 6. Sulfur combines chemically with oxygen to produce an oxide of sulfur.
7. A sample of an unknown compound with a mass of 2.571 g was found to contain 1.102 g of carbon and 1.469 g of oxygen. What is its empirical formula?
8. One of the oxides of iron, “black iron oxide,” occurs naturally in the mineral magnetite. When a 2.448 g sample was analyzed, it was found to have 1.771 g of Fe and 0.6766 g of O. Calculate the empirical formula of this compound. 9. A 2.012 g sample of a compound was found to consist of 0.5219 g of nitrogen
and 1.4901 g of oxygen. Calculate its empirical formula.
10. What is the percentage composition of an ore that contains 1.85 g of aluminum and 1.65 g of oxygen?
11. A compound consisting of aluminum and chlorine weighs 3.60 g. What is the weight of chlorine in the compound? What is the percent composition of the compound?
12. Calculate the percent composition of C6H12O6.
13. Calculate the percent composition of (NH4)2CO3.
14. A sample of zinc sulfide having a mass of 3.04 g was analyzed and found to contain 1.00 g of sulfur, and the remainder zinc. If a prospector discovered 1000 kg of pure zinc sulfide, what would be the maximum amount of zinc that could be recovered?
15. How many moles of oxygen atoms are combined with 4.20 moles of chlorine atoms in Cl7O2?
16. How many molecules of Na2SO4 are there in 2.7 g of Na2SO4?
17. How many moles of nitrogen atoms are combined with 8.60 moles of oxygen atoms in dinitrogen pentoxide, N2O5?
18. How many grams of Cl are needed to combine with 24.4 g of Si to make silicon tetrachloride?
19. How many grams of iron are needed to combine with 25.5 g or O to make Fe2O3?
20. Assume that 2.40 moles of oxygen are reacted with 2.46 moles of
magnesium to produce magnesium oxide, MgO. Determine which reactant is in excess and by what amount, and calculate the number of moles of product formed.
21. Assume that 13.1 g of potassium are reacted with 18.0 g of oxygen to
produce potassium oxide, K2O. Determine which reactant is in excess and by
what amount, in m0oles, and calculate the number of gram of products formed.
22. A 5.00 g sample of mercury (II) oxide, a red powder, is heated in a open 12.00 g test tube. After prolonged heating, the mass of the test tube and its contents is 16.63 g. A silver liquid remains.
i. Calculate the % composition of the mercury (II) oxide. ii.What mass of mercury will combine exactly with 500 g of
oxygen to form mercury (II) oxide?
iii. 2.00 g of mercury and 0.120 g of oxygen are heated in a closed container until the reaction is complete. What
substances remain? Determine the masses of each. 23. A chemist wishes to determine the composition of copper (I) sulfide. She
a. What is the mass and identity of the black powder?
b. Calculate the % of copper in the compound formed in the reaction. c. Calculate the mass of sulfur burned off in the fume hood.
d. How many grams of sulfur is needed to react with 60 g of copper to make the copper-sulfur compound?
24. 3.15 g of magnesium combines with 1.98 g of oxygen to make magnesium oxide.
a. Calculate the empirical formula.
b. Calculate the % composition of magnesium oxide.
c. How many grams of oxygen are consumed when 1000 g of magnesium is burned?
d. If 4.0 g of magnesium and 1.5 g of oxygen are heated in a closed container, how much magnesium oxide forms?
25. When 2.70 g of aluminum is completely burned, 5.10 g of aluminum oxide forms.
a. Calculate the percent composition of aluminum oxide.
b. Calculate the mass of oxygen needed to burn 480 g of aluminum.
c. Calculate the mass of aluminum oxide that forms when 32 g of aluminum are completely oxidized.
d. 6.0 g of aluminum and 5.7 g of oxygen are heated together in a closed container until reaction is complete. What substances will remain? Calculate their original masses.
26. When 400 g of water is broken down into its constituent elements, 44.4 g of hydrogen forms.
a. Determine the percent composition of water.
b. What mass of hydrogen will combine exactly with 15 g of oxygen to make water?
Unit 3: Solutions and Their Concentrations (Chapters 8-9-10)
1. Define the following terms: Homogeneous mixture, Heterogeneous mixture, Miscibility, Solute, Solvent, Ionic Compound, Covalent Compound, Saturated Mixture, Unsaturated Mixture, Supersaturated Mixture, Concentration,
Molarity, Molality, Heat of Solution, Precipitate, Dissociation, Net Ionic Equation, spectator ion
2. Explain how ionic substances dissolve in water. 3. Explain how covalent substances dissolve in water.
4. What is a dipole? Give three examples of molecules that possess a dipole. 5. What is a dipole-dipole attraction?
6. What is hydrogen bonding and how does it occur? Give an example of a compound that undergoes hydrogen bonding.
7. How do intermolecular forces affect the boiling points of compounds? Which intermolecular force that we learned about is the strongest? Weakest? 8. Explain the effects of temperature on the solubility of a compound. 9. Explain some of the factors that affect solubility.
10. What is the molarity of a solution made by adding 295.6 g of potassium sulfide, K2S, to sufficient water to make 10.00 L of solution.
11. Calculate the molality of a solution made by dissolving 31.5 g of NaNO3 in
12. How many grams of NaCl are present in a 1.21 M solution made by using 15.0 L of water?
13. What is the molarity of a solution made by adding 16.5 g of sodium hydroxide, NaOH, to sufficient water to make 2.7 L of solution.
14. Calculate the molality of a solution made by dissolving 13.4 g of (NH4)2SO3 in
567 g of water.
15. How many grams of KCl are present in a 0.83 M solution made by using 10.0 L of water?
16. How does the solubility of the products of a reaction affect the net ionic equation?
17. For each of the following reactions:
i) Predict the products of the reaction.
ii) Determine whether or not a precipitate forms.
iii) Write the net ionic equation for the reaction and identify the spectator ions.
a. MgSO4(aq) + CaBr2(aq)
b. FeSO4(aq) + (NH4)2S(aq)
c. CuSO4(aq) + FeCl3(aq)
d. HNO3(aq) + CaBr2(aq)
e. Al2(SO4)3(aq) +CaCl2(aq)
18. Define acid and base according to both Arrhenius and Bronsted-Lowry. 19. List 4 properties of acids and 4 properties of bases.
20. What is the difference between a strong acid and a weak acid? 21. Define neutralization.
22. Identify conjugate acid-base pairs. (Try p. 380 # 10)
23. What is an acid-base indicator? Give three examples of indicators and explain how you would use them.
24. Sketch a titration curve for the titration of H2SO4 with NaOH.
25. Write the ionization equations for H2SO4 and HNO3 dissolving in water.
26. How much sodium carbonate would you have to weigh out to make up 500 mL of a 0.001 M solution?
27. Find the mass of sodium hydroxide in 100 mL of a 0.01 M solution of sodium hydroxide.
28. What volume of 0.1 M NaOH will just neutralize 25 mL of 0.08 M H2SO4?
29. Find the molarity of an HCl solution if 24 mL reacts completely with 0.207 g of sodium carbonate.
30. What is the concentration of NaOH, given that 25 mL of it was neutralized by titration with 10 mL of 0.15 M HCl?
31. 20 mL of 0.1 M NaOH was neutralized by 25 mL of HCl. What was the concentration of the acid?
32. 20 mL of 0.1 M NaOH was neutralized by 25 mL of H2SO4. What was the
concentration of the acid?
33. Calculate the volume of 0.50 M sodium hydroxide solution required to neutralize 150 .0 mL of 1.5 M nitric acid.
34. Calculate the volume of 0.50 M sodium hydroxide solution required to neutralize 10.0 mL of 0.25 M sulfuric acid.
35. Calculate the molarity of an acetic acid (CH3COOH) solution if 34.57 mL of
this solution are needed to neutralize 25.17 mL of 0.1025 M sodium hydroxide. CH3COOH(aq) + NaOH(aq) NaCH3COO(aq) + H2O(aq)
1. 2. 3. 4.
How much water would you add to obtain a solution of the required concentration
?
34. How can we prepare 200 mL of 0.0500 M K2Cr2O7 from 0.300 M K2Cr2O7?
35. What would be the concentration (in mol/L) of a solution made by adding 25.0 mL of water to 10.0 mL of a 0.25 M solution?
Unit 4: Gases and Atmospheric Chemistry (Chapters 11-12)
1. State the four postulates of the kinetic molecular theory of gases. 2. What is pressure? How is pressure calculated?
3. Use the concepts of volume and temperature to explain why bike tires can go flat over the winter?
4. You are driving up to the mountains to go skiing. If you had a balloon in the car with you, it would get bigger as you travel higher in altitude. Explain why this would happen.
5. Explain the concept of gas pressure in terms of the kinetic molecular theory. 6. Use the concepts of pressure and temperature to explain why it is dangerous
to throw an aerosol can into a fire. Which gas law is being illustrated? 7. A balloon can burst when too much air is added. Describe what happens to
air pressure, volume, temperature, number of moles, and the balloon itself as it inflates and finally bursts.
8. Complete the following tables:
a. Gay-Lussac’s Law Problems:
Initial Conditions Final Conditions
P1 T1 P2 T2
95 kPa -15 C 27 C
100 C 2.00 atm -170 C
b. Boyle’s Law Problems:
Initial Conditions Final Conditions
P1 V1 P2 V2
95 kPa 205 mL 0.659 L
1.00 atm 5.00 L 2.00 atm
c. Charles’s Law Problems:
Initial Conditions Final Conditions
V1 T1 V2 T2
4.20 L -15 C -170 C
15.6 L 100 C 10.8 L
9. 257 mL of oxygen in a balloon goes from 17 C to 42 C from being out in the sun. The pressure in the balloon does not change as the temperature
10. What volume of air at standard pressure gets packed into a 30.0 L SCUBA tank at the same temperature and 5 atm?
11. An enormous (57, 400 cubic metre) expandable helium balloon at 22 C is heated up by a fire under it and by the action of the sun heating its dark plastic covering. To what temperature must the air in the balloon be heated to in order to fill out to 66, 500 cubic metres so that it can lift its cargo? 12. Exactly one mole of oxygen gas was collected in the laboratory at a
temperature of 24.0 C and a pressure of exactly 100.0 kPa. The volume occupied by the mole of gas at this temperature and pressure was measured as 24.686 L. What is the value of R?
13. Calculate the pressure of 2.35 g of nitrogen gas at 18 C and occupying 4.87 L.
14. What will be the volume of 1.2 g of oxygen gas at standard temperature and 790 mmHg?
15. What is the temperature of 3.2 moles of helium gas that occupies 3.3 L at 1.10 atm.
16. How many moles of argon are there in a 18.4 L sample of gas at 760 mmHg and 29 C?
17. What is the mass of 15 L of chlorine gas at STP?
18. What is the mass of 25 L of fluorine gas at 2.85 atm and 450 C?
19. How many litres of ammonia at STP are produced when 10 g of hydrogen is combined with nitrogen?
20. How many mililitres of hydrogen at 0 C and 1400 mmHg are produced if 15 g of magnesium reacts with sulfuric acid?
21. A 250 Kg tank of liquid butane (C4H10) burns to produce carbon dioxide at
120 C and water. What volume of carbon dioxide is produced at 1 atm? 22. How many litres of product at 960 mmHg and 0 C are produced by the
burning of three litres of acetylene (C2H2) at 5 atm and 20 C?
23. If 0.515 g of magnesium is added to HCl, it makes hydrogen gas and magnesium chloride. The hydrogen is collected at 23 C and 735 mmHg. What is the volume of hydrogen produced?
24. Isopropyl alcohol, C3H7OH, makes an excellent fuel for cars. What volume of
