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Grade 11 Chemistry Review
Unit 1: Matter and Chemical Bonding (text chapters 1, 2, 3)
1) Clearly define and give an example of each of the following:
a) physical change
b) chemical change c) physical property d) chemical property
e) pure substance
f) element
g) compound
h) solution
i) solvent
j) solute
k) mechanical mixture
l) homogeneous mixture
2) Describe the location, relative mass, and charge of the three main components of the atom.
3) a) According to Lavoisier, if 35.0 g of Iron reacts with Oxygen to produce a substance with a total
mass of 50.0 g, what mass of Oxygen must have reacted?
b) In another experiment, 58.4 g of Iron reacts with Oxygen to produce a substance with a total mass
of 75.5 g. According to Proust, is this substance the same as the substance in question 3a?
Explain.
4) Define the following terms: mass number, atomic number, isotopes, average atomic mass.
5) Complete the following table.
Isotope
109
47
35
17
# of protons
# of electrons
# of neutrons
Ag
Cl
Mg-25
6) Complete the following table. All atoms are neutral unless otherwise indicated.
Element
# of electrons in
1st energy level
# of electrons in
2nd energy level
# of electrons in
3rd energy level
# of electrons in
4th energy level
Sulfur
Calcium
Neon
Lithium
Sulfur 2Aluminum 3+
7) According to current atomic theory, what is the maximum number of electrons that can exist in the
following energy levels:
a) 1st energy level
b) 2nd energy level
c) 3rd energy level
d) 4th energy level
8) Construct Bohr diagrams for the following atoms: a) N,
9) Define the following terms:
b) Ca, c) O2-, d) Mg2+
a) first ionization energy
c) electronegativity
b) electron affinity
d) atomic radius
10) Describe how each of the properties in question 9 changes as:
a) one moves from left to right across the periodic table,
b) as one moves from top to bottom down a family in the periodic table.
11) Describe where in the periodic table you would find the following:
a) metal elements
b) non-metal elements
c) Noble Gases
d) Alkali Metals
e) Halogens
f) Metalloids
g) transition elements
h) most reactive non-metal element
i) most reactive metal element
12) a) Draw a Lewis dot diagram for a typical atom in the Oxygen family of elements.
b) What is the typical ionic charge of these atoms.
c) Explain why all the atoms in this family form stable ions with this charge.
13) The Alkali Metals are a very reactive family of metals.
a) Explain what happens to these atoms when they react with an atom of Chlorine.
b) Why do all atoms in this family behave in this manner with Chlorine?
c) Potassium will be a much more reactive metal than Lithium. Explain why this is so.
14) Distinguish between the following types of chemical bonds:
a) ionic bond b) covalent bond c) polar covalent bond d) co-ordinate covalent bond
15) Consider the compound Sodium acetate.
Classify each bond as
a) purely covalent
b) polar covalent
c) ionic
16) i) Use Lewis Dot diagrams to illustrate the bonding between atoms of the following elements.
Indicate the type of bond in each case (purely covalent, polar covalent (indicate partial positive and
negative poles), and ionic (indicate the positive and negative ion).
ii) Give the correct IUPAC name and chemical formula of the resulting compound.
a) Oxygen and Sodium
b) Carbon and Fluorine
c) Magnesium and Nitrogen
d) Aluminum and phosphorus
e) Hydrogen and Sulfur
d) Oxygen and itself
17) Give the correct IUPAC name for the following compounds:
b) SiO2
c) N2O4
d) Cu(NO3)2 e) CCl4
a) PbS2
g) ZnCl2
h) Fe2(SO4)3 i) H2CO3(aq) j) H2S(aq)
k) KI
m) KClO4
n) CH3CO2H o) NO2
p) CO
q) MgCO3
s) H3PO3(aq) t) HNO3(aq)
u) NCl(aq)
v) PBr5
w) H2SO4(aq)
y) SF6
z) MgSO4A7H2O
f) AlPO4
l) CaS
r) Co(NO2)2
x) HClO3(aq)
18) Write the correct chemical formula for the following compounds
a) Zinc oxide
b) Mercury (I) nitride
d) Lead (II) carbonate
e) Lithium iodide
g) Silver chloride
h) Nickel (III) sulfide
j) diphosphorous trioxide
k) Copper (II) sulfate
m) aqueous hydrogen nitrite
n) ammonia
p) Beryllium chlorate
q) bromic acid
s) hydrofluoric acid
t) dinitrogen monoxide
v) Cobalt (III) phosphide
w) Boron trichloride
y) Gold (I) hypochlorite
z) Barium hydroxide octahydrate
c) Calcium nitrate
f) Tin (II) acetate
i) Carbon dioxide
l) Aluminum nitrate
o) ammonium phosphate
r) water
u) Rubidium sulfide
x) Iron (III) sulfite
19) Identify the shape of the following molecules and characterize them as “polar” or “non-polar.”
If the molecule is polar, identify the side that would be slightly negative.
a) NH3 b) CCl4 c) PF3 d) CO2 e) HCN f) HCF3 g) BH3 h) HI i) O2 j) H2O k) PH3
20) i) Draw structural diagrams for the following molecules.
ii) Identify any co-ordinate covalent bonds present.
iii) Show how hydrogen ions would bond to the molecules in ‘d’ and ‘e’ to form neutral molecules.
iv) How would the ions in ‘d’ and ‘e’ form in the first place?
a) NH4+ b) SO2 c) SO3 d) SO32- e) SO42- f) C2H6O
21) Balance the following skeleton chemical equations.
a) H2O2 —> H2O + O2
b) Na3PO4 + CaCl2 —> NaCl + Ca3(PO4)2
c) NH3 + O2 —> N2 + H2O
d) C5H6 + O2 —> CO2 + H2O
22) Predict the products of the following reactions and balance the resulting chemical equation.
If there is no reaction, write ‘NR.’
a) P4 + O2 —>
b) Zn(s) + FeCl3(aq) —>
c) HCl(aq) + Al(OH)3(aq) —>
d) KBr(aq) + (NH4)2SO4(aq) —>
e) CaCO3 + Energy —>
f) Ag + Cl2 —>
g) Pb(s) + Mg(NO3)2(aq) —>
h) NaCO3(aq) + CaCl2(aq) —>
i) NaHCO3 + Energy —>
23) Classify the type of reaction involved in each of the following and state if it is endothermic or
exothermic.
a) Ca(OH)2 + H2SO4(aq) —> 2H2O + CaSO4(aq) + Energy
b) P2O5 + 3H2O(l) —> 2H3PO4(aq) + Energy
c) 2Mg(s) + O2(g) —> 2MgO(s) + Energy
d) 2Al(s) + 3Cu(NO3)2(aq) —> 3Cu(s) + 2Al(NO3)3(aq) + Energy
e) Zn(HCO3)2(s) + Energy —> ZnO(s) + 2CO2(g) + H2O(g)
f) 2AgNO3(aq) + NiCl2(aq) —> 2AgCl + Ni(NO3)2(aq) + Energy
g) 2LiCl(s) + Energy —> 2Li(s) + Cl2(g)
Unit 2: Quantities in Chemical Reactions (text chapters 4 and 5)
1) Determine the molar mass of the following pure substances.
a) Ca
b) Mg(OH)2
c) Zn
d) N2
e) CH3CH2CH3
f) (NH4)2SO4
2) If 0.25 moles of a compound has a mass of 30.0 g, what is the molar mass of the compound?
3) What mass of CaCl2 would you have if you had 1.75 moles of this compound?
4) How many moles of KCN would you have if you had 25.4 g of this compound?
5) How many particles are there in 1 mole of anything?
6) If you had 6.5 g of liquid Nitrogen and allowed it to fill a balloon as Nitrogen gas:
a) How many Nitrogen molecules would you have?
b) How many Nitrogen atoms would you have?
c) Why are these two numbers not the same?
7) What is the percent composition by mass of the compounds:
a) H2O
b) CO2
c) HCN
d) Al2(CO3)3
8) A compound is found to have the following percentage composition by mass:
30.57 % Carbon, 3.83 % Hydrogen, 45.22 % Chlorine, 20.38 % Oxygen.
a) Determine the empirical formula for this compound.
b) Based on a molar mass of 157.0 g, what is the molecular formula of this compound.
c) Give at least one method that might have been employed to determine the molar mass.
9) For the following reaction: 3Mg(s) + N2(g) —> Mg3N2(s)
a) If 109.5 g of Mg and 70.00 g of N2 gas are reacted in this manner, determine the limiting reactant.
b) Using the data from part ‘a,’ what mass of Magnesium nitride should be produced?
c) If the reaction in part ‘a’ took place and 148.5 g of Mg3N2(s) is collected, what is the percent yield?
10) For the following reaction: 3Ca(s) + 2H3PO4(aq) —> Ca3(PO4)2(s) + H2(g)
a) What mass of Calcium would be needed to produce 8.5 moles of Hydrogen gas?
b) If 2.35 g of Calcium phosphate are produced in this reaction, what mass of Calcium must have
reacted?
c) How many moles of phosphoric acid will be needed to completely react with 15.5 g of Ca?
11) In the following reaction: 3(NH4)2CO3(aq) + 2Al(NO3)3(aq) —> Al2(CO3)3(s) + 3NH4NO3(aq),
what mass of Aluminum carbonate precipitate would be produced if 2.50 g and ammonium
carbonate reacts completely?
Unit 3: Solutions and Solubility (text chapters 6, 7, 8)
1) Write the ionic equation for the dissolution of MgCl2(s) in water.
2) The single displacement reaction of aqueous Lead (II) nitrate with solid Aluminum metal produces
solid Lead metal and aqueous Aluminum nitrate.
a) Write a balanced chemical equation for this reaction.
b) Write the ionic equation for this reaction.
c) Write the net ionic equation for this reaction.
3) Write the following: a) balanced equation, b) ionic equation, and c) net ionic equation, for the
following neutralization reaction:
2HCl(aq) + Ba(OH)2(aq) —> BaCl2(aq) + 2H2O(l)
4) If a solution contains both dissolved sulfate (SO42-)ions and dissolved sulfide (S2-) ions, use the
solubility rules to determine a dissolved metal ion that could be added to the solution to remove the
sulfate ions by precipitation while leaving the sulfide ions in solution.
5) Determine the molar concentration of a solution prepared by dissolving 2.56 g of CuSO4 into
750.0 mL of solution.
6) What mass of Sr(NO3)2 would be needed to prepare 500.0 mL of 0.500 mol/L solution?
7) A stock solution of HCl(aq) has a molar concentration of 12.0 mol/L. What volume of this stock
solution is needed to prepare 250.0 mL of 0.500 mol/L solution. Describe how you would accurately
prepare this dilute solution.
8) If 50.0 mL of 1.50 mol/L NaCl solution is diluted to a new volume of 750.0 mL, what will be the
concentration of this new solution?
9) For the reaction between Lead (II) nitrate and Potassium chloride:
Pb(NO3)2(aq) + 2KCl(aq) —> PbCl2(s) + 2KNO3(aq)
a) If you want to completely react 100.0 mL of 0.50 mol/L lead (II) nitrate solution, what volume of
0.35 mol/L Potassium chloride will you need?
b) If 55.5 mL of 0.25 mol/L Potassium chloride completely reacts, what mass of Lead (II) chloride
should be produced?
10) A solution of Lithium hydroxide is titrated using a standard solution of 0.150 mol/L sulfuric acid.
2LiOH(aq) + H2SO4(aq) —> 2H2O(l) + Li2SO4(aq)
a) If 25.0 mL of the acid are needed to neutralize 10.0 mL of the base, what is the concentration of
the base?
b) Write an ionic and net ionic equation for this reaction.
11) Identify a set of test solutions and a correct sequence for adding them to an unknown solution to
test for the presence of one or more of the following metal ions: Mercury (2+), Silver, and Barium.
12) What is the concentration of sulfate ions in a 2 mol/L solution of Aluminum sulfate?
13) Identify the type of solution (unsaturated, saturated,
supersaturated) for the following in the graph to the right:
a) 0.7 mol/L @ 70 oC
b) 0.4 mol/L @ 34 oC
c) 1.0 mol/L @ 90 oC
14) Describe what will take place as a 0.8 mol/L solution @
85 oC is cooled to 20 oC while being agitated.
15) How would you prepare a supersaturated solution?
16) Calculate the pH of the following solutions:
a) .01 mol/L of HCl(aq), b) 0.001 mol/L of H2SO4(aq),
c) 2 mol/L of acetic acid (strength is 13%)
17) Give the [H+] concentration in the following solutions:
a) pH = 4,
b) pH=10.5,
c) pOH = 6
18) a) Why is the pH of a 0.1 mol/L solution of HNO3 equal to 1.0 when the pH of a 0.1 mol/L solution
of HCN is 5.0?
19) a) Give Arrhenius and Brønsted-Lowry definitions for acids and bases.
b) What improvements did the Brønsted-Lowry theory make compared to the Arrhenius theory?
20) Identify the Brønsted-Lowry acid, base, conjugate acid, and conjugate base for the following
reaction: NaHSO4(aq) + NaHCO3(aq) —> Na2SO42-(aq) + H2CO3(aq)
21) Write chemical equations to show that NaHCO3 is amphoteric (can act as an acid or base).