High Level High Quality High Efficiency
One-on-One or Group Instruction
YU: CHEM 2020, CHEM 3020; UofT: CHM138H, CHM247S, CHM249S, CHM348F, CHM343S, CHM342F; McM: 2OA3, 2OB3, 2OC3, 2OD3, 3D03, 4D03; UWO: CHEM2213a, CHEM2223b.
* YU: York University;
UofT: University of Toronto;
McM: McMaster University;
UWO: University of Western Ontario
© 2009 - 2013 MPC Academy. All Rights Reserved. For registration or inquiry, please email us at email@example.com or call us at (416)879-3919
Chemical System and Equilibrium
1.The oxidation of sulfur dioxide to sulfur
trioxide is an important reaction. At 1000 K,
the value of Kc is 3.6E-3.
2SO2(g) + O2(g) = 2SO3(g)
A closed flask originally contains 1.7 mol/L
SO2(g) and 1.7 mol/L O2(g). What is [SO3] at equilibrium when the reaction vessel is
maintained at 1000 K?
2. Ethanol and propanoic acid react to form the ester ethyl propanoate, which has the odour of bananas.
CH3CH2OH(l) + CH3CH2COOH(l) = CH3CH2COOCH2CH3 + H2O(l)
At 50C, Kc for this reaction is 7.5. If 30.0 g of ethanol is mixed with 40.0 g of propanoic acid, what mass of ethyl propanoate will be present in the equilibrium mixture at 50C?
Hint: Calculate the initial amounts of the
reactants. Then solve the equilibrium equation using amounts instead of concentrations. The volume of the mixture does not affect the calculation.
3. 0.50 mol of CO(g) and 0.50 mol of water(g) are placed in a 10 L container at 700 K.
The following reaction occurs.
CO(g) + H2O(g) = H2(g) + CO2(g)
Kc = 8.3
What is the concentration of each gas that is present at equilibrium?
4. Sulfur trioxide gas reacts with gaseous
hydrogen fluoride to produce gaseous sulfur
hexafluoride and water vapour. The value of Kc is 6.3E-3.
(a) Write a balanced chemical equation for
(b) 2.9 mol of sulfur trioxide is mixed with
9.1 mol of hydrogen fluoride in a 4.7 L flask.
Set up an equation to determine the equilibrium concentration of sodium hexafluoride.
(c) Explain why you are likely unable to solve
1. The oxidation number of carbon in sodium carbonate is__.
2. Why rusting is a redox reaction?
3. What are primary batteries?
4. The sum of the oxidation numbers of all the elements in a compound is ___.
5. What is the reaction occurring in a lead-acid battery when recharged? Why is the density of the electrolyte solution in the battery greatest when it is fully recharged?
6. Does the fact that you can assign oxidation numbers of +1 to hydrogen and -2 to oxygen in
water mean that water is an ionic substance? Explain.
7. What are the products from the electrolysis of a 1 mol/L solution of hydrochloric acid.
8. Suppose you decide to protect a piece of iron from rusting, by covering it with a layer of lead. (a) Would the iron rust if the lead layer completely covered the iron? Explain. (b) Would the iron rust if the lead layer partially covered the iron? Explain.
9. The percent by mass of tin in an alloy can be found by dissolving a sample of the alloy in an acid to form aqueous tin(II) ions. Titrating with cerium(IV) ions produces aqueous tin(IV) and cerium(III) ions. A 1.475 g sample of an alloy was dissolved in an acid and reacted completely with 24.38 mL of a 0.2113 mol/L cerium(IV) nitrate solution. Calculate the percent by mass of tin in the alloy.
10. In a galvanic cell, the mass of the magnesium anode decreased by 3.38 g while the cell produced electricity. (a) Calculate the quantity of electricity produced by the cell. (b) If the constant current flowing through the external circuit was 100 mA, for how many hours did the cell produce electricity?
1. Draw structures according names:
h) methyl ethanoate
i) methyl propyl ether
2. What is the difference between a saturated and unsaturated hydrocarbon? Which is benzene?
5. Identify each pair as structural isomers, geometric isomers, or neither
a) cyclopentane, pentane
b) 1,1-dichloroethene, trans-1,2-dichloroethene
c) hexanoic acid, propyl propanoate
d) cis-1,2-dichlorocyclopentane, trans-1,2-dichlorocyclopentane
6. C2H6O exists as two structural isomers. Draw each isomer and name each. Predict the relative boiling point and solubility in water of the two isomers. Explain.
7.Indicate if the functional group contains a cabonyl, carboxyl or hydroxyl group:
d) carboxylic acid
e) carboxyl group
f) carbonyl group
g) hydroxyl group
8.Indicate how the following molecule can be made by drawing the reaction, naming all products and reactants, and naming the type of reaction that is occurring.
a) 1,2-dichlorocyclopentane (from a hydrocarbon)
b) 1,1-dichlorocyclopentane (from a hydrocarbon)
c) A condensation reaction to make butyl methanoate
9. Using bromine aqueous solution, how can you easily distinguish between pentane, 1-pentene, and 1-pentyne?
10. What is wrong with the following names?
1. Discuss the bonding that is present in NaClO(s). Indicate ionic or covalent bond in the molecule
Determine the molecular shape of Is it a polar molecule?
2. Use VSEPR theory to determine molecular shape of PF2Cl, PCl5, SOCl2, CCl4, XeF6, PF2Cl3, CH3Cl, and CHCl3.
3. What types of intermolecular forces must be
broken to melt solid samples of the following?
(a) NH3 (b) NaI (c) Fe (d) CH4
4. In which compound, H2O or in NH3, will the
hydrogen bonding be stronger? Explain.
5. Draw a Lewis structure for KBrO.
6. In what cases is the name of the molecular
shape the same as the name of the electron
7. How can a molecule with polar covalent bonds
not be a polar molecule?
8. Compare the molecules SF4 and SiF4 with
respect to molecular shape and molecular
9. Discuss the intermolecular and intramolecular
forces in CH3OH and CH3SH. Based upon the
bonding between molecules, which of these two
compounds would have a lower boiling point?
10. Explain the significance, in terms of absorption or emission of photons, of the following statements.
(a) An electron moves from n = 2 to n = 5.
(b) An electron moves from n = 4 to n = 1.
(c) An electron moves from n = 4 to n = 3.
11. Arrange the following in order from lowest energy to highest energy, and justify your sequence: n = 7, n = 2, n = 5, n = 4, n = 1.
12. Explain why it is not possible to measure the
size of an atom directly.
13. Why are there no p block elements in period 1
of the periodic table?
14. At one time, scandium (Z = 21) was placed
in the same group of the periodic table as
aluminum (Z = 13).
(a) Use the aufbau principle to write the ground
state electron configurations for atoms of
these two elements.
(b) What ionic charge would each have in
common? How might this have led chemists
to place them originally in the same group?
(c) Identify and explain all the evidence you have.
Energy Changes and Rates of Reactions
1. Consider the following thermochemical equation. 25.9 kJ + 12H2(g) + 12I2(g) = HI(g)
(a) What is the enthalpy change for this reaction?
(b) How much energy is needed for the reaction of 2.0 g of iodine with excess hydrogen?
(c) Draw and label an enthalpy diagram that corresponds to the given thermochemical equation.
2. Dilution of concentrated sulfuric acid is extremely exothermic. Design an experiment to measure the enthalpy change (in kJ/mol). Assume that you have access to any
equipment in your school's chemistry laboratory.
(a) What equipment and chemicals do you need?
(b) Write a step-by-step procedure.
(c) Set up an appropriate data table.
(d) State any information that you need.
3. A chemist mixes 100.0 mL of 0.150 mol/L potassium hydroxide with 100.0 mL of 0.050 mol/L phosphoric acid in a constant-pressure
calorimeter. The temperature of the reactants is 21.01 C. The temperature of the products is 21.85C.
(a) Write a thermochemical equation for the reaction.
(b) Explain why a bomb calorimeter may not provide accurate results for determining the enthalpy of a reaction.
4. Use the following equation to answer the questions below.
CH3OH(l) + 1.5O2(g) = CO2(g) + 2H2O(g)
(a) Calculate the enthalpy change of the complete combustion of 2.0 moles of methanol, using enthalpies of formation.
(b) How much energy is released when 125 g of methanol undergoes complete combustion?
5. A 10.0 g sample of pure acetic acid, CH3CO2H, is completely burned. The heat released warms 2.50 L of water from 23.3 C to 39.6 C. Assuming that no heat was lost to the calorimeter, what is the enthalpy change of the complete combustion of acetic acid? Express your answer in units of kJ/g and kJ/mol.
6. Magnesium metal reacts with hydrochloric acid to produce magnesium chloride and hydrogen gas.
Mg(s) + 2HCl(aq) = MgCl2(aq) + H2(g)
Over an interval of 1.00 s, the mass of Mg(s) changes by -0.011 g.
(a) What is the corresponding rate of consumption of HCl(aq) (in mol/s)?
(b) Calculate the corresponding rate of production of H2(g) (in L/s)
7. The rate of decomposition of hydrogen peroxide was studied at a particular temperature. H2O2(aq) = H2O (l) + 12O2(g)
(a) The initial concentration of hydrogen
peroxide was 0.200 mol/L. 10.0 s later,
it was measured to be 0.196 mol/L. What
was the initial rate of the reaction?
(b) 0.500 L of hydrogen peroxide solution was
used for the experiment. What mass was lost
as O2 bubbled out of solution in this initial
10.0 s interval?
This course enables students to deepen their understanding of chemistry through instructing the topics listed below with explanation of typical question examples. Each session will have homework and check/discussion of the questions. Students will develop problem-solving skills as they investigate chemical reactions, at the same time refining their ability to communicate scientific information. About 400 slides and hundreds of practice questions and answers are available.
1) Organic Chemistry: Bonding and shape of Organic Molecules. Nomenclature of organic compounds using the IUPAC system and common name. Chemical reactions of various organic compounds.
2) Structure and Properties: Quantum mechanical theory and how types of chemical bonding account for the properties of ionic, molecular, covalent network, and metallic substances. Products and technologies whose development has depended on understanding molecular structure, and technologies that have advanced the knowledge of atomic and molecular theory. The properties of solids and liquids, and the shape of simple molecules.
3) Energy Changes and Rates of Reaction: The energy transformations and kinetics of chemical changes. Determine energy changes for physical and chemical processes and rates of reaction, using experimental data and calculations.
4) Chemical System and Equilibrium: Le Chatelier's principle and solution equilibria. Investigate the behaviour of different equilibrium systems, and solve problems involving the law of chemical equilibrium. The importance of chemical equilibrium in various systems, including ecological, biological, and technological systems.
5) Electrochemistry: Fundamental concepts related to oxidation-reduction and the interconversion of chemical and electrical energy. The functioning of simple galvanic and electrolytic cells and their description by equations; Quantitative problems related to electrolysis and their solution. Some uses of batteries and fuel cells and the importance of electrochemical technology to the production and protection of metals.