Download SCH 3U – Chemistry 11 - Exam Review

SCH3U
Exam Tuesday, January 28, 8:30 am
Final Exam Review
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This review is a good sample of questions covered in this unit. It is by no means the only
topics/problem types that you should know. You should also review tests, read over notes,
make a list of problem types for each unit and make a list of formulas you need to memorize.
Include units and consider significant digits for problem solving.
Answers are available in the classroom.
For +1% of your final mark, answer all questions on lined paper (not typed) and hand in on
the day of the exam.
Unit 1 – Matter and Bonding
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How many significant digits are in the following: a) 1.00 b) 0.0987 c) 900 d) 308.009
Compare accuracy and precision.
Give the atomic notation(standard atomic notation) for nitrogen.
Describe the parts of an atom
What do the group 17 elements have in common in terms of their electrons?
Define and give the trends for a) Atomic radius, b) Ionization energy c) Electron affinity, and d) Electronegativity
Compare properties of ionic and covalent compounds
What type of bond will the following form to make (ionic, polar covalent, non-polar covalent)? Include your calculations:
a) Ca-Cl b) C-O c) Br2
9. Identify whether each of the following are polar or non-polar molecules. Explain your reasoning and include a Lewis
diagram and a VSEPR diagram a) CO2 b) NH3 c) H2
10. Natural neon contains three isotopes: Ne-20 with a mass of 20amu, Ne-21with a mass of 21amu, and Ne-22with a mass of
22 amu. In a sample, 90.92% of the atoms are Ne-20, 0.26% of the atoms are Ne-21, and 8.82% of the atoms are Ne-22.
Calculate the average atomic mass of neon.
11. Complete the following for each of the molecules listed below: a) Draw the Lewis Structure b) Draw and name the VSEPR
structure c) determine if the molecule is Polar or Non-polar 1) BeF3 2) CH4 3) H2O 4) NH3 5) N2
Unit 2 – Nomenclature/Chemical Reactions
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Give the formula/name of the following: a) magnesium chloride b) FeCl 2 c) cuprous fluoride d) diphosphorus
pentasulphide e) Ba3(CO3)2 f) copper (II) hydroxide g) iron (II) sulfite h) chromium(III) nitrate nonahydrate i) potassium
thiosulphate j) hydrofluoric acid k) H3PO4 l) CuSO4.5H2O l) sulfurous acid
Balancing reactions: p. 118 # 9 c, d , p. 156 # 34
Provide an example of a synthesis, decomposition, and combustion reaction.
p. 127 # 21 f, g, h
Activity series and SDR: p. 131 # 24 a, b, d, e, 23
Predicting DDR: p. 140 # 2 a, b, c (consider solubilities and neutralization reactions)
p. 141 # 4
Unit 3 – Chemical Quantities
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Define a mole.
How many molecules are in a) 99.00 grams of zinc b) 9.0 x 10-3 mg of lead?
What is the mass of 0.89 moles of hydrogen gas?
Percentage composition: p. 201 # 2 p. 204 # 6
Empirical formula: p. 209 # 10 p. 211 #15
Molecular formula: p. 218 # 20
The combustion (reacting with oxygen) of propane (C3H8 (g)) produces liquid water and carbon dioxide. Write the balanced
chemical reaction and answer the following questions a) if 10.0 g of propane reacts with excess oxygen, what mass of
water will be produced b) if 4.5 g of propane reacts with 3.0 g of oxygen, what mass of carbon dioxide will be produced?
8. Another limiting reactant problem: p. 259 # 6
9. How many atoms of oxygen are in 1.87 g of aluminum carbonate?
10. Percentage yield: p. 262 # 31, 32
SCH3U
Exam Tuesday, January 28, 8:30 am
Unit 4 – Solutions and Solubility
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Define the following: solution, solvent, solute, miscible, immiscible, saturated solution, unsaturated solution, super
saturated solution.
Describe three factors that affect the RATE of dissolving
Describe the dissolving process.
Describe why salts commonly dissolve in water.
Describe why polar molecules dissolve in water.
Explain “like dissolves like”
What factors determine the solubility of solutes.
What is the mass of 13 mL water? What is the mass of 1.2 L of water?
m/m %, v/v %, ppm, and ppb: p. 308 # 9 p. 310 # 11 p. 15
What is the concentration of a 200 mL solution with 4.8 g of NaCl?
What mass of aluminum is in a 45 mL 1.0 x 10 -2 mol/L solution of aluminum sulphate?
A 1350 mL solution contains 3.00 x 1024 ions of potassium. What mass is potassium carbonate was originally dissolved
in the solution?
Diluting solutions: p. 324 # 3, 5
Solubility chart: p. 335 # 3
Predicting a precipitate: p. 339 # 4 i, j, k, l
Describe the technique of qualitative analysis.
Net ionic equations: p. 347 # 2
Limiting reactant: p. 355 # 12, 13
p. 356 # 7
Unit 5 – Acids and Bases
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Compare four observable properties of acids and bases.
Define a) an Arrhenius acid and base b) a Bronsted-Lowry acid and base (which theory is more current?)
Identify the acid base conjugate pairs in HF(aq) + H2O(l) -> H3O+(aq) + F–(aq)
What is the conjugate acid of NH3?
Define a strong acid and weak acid. Give an example of each.
Define a strong and weak base (careful, this one is not as simple as #24). Give and example of each.
What is the pH of a) 0.97 mol/L HBr b) 0.065 mol/L of NaOH
Define the following: indicator, end point, equivalence point
Titration problems: p. 398 # 13
Unit 6 – Gasses
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Compare the forces between and motion of particles of a) solids b) liquids c) gases
What are the assumptions of KMT?
Define pressure
Convert the following to kPa a) 700 torr b) 1.15 atm c) 15.0 psi d) 100 000 Pa
Define the following Laws: a) Boyle’s b) Charles’ c) Gay-Lussac’s d) Combined gas e) Avogadro’s g)Ideal gas
Describe the Kelvin scale and define absolute zero.
What is the Kelvin temperature equivalent to 35 °C?
p. 435 # 6, p. 446 # 8. 10 p. 450 # 15, p. 451 # 2, p. 457 # 21
How do real gasses behave that is different from KMT?
Ideal Gas Law p. 488# 13,14
Density of gasses (apply what you know!): p.493#16,17
Gas Law Stoichiometry p.506#30,31, p.511#35,36