SCH 3U EXAM REVIEW
This is not a complete review of all of the topics covered.
This assignment is to get you started on your review. It does not cover all of the topics that were covered within each unit. Answer each section on another piece of paper. For questions involving calculations, be sure to show all of your work that lead you to the final answer. Be aware of your form, significant figures, units and therefore statements.
It is NOT expected that you would do all of the following questions. This list simply provides questions that you SHOULD be able to do. As you review your notes, labs, quizzes and tests, you will find concepts or problem types that you may be a little uncertain of. The list below from the text provides additional practice questions on any concepts you require. Answers to many of the questions are provided in the unit answer keys under the Exam Review tab on www.manningsscience.com.
Good luck on the exam! If you have done the work required during the course, a thorough review should put you in an excellent position to do well on the exam.
Units of Study (refer to the individual unit expectations when studying)
Unit / Concepts / Text Reference1: Matter and Chemical Bonding / · Structure of the atom & ion (5 ways to represent an atom & 4 ways to represent an ion)
o Calculating subatomic particles
· electron configurations (s, p, d, f blocks)
· Periodic table trends (IE, EA, electronegativity & atomic radius) – explain using keywords (nuclear charge, electron shielding & effective nuclear charge
· Intramolecular forces: Chemical bonds (ionic, covalent & metallic) both polar & non-polar (show dipoles and partial charges)
· Lewis Structures for covalent & ionic compounds
· Structural diagrams for covalent compounds (using formal charge)
· VSEPR & shapes
· Intermolecular forces
· Physical & Chemical properties of ionic & covalent compounds
· Naming compounds & writing chemical formulas
o Molecular (use prefixes)
o Ionic – multivalent, polyatomic & binary
§ Per, hypo, ite, ate
o Acids (binary & oxyacids)– ic & ous acids
· Calculating average atomic mass using isotope abundances / pg. 21 #3, 10
pg. 30 #13
pg. 40 #2, 9, 11, 13
pg. 48-49, skip #12 & 19
pg. 63 #3, 5, 12, 14,
pg. 75 #4, 5, 6, 14
Self-Assessment pg. 92-93 1-25
2: Chemical Reactions / · Balancing chemical equations (with states)
· Word equations
· Identifying types of reaction
· Predicting products of reactions
· Nuclear reactions/half-life (radioactive decay)
· Total and net ionic equations / pg. 121 #1, 2, 10, 11, 14
pg. 136 #5, 15
pg. 145 #5
pg. 158-159 skip #13, 14, 21, 23, 24, 25
pg. 180 #7, 15
pg. 204-205 #3, 4, 5, 6, 11, 15
pg. 414
3: Quantities in Chemical Reactions / · Significant digits
· Defining the mole
· Calculations involving Avogadro’s number (particles, atoms, ions…)
· Calculations involving the mole/molar mass
· % composition
· Percentage Yield (theoretical vs. actual)
· Finding empirical formula/molecular formula
· Balancing equations & stoichiometry
· Limiting reagent, excess reagent / pg. 232 #7, 8,
pg. 243 #4, 8, 11
pg. 254-255 skip #15, 19, 23, 24, 25
pg. 267 #6, 13
pg. 279 #3, 7, 10
pg. 292-293 skip #11, 13, 24, 25
pg. 305 #1, 2, 3, 5, 10
pg. 321 #5
pg. 334-335 #1-25
pg. 414 #2, 3, 4
4: Solutions and Solubility / · Types of solutions (saturated, unsaturated & supersaturated
· Solubility (solubility curves) – relating to types of solutions
· Concentration – molarity
· Dilution
· Methods of making a solution (from solid & stock solution)
· solution stoichiometry
· acids & bases; strong & weak, pH
· titrations – calculations involving them / pg. 370 #2, 7, 9,
pg. 390 #3, 4, 5, 6, 9
pg. 421` #1, 2, 3, 6
pg. 463 #1, 9, 10, 11, 12
pg. 470 #2, 3, 4, 5, 7, 8, 10, 12
pg. 492-493 #2, 34, 8, 10, 23
Format for the Exam:
You will be given a formula sheet and periodic table. The exam will be composed of:
Part A: Multiple Choice ~ 25 marks
Part B: Short Answer - Inquiry ~ 40 marks (problems/calculations, problem solving/explaining,
reading graphs)
Part C: Short Answer – Communication ~25 marks (comparisons, diagrams, communication knowledge using keywords, writing chemical formula/naming)
PRACTICE:
Part A: Definitions Define the following words and provide examples when appropriate. Diagrams may also be helpful.
Electronegativity / Chemical change / SolubilityEnergy level / Diatomic molecule / Miscible
Electrolyte / Titration / Precipitate
Periodic trend / Dilution / radioisotope
Valence electron / Concentration / Neutralization
Stable octet / Aqueous solution / Endpoint
Electron shielding / oxyacid / Molar volume
Isotope / Law of conservation of mass / Equivalence point
Spectator ions / Effective nuclear charge
Part B: Comparison Compare the following sets of terms. (similarity/difference/ examples for each)
atom vs. ion vs. polyatomic ion / electron affinity vs. electronegativityatomic mass vs. molar mass / Isotope vs radioisotope
alpha vs. beta vs. gamma radiation / first vs. second ionization energy
relative atomic mass vs. mass number / saturated vs. supersaturated
electron affinity vs. ionization energy / actual and theoretical yield
London (dispersion) forces vs. dipole-dipole vs. hydrogen bonding / strong vs. weak acid (strong/weak base)
polar and non-polar molecules / Empirical formula vs molecular formula
anhydrous vs. hydrate / Structural diagram vs Lewis structure
ionic compounds vs. molecular compounds / Intramolecular vs intermolecular
Part C: Diagrams
1. Represent each atom/ion in as many ways as possible (eg: Bohr-Rutherford, standard atomic notation): potassium, chloride, aluminum and sulfide.
2. Compare the atomic radius, ionic radius and the 1st and 2nd ionization energies for the following; Mg, Ne, Ca 2+, F 1- - use keywords from class to explain them
3. Write out the electron configuration for silicon, and zinc. Write the short hand (noble gas configuration) for iron, and strontium.
4. Draw the Lewis structure and structural diagrams for the following molecules; Cl2, HCl, HCN, C2H2, C2H4, NH3, CH4, CH2Cl2 OF2, H2S, BF3
-indicate whether the structures in the above list are;
-polar or non-polar (add dipole & partial charges if polar), and predict the shape of each molecule, and predict the physical properties of each substance based upon their intermolecular forces
Part D: Problems
Problem Type: Molecular Mass - % Composition - Simplest Formula
1. Calculate the molar mass of a) F2 38 g/mol
b) Al2(SO4)3 342 g/mol
c) CuSO4.5H2O 249.5 g/mol
2. What is the mass of a) 5.0 moles of boron? 54 g
b) 0.25 moles of glucose, C6H12O6? 45 g
3. Calculate the number of moles in 21 g of aluminum.
How many aluminum atoms does this represent? 0.78 mol, 4.7 x 1023 atoms
4. Determine the number of moles in 115 g of ethanol, C2H5OH. 2.5 mol
5. How many moles of C atoms are there in 6 moles of CH3COOH? 12 mol
6. Calculate % a) of Ca in CaCO3. 40%
b) composition of CaCO3. 12% C, 48% O
c) of water in Na2CO3.10H2O. 62.9%
8. 9.23 g of calcium are heated in an excess of nitrogen. The final product has a mass of 11.38 g.
Determine its formula. Ca3N2
9. Many crystalline compounds contain water of hydration that is driven off when the compound is heated. The loss of mass in heating can be used to determine the simplest formula. For example, a hydrate of cobalt chloride, CoCl2.XH2O, weighing 0.809 g was heated until all the combined water was expelled. The dry powder remaining weighed 0.442 g. Determine the formula of the hydrate. CaCl2.6H2O
10. A sample of an organic compound containing carbon, hydrogen and oxygen (CxHyOz), which weighed 12.13 mg, gave 30.6 mg of CO2 and 5.36 mg of H2O on combustion. Determine the simplest (empirical) formula of this compound. C7H6O2
Problem Type: Problems based on Concentrations of Solutions
1. Calculate the # of grams of KMnO4 required to prepare 6.0 L of a 0.050 M solution. 47 g
2. What volume of 18 M acid is needed to prepare 5.0 litres of 6.0 M acid? 1.7 L
3. How would you prepare 150 mL of 0.40 mol/L Na2SO4 from a 2.0 mol/L solution? (calculate & explain)
4. If 3.0 litres of 6.0 mol/L HCl are added to 2.0 litres of 1.5 mol/L HCl, what is the resulting concentration? (Assume the final volume to be exactly 5.0 litres). 4.2 mol/L
5.
Problem Type: Balancing Equations
Balance the following equations and name the type of chemical reaction. (add states)
1. AgNO3 + MgBr2 AgBr + Mg(NO3)2
2. N2 + O2 NO2
3. P + I2 PI3
4. Al + MnO2 Mn + Al2O3
5. C6H6 + O2 CO2 + H2O
6. Al(OH)3 + H2SO4 Al2(SO4)3 + H2O
8. Cu(OH)2 + H3PO4 Cu3(PO4)2 + H2O
Problem Type: Balancing Equations ( cont’d)
Determine the products of the following reactions (if a reaction occurs), balance the equation and name the type of reaction: if no reaction just write NR and state why there is no reaction.
Na(s) + FeCl2(aq) à
KOH(aq) + H2SO4 (aq)à
Zn(s) + LiOH (aq)à
C2H5OH(l) + O2(g) à
Complete a total ionic equation, circle the spectator ions and complete a net ionic equation for the following: Add in the states of the reactants and products. (You need to balance the equations first!)
Al(OH)3 + H2SO4 Al2(SO4)3 + H2O
AgNO3 + MgBr2 AgBr + Mg(NO3)2
Problem Type: Stoichiometry Problems - Mass-Mass Relationship
1. How many grams of calcium carbide are needed to produce 6.5 g of acetylene?
The reaction is: CaC2 + 2H2O C2H2 + Ca(OH)2 16 g
2. How many grams of nitrogen dioxide may theoretically be produced when copper reacts completely with 21 g of concentrated nitric acid? (Assume that the concentrated acid is 100% HNO3). 7.7 g
3. Enough sulphuric acid was added to 100 mL of a solution of barium chloride to convert all the barium to insoluble barium sulphate. If 0.854 g of barium sulphate was formed, calculate the mass of barium chloride in 100 mL of solution.
0.762 g
4. How much ethyl alcohol, C2H5OH, can theoretically be produced by the alcoholic fermentation of 5.0 kg of dextrose sugar, C6H12O6?
C6H12O6 zymase 2C2H5OH + 2CO2 2.6 kg
5. How much limestone, which contains 90% calcium carbonate, must be used to produce 5.0 x 102 kg of calcium oxide? CaCO3 CaO + CO2 9.9 x 102 kg
6. Nitric acid reacts with silver according to the equation
4HNO3 + 3Ag 3AgNO3 + NO + 2H2O 170 g AgNO3
Calculate the number of grams of AgNO3 and NO produced when 126 g of HNO3 is added to 108 g of Ag.
Problem Type: Naming Compounds and providing chemical formula(e)
# / Formula / Name / # / Formula / Name1. / Cu2SO3 / 16. / Ag2S2O3
2. / Pb(MnO4)2 / 17. / Zn(IO4)2
3. / HF(aq) / 18. / H2CrO4(aq)
4. / (NH4)2SO4 / 19. / Sb2S3
5. / Ca(HCO3)2 / 20. / Hg2(CH3COO)2
6. / Be(NO2)2 / 21. / MnCO3
7. / HIO3(aq) / 22. / As(CN)3
8. / MgSO4*5H2O / 23. / HSCN(aq)
9. / KHC2O4 / 24. / Ba(BrO)2
10. / Cu(OH)2 / 25. / CuSO4 *4H20
11. / Fe2(CrO4)3 / 26. / AlH3
12. / (NH4)2Cr2O7 / 27. / SO2
13. / V(CO3)3 / 28. / N2O
15. / P2O5 / 30. / BCl3
More naming and formulae
# / Formula / Name / # / Formula / Name1. / gold(III) fluoride / 16. / cobalt (II) borate
2. / manganese(III) hydroxide / 17. / silver chromate
3. / sodium chlorite / 18. / tin(II) oxalate
4. / aluminum acetate / 19. / ammonium thiocyanate
5. / tin(IV) chloride / 20. / acetic acid
6. / zinc phosphide / 21. / ammoniun perchlorate
7. / aluminum bromate / 22. / sulphuric acid
8. / carbon tetrachloride / 23. / antimony(V) permanganate
9. / hypophosphorous acid / 24. / nitrogen triodide
10. / boric acid / 25. / barium sulfate hexahydrate
11. / iron (III) chromate / 26. / nitric acid