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1) Matter and Chemical Bonding: Including the relationship between periodic tendencies, types of chemical bonding, and the properties of ionic and molecular compounds; The type of chemical reactions and the reactivity of starting materials, symbols and formulae to represent the structure and bonding of chemical substances.

Over 400 instructing slides and hundreds of practice questions and answers are available. The course focuses on the basic concepts and theories of chemistry and allows students to study the behaviors of solids, liquids, gases, and solutions; explore changes and relationships in chemical systems; and understand how chemistry is related to products and processes that affect our lives and our environment. The course covers 11 chapters. The topics are detailed as follows:

G11 Chemistry Review Questions

Chemical Equations and Stoichiometry


1.Find the simplest formula for the compound with composition:
a) 38.7% carbon, 9.7% hydrogen and 51.6% oxygen (CH3O)

b) 82.4% nitrogen and 17.6% hydrogen (NH3)

2. A certain compound is 40.0% carbon, 6.7% hydrogen and 53.3% oxygen by weight. One mole of this substance weighs 180 grams. What is the molecular formula of the compound? (C6H12O6)

3. For the decomposition reaction of potassium chlorate: a) What volume of oxygen gas, measured at STP, is formed by the reaction of 75.0 g of potassium
chlorate? (20.6 L) b) How many moles of potassium chlorate must react in order to form 2.5 g of potassium chloride? (0.034 mol)

4. For the reaction: 6 CO2(g) + 6 H2O(g) � C6H12O6(s) + 6 O2(g)
a) Which substance is in excess when 100 g of carbon dioxide and 50.0 g of water are reacted?
(water)
b) How much glucose will be produced by the reaction in part (a)? (68.2 g)

5. For octane compustion reaction, a) What mass of carbon dioxide will be produced when 180 g of octane are completely burned? (555 g) b) What volume of carbon dioxide, at STP, will be produced when 62.7 g of oxygen are completely reacted? (28.1 L)

Kinetic Molecular Theory and the Gas Laws

1. Define or explain the following terms:
a) Kinetic Molecular Theory
b) Kinetic energy
c) STP and SATP
d) Potential Energy
e) Dalton's Law of Partial Pressure

2. Explain these observations in terms of the Kinetic Molecular Theory.
a) Gases are much less dense than solids/liquids.
b) Solids can not be significantly compressed.
c) The volume of a gas increases as it is heated.

3. Describe what is happening to the motion of the particles of a gas as they are cooled, assuming that there is no change in volume and no change in state.

4. When the temperature of a substance increases, the ________ (potential/kinetic) energy of the substance ______ (increases/decreases).

5. When a substance changes state from liquid to gas, the ___________ (potential/kinetic) energy of
the substance ________ (increases/decreases).

6. Calculate the volume of 1.00 mol of chlorine gas at STP. (22.4 L).

7. What is the volume of 250.0 g of hydrogen gas, at 25 degree C and 100.0 kPa? (3065 L)

8. At what Celsius temperature will 10.0 grams of ammonia exert 700.0 mmHg pressure in an 8.0 L container? (-120 C)

9. 10.0 g of gas occupies 8.00 L at STP. What volume would this gas occupy at 273 C and 160 kPa? (10.1 L)

10. A gas measuring 0.50 L at 99.0 kPa and 87 C is heated to 127 C inside a rigid container that can not expand. What is the new pressure of the gas? (110 kPa)

Solutions, Acids And Bases

1. Define or explain these terms.

a) Saturated and Unsaturated solution
b) Solute
c) Solvent
f) Super-saturated
d) Solubility 
e) Molar Concentration
f) Acid
g) pH
h) Base and Acid
li Acid-base indicator

2. Explain following observations:
a) Acids react with metals such as magnesium and calcium producing gas.

b) Acids react with compounds containing carbonate ions to produce gas.

c) An acid reacts with a base producing and a _______. The resulting solution is neither acidic nor basic, thus the reaction is called _______.

d) Sea water has a pH of 8. Sea water is (acidic, basic, neutral).

e) Lemon juice is very acidic. The pH might be (2, 5, 7, 9, 13).

3. Write a balanced chemical equation showing the reaction between:
a) magnesium and hydrochloric acid
b) zinc and acetic acid
c) sulfuric acid and calcium carbonate

4. a) What is the difference between a strong acid and a weak acid. Give an example of each.
b) How is a concentrated solution different than a dilute solution? c) How is a "strong" acid different than a "concentrated" acid?

5. 80.0 g of lithium hydroxide is dissolved in enough water to make 500 mL of solution. What is the molar concentration of the solution? (6.68 mol/L)

6. What mass of sodium acetate is present in 600 mL of a 4.00 M solution? (197 g)

7. Write balanced chemical equations for the reactions between the following solutions. Indicate any precipitates that will form.
a) potassium sulfate and copper (II) nitrate
b) ammonium bromide and lead (II) chlorate
c) lithium sulfide and barium acetate
d) sodium chloride and magnesium nitrate

Grade 11 Chemistry

Matter

1. Define or explain the following terms:

a)   Chemical Property 
b)   Element Solvent
c)   Matter 
d)   Physical Change 
e)   Compound 
f)    Solute
g)   Qualitative property
h)   Chemical Change 
i)    Solution
j)    Metal
k)   Quantitative property 
l)    Pure Substance 
m)  Homogeneous Non-metal
n)   Physical Property Mixture 
o)   Heterogeneous Metalloid

2. Explain the Difference
a) a compound and a solution
b) a molecule and a mole
c) a mixture and a solution
d) a chemical change and a physical change
c) a physical property and a chemical property
 f) a metal and a non-metal

3. Classify each of the following as either an element, compound, solution, or mechanical mixture. Which items are pure substances, and which ones are mixtures?

a) an iron bar
b) iron (III) oxide (rust)
c) ozone
d) concrete (sand, gravel and lime)
e) brass (copper and zinc)
f) ammonia gas
g) 24 karat (pure) gold
h) tap water
i) sulfur
j) distilled water
k) granite rock
r) silver nitrate

4. What are five physical properties and fure (or more) chemical properties for the element
aluminum? What is aluminum’s electron configuration? Is it a metal or a non-metal?

5. What are the four indications (signs) that a chemical change has taken place?

6. Classify each of the following as physical or chemical changes. If it is a chemical change, what is the evidence that a chemical change had occurred?

a) toasting bread ________
b) breaking glass ________
c) sugar dissolving in coffee ________
d) tarnishing of silver ________
e) allowing pop to go flat ________
f) grinding coffee beans ________
g) boiling water ________
h) heating platinum until it glows red ________
i) alka-seltzer fizzing in water ________
j) explosion of nitroglycerin ________
k) evapouration of water ________
l) firing a cap pistol ________
m) boiling an egg ________
n) zinc dissolving in hydrochloric acid ________

7. What is the difference in electronegativity between bonding atoms in:
ionic compounds? _________, pure covalent compounds? __________, polar covalent compounds? _________

8. Write the balanced nuclear reactions to show the:
a) alpha decay of U – 238
b) alpha decay of Am-243
c) beta decay of I-131
d) beta decay of C-14

Atomic Theory and Chemical Bonds

1. Define or explain the following terms:

a) Atom
b) Radio-isotope
c) Chemical Groups
d) Electron affinity
e) Nucleus Alpha-decay
f) Chemical Period
g) Chemical bond
h) Proton Beta-decay
i) Isoelectronic Stable
j) Octet Rule
k) Neutron Metal
l) Halogen Ionic Bond
m) Electron
n) Non-metal
o) Alkali Metal
p) Covalent Bond
q) Atomic Number
r) Metalloid
s) Alkaline Earth metal
t) Electron configuration
u) Noble gas
v) Polar covalent bond
w) Isotope
x) Orbital
y) Ionization energy

2. Give an example for each of following questions
a) Atoms and ions are isoelectronic
b) Isotopes

3. What one piece of information about an atom determines its chemical properties?

4. What holds the electrons in an atom close in to the nucleus?

5. What happens to the potential energy of an electron when it moves further from the nucleus?

6. One isotope of uranium is “U-238”. What does the number represent? How many neutrons does a U-238 atom have?

7. Write electron configurations for Zn, Ar, Ag and I.

8. List any three ions or atoms that are isoelectronic with a Ca2+ ion. Repeat the question for a P3- ion.

9. What are the trends on the Periodic Table for the following characteristics:
a) Ionization energy
b) Electron affinity
c) Electronegativity
d) Reactivity of Metals; Reactivity of Non-metals
e) Metallic characteristics

10. Answer the following questions about quantum levels and electron orbitals:
a) How is an orbital different from an orbit?
b) For n=1, What is the maximum number of electrons in this principle quantum level? _________
How many orbitals are there? ___________
How many types of orbitals are there? __________ List them: ______________________________
c) For n=2, What is the maximum number of electrons in this principle quantum level? _________
How many orbitals are there? ___________
How many types of orbitals are there? __________ List them: ______________________________
d) For n=3, What is the maximum number of electrons in this principle quantum level? _________
How many orbitals are there? ___________
How many types of orbitals are there? __________ List them: ______________________________
e) For n=4, What is the maximum number of electrons in this principle quantum level? _________
How many orbitals are there? ___________
How many types of orbitals are there? __________ List them: ______________________________

11. What is the maximum number of electrons that can be designated (held in): 3s _________, 2p _________,
4d _________, 1s ________ , 4f _________, 6s _________, 3d __________, 5p __________

Moles and Molecular Formulas

1.a) One molecule of acetic acid contains _____________ atoms.

b) The molar mass of water is _______.

c) 30 grams of water is ____________________ moles.

d) 30 grams of water contains ____________________ molecules.

e) 30 grams of water contains ___atoms.

f) 30 grams of water contains ____________atoms of hydrogen.

g) 30 grams of water contains ____________ grams of oxygen.

2.a) What is the molar mass of CO ? _________________

b) 220 g of CO contain ______molecules.

c) 4.8E24 molecules of CO weigh __________________ grams.

d) 89.6 L of carbon dioxide at STP is ______ moles weighing _______grams.

e) 8.8 g of carbon dioxide is ____ moles and contains __________molecules.

f) The volume of 8.8 g of carbon dioxide at STP is ________________ litres.

g) In 44 g of carbon dioxide there are __________________ grams of carbon.

h) In 220 grams of carbon dioxide there are _________ grams of carbon and _________ grams of oxygen.

i) In 220 grams of carbon dioxide there are ______________ oxygen atoms.

j) 100 L of CO contain _____ grams of carbon and _______ grams of oxygen.

3. How many moles are there in:
a) 1.5 g of NaCl

b) 100 g of potassium nitrate

g) 300 mL of a 2.00 M NaOH solution

4. Find the percentage by mass of nitrogen in aluminum nitrate (19.7%)

5. Calculate the mass of 0.0250 mol of NaF. (1.05 g)

6. Calculate the mass of 1.50 L of argon gas, Ar, at STP. (2.67g)

7. A sample of a chemical was analyzed and found to contain 138 grams of sodium, 36 grams of carbon and 144 grams of oxygen. Determine the simplest formula for the compound. (sodium carbonate)

8. A chemist analyzes a sample of rock from the centre of the earth. It contains 18.61 g of iron and 8.00 g of oxygen. What is the simplest formula for the iron compound in the rock? (iron(III) oxide)

9. Analysis of an organic compound shows that it contains 61.02% carbon, 11.86% hydrogen and 27.12% oxygen. What is the simplest formula of the compound? If the molar mass of the compound is 118.1 g/mol, what is the molecular formula of the compound? (simplest formula is C3H7O, and the molecular formula is C6H14O2).