Download chemistry 11 exam review

Unit #1 Matter, Chemical Trends, And Chemical Bonding (p8- 105)
 Significant Figures
 The Structure of the Atom (atomic mass, atomic number, #p, #e, #n)
 Isotopes, Radioisotopes, Half-life
 The Periodic Table
 Periodic Table Trends (electronegativity, ionization energy, atomic radius, reactivity, electron
affinity)
 Ionic and Molecular Compound Physical Properties
 Lewis Dot Diagrams (electron dot diagrams)and Structure Formulas
 Ionic and Covalent Bonding
 Polarity of Bonds and Polar Molecules
 Intermolecular Bonds (Dipole-dipole, VanderWaals, LDF, Hydrogen Bonding)
 The Names and Formulas of all Compounds
Practice:
1. How many significant digits are in a) 200 152 ......... 1.0040 x 106 ........... 0.0048 ............ 7.0 .........
2. 18.987 g / 3.06 = ............................... ............... which has ................. significant digits
3. What is an example of a physical property of matter
4. What is an example of a chemical property of matter
5. What is an example of a physical change
6. What is an example of a chemical change
7. What is an example of a quantitative property of matter
8. If a mixture has only one phase it is said to be .............................................. 2 phases ?
..................................................
9. Isotopes of the elements contain different numbers of ..........................but the same # of
............................. & ...........................
10. An alpha particle contains ..................... protons & ................ neutrons and weighs ................,
symbol = ...................
11. When an isotope emits a beta particle the atomic# of the substance ..............................& the at
mass ..............................
12. The isotope cobalt-60 ( 6027Co ) has .................... protons & ............. neutrons in its nucleus with
............... electrons around it
13. The isotopic ion 5625Mn 2+ has............ protons & .............. neutrons in its nucleus with ....................
electrons around it.
14.Fill in the following table
Chemical Formula
Name of substance
Identify the Type of
Molecule
HNO3(aq)
Aqueous Ammonium hydroxide
N2H4
Iron(III) sulphide
Copper(II)chromate pentahydrate
15. Complete the table below.
Element
name
fluorine
Element
symbol
Atomic
number
Group
number
Family
name
Period
number
noble gas
3
Metal or
nonmetal
Ba
16. Use electron dot diagrams to explain the formula for Mg3P2.
17. Draw the Lewis diagram for carbon dioxide, water and methane, nitrogen triiodide, and oxygen
gas. Include dipoles, type of bonding ( intermolecular and intramolecular), and whether the
molecule is polar or nonpolar.
18. List all families on the periodic table and state two properties of each.
19. Explain all periodic trends, state the trend for periods and groups.
4.
Unit #2 Chemical Reactions (p106- 153)
 Conservation of Mass
 Write Word Equations, and Balanced Chemical Equations with States.
 Types of Reactions (synthesis, decomposition, single and double displacement, incomplete and
complete combustion)
 Reactivity Series for single displacement reactions
 Solubilty Chart for double displacement reactions
 Neutralization Reactions
 Chemical Reactions for the formation of acids and bases from metal oxides and non-metal oxides
Practice:
1. Predict the products and write a balanced chemical equation for the following chemical reaction:
Hydrochloric acid is neutralized by an aluminum hydroxide suspension.
2. Predict the products and write a balanced chemical equation for the following chemical reaction.
Charcoal (pure carbon) is burned in a barbecue.
3. Complete the following chemical reaction equation, including states of matter and balancing:
FeCl3(aq) + Zn(s) 

4. Balance:
a)
Cu
+
O2

CuO
b)
P4
+
H2

PH3
c)
Al
+
CuSO4 
Al2(SO4)3
+
Cu
d)
H2SO4 +
Al

Al2(SO4)3
+
H2
e)
Na
+
H2O 
NaOH
+
H2
f)
NH3 +
O2

NO2 +
H2O
g)
Na2CO3 +
H3PO4
 Na3PO4
+
H2O +
CO2
h)
FeS +
O2

Fe2O3
+
SO2
i)
Na2CO3 +
FeCl3 + H2O 
Fe(OH)3 + NaCl +
CO2
5. Identify the type the reaction, complete & balance the following reaction equations. Include
states
a) A chunk of sodium is dropped into a cup of water.
b) Some butane is burned in excess oxygen.
c) Some butane is burned in limited oxygen.
d) A solution of ammonium phosphate is added to a solution of copper(II) sulfate.
e) A hydrated salt, magnesium sulfate heptahydrate, is heated vigorously over a burner.
f) A solution of platinum(III) nitrate is stirred with a zinc-plated spoon.
g) A piece of magnesium is left overnight in a solution of acetic acid
h) H3PO3 (aq)+ NH4OH (aq)→
i) CaCO3 (aq) + H3PO4 (aq)→
j) Cu (s) + H2O (l)→
kj) Al (s) + Cu(NO3)2 (aq)→
Unit #3 Quantities in Chemical Reactions (p158- 259)
 Isotopic Abundance Calculations
 Law of Definite Proportions
 The Mole (Avogadro’s Number)
 Molar Mass Calculations
 Calculation of particles (molecules, atoms)
 Mole to Mass Calculations
 Percentage Composition
 Empirical and Molecular Formulas
 Stoichiometry (mass to mass calculations)
 Limiting and Excess Reactants
 % Yield
Practice:
1. Convert 10.0 g of copper to mol.
2. How many mol is 24.7 g of aluminum?
3. 42.8 g of oxygen gas is how many mol of oxygen gas?
4. Convert 4.62 mol of zinc hydroxide to grams.
5. How many grams is 8.32 mol of calcium phosphate?
6. Determine the number of molecules found in 5.00 mol of water.
7. Determine the number of hydrogen atoms found in 2.50 mol of methane, CH4.
8. Determine the number of atoms in 0.25 mol of sodium.
9. Consider the following balanced equation: Ca(OH)2 + 2HCl
CaCl2 + 2H2O
Determine the number of moles of calcium chloride produced when 8.0 mol of HCl is combined
with excess Ca(OH)2.
10. What mass of FeCO3 is required to form 28.6 g of Fe2O3 ?
Balanced equation: 4FeCO3 + O2
2Fe2O3 + 4CO2
11. If 17.5 g of C is dissolved in water and added to a solution containing 39.5 g of SO2, the reaction
occurs:
5C + 2SO2  CS2 + 4CO
i) Determine the limiting reactant.
ii) Determine the mass of CO produced.
12. a) Balance the following equation: KO2(s) + CO2(g)  K2CO3(s) + O2(aq)
b) If 1.50 g of KO2 was present in the reacting mixture, with excess CO2, calculate the theoretical
yield of K2CO3.
c) Determine the percentage yield of 1.2 g of K2CO3.
13. Magnesium hydroxide, Mg(OH)2, can be commonly found in antacids. Calculate the percentage
composition, by mass, of each element in magnesium hydroxide.
14. What is the empirical formula of a compound whose percentage composition is found to be 2.20%
hydrogen, 26.7% carbon, and 71.1% oxygen?
15. Analysis of an air pollutant reveals that the compound is 30.4% nitrogen and 69.6% oxygen. The
molar mass of this compound is 92.0 g/mol. What is the molecular formula of this compound?
(N2O4)
Unit #4 Solutions and Solubility (p264- 411)
 Solution Terminology (solute, solvent, homogeneous, heterogeneous, electrolytes)
 Water as a Universal solvent
 Dissolving (likes dissolve likes)
 Solution Concentration (%w/w, 5v/v, 5w/v, ppm, molar concentration)
 Water Treatment
 Solution preparation (from a solid)
 Dilution Problems(ccvc = cDvD)
 Solubility curves (unsaturated, saturated, supersaturated)
 Reactions in Solutions (Complete and net ionic equations)
 Qualitative Chemical Analysis(flame test, conductivity, colour of solution, sequential
flowchart(ppt test))
 Acids/Bases Properties
 pH and litmus paper
 Dilution and pH
 Strong and Weak Acids
 Arrhenius, Lowry-Bronsted Definitions
 Acid-Base Titrations
Practice:
v
1. Determine the concentration of a solution containing 5.6 g sodium nitrate in 250 mL of water.
Express your answer in mol/L.
2. Consider the following reaction: Barium chloride solution is mixed with potassium sulphate
solution to produce a precipitate barium sulphate and potassium chloride.
For this reaction, write
(a) a balanced chemical equation
(b) a total ionic equation
(c) a net ionic equation
3. a)Describe how you would prepare 100 mL of a 0.5 mol/L solution of sodium chloride. Include
your calculation and the procedure.
b)In a titration 20.0 mL of 0.10 mol/L sodium hydroxide was neutralized by an average of 9.3 mL
of hydrochloric acid. Determine the concentration of the acid.
4. A solution of copper(II) sulphate is used at blood donor clinics to test donor blood for low iron.
Calculate the concentration of this solution, in mol/L, if 125 g of CuSO4 crystals are dissolved in
1.0 L of water.
5. What volume of a 17.4 mol/L acetic acid stock solution is required to make 2.0 L of a 1.5 mol/L
acetic acid solution?
6. For a chemical analysis, 750 mL of a 0.480 mol/L potassium permanganate solution is to be
prepared. Calculate the mass of potassium permanganate crystals that will be dissolved to make
this solution.
7. Calculate the pH of vinegar which has a hydrogen ion concentration of 5.2  10–2 mol/L.
8. A swimming pool has a pH of 7.5. Calculate the hydrogen ion concentration in the pool.
9.
Draw a flow chart illustrating how you could analyze a solution that may contain Pb2+
and/or Sr2+.
10. Use the Bronsted-Lowry concept to identify the reactants as acids or bases in the following
reactions. In your answers, indicate the conjugate acid-base pair
a) HSO-4 (aq) + HCO3-(aq)  SO42-(aq) + H2CO3(aq)
b) HS- (aq) + HCO3-(aq)  H2S(aq) + CO32-(aq)
c)
11. Complete the following chart
Substance
Litmus
Paper
pH
Conductivity
Acid/Base/Neutral?
HCl
Ca(OH)2
NH3
CH3COOH
KNO3
12. Will ammonia dissolve in water? Will carbon dioxide dissolve in ammonia? Explain.
Provide the required diagrams.
13. A sample of well water contains 0.24 ppm of dissolved iron(III) sulfate form the
surrounding rocks. What mass of iron(II) sulfate is present in 1.2L of water in the kettle?
12.
Unit #5 Gases and Atmospheric Chemistry (p417- 499)
 Kinetic Molecular Theory
 Properties of the States
 Pressure and temperature Conversions
 Gas Laws (Boyles, Charles, Gay-Lussac, Combined Gas Law, Ideal Gas Law, Dalton’s Law of
Partial Pressure)
 Stoichiometry involving Gases
 Air Quality
Practice:
1. What volume does 2.50 mol of oxygen gas occupy at 3.60 atm and 400C? (38.3 L)
2. What temperature is 4.53 mol of gas that occupies 5.78L at 704 torr? (14.4 K)
3. How many mol of gas will occupy 7.42 L at 373K and 200.4 kPa? (0.479 mol)
4. What is the pressure (in atm) of 2.45 x 10-2 mol of gas that occupies 5.68 mL at -35C? (84.2atm)
5. What is the new volume of 750.0 mL of gas at 700.0 torr when the pressure is changed to 800.0 torr?
(656.3 torr)
6. 45 mL of gas at 15C and 790 torr is changed to 23C and 810 torr. What is the new volume? (45
mL)
7. 175 mL of gas at –30.0C and 2.57 atm is changed to standard conditions. What is the new volume?
(505 mL)
8. What pressure is needed to change 130 mL of gas at 740 torr to 150 mL? (641 torr)
9. What temperature change is needed to change 1.0 L of gas at 10.0C and 800.0 torr to 0.50 L and
760 torr? (-138C change to bring your final temperature to 134K)
10. A 1.0 L rubber bladder is filled with carbon dioxide gas in a 25C room (Pressure = 745 mm Hg).
What volume will the gas occupy when it is taken out into the open air where the temperature is 12C and the pressure is 742 mm Hg? (0.96 L)
11. A balloon containing 425 mL of air escapes from a little boy watching a parade. The temperature is
32C at street level and the pressure is 742 torr. When the balloon stops rising, its volume has
become 895 mL although the atmospheric pressure has decreased by only 300.0 torr. What is the
temperature at this level? (383 K = 110C)
12. To what temperature must 15 L of oxygen gas at 0.0C be heated at 1.00 atm pressure in order to
occupy a volume of 23 L, assuming that the pressure increases by 5.0 mm Hg? (421 K = 148C)
13. Determine the volume of a balloon at SATP assuming that it occupies a volume of 20.0 L at a
temperature of 12ºC and a pressure of 135 kPa.
14. What is the pressure reading on a scuba tank if the tank is filled with N2 gas at a partial pressure of
115 atm and O2 gas at a partial pressure of 30 atm?
15. Butane from a lighter undergoes combustion in the following manner:
2C4H10 (g) + 13O2 (g)  8CO2 (g) + 10 H2O (g)
What volume of butane was burned to produce 325 mL of CO2 (g)?
16. What amount of oxygen, in moles, is available for a combustion reaction in a volume of 12.5 L at
STP?
17. Potassium metal with a total mass of 15.0 g is dropped into a beaker of water. What volume of
hydrogen gas will be produced if the temperature is 15ºC and the pressure is 100 kPa?
Chemical equation: 2K(s) + 2H 2KOH(aq) + H2 (g)
18. What is the initial pressure of a gas if it occupied a volume of 375 mL, but now occupies a volume
of 1.25 L at a pressure of 95.5 kPa. Assume that the temperature remains constant throughout the
process.
19. Magnesium was added to hydrochloric acid, HCl, and produced 5.25 L of H2 gas at a temperature
of 325 K and a pressure of 100 kPa. What mass of Mg was used in this single displacement
reaction?