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Unit #1 Structure and Properties of Matter (p160-291)
 Models of the Atoms
 Atomic Spectrum
 Quantum Mechanical Model of the Atom
 Quantum Numbers
 Energy Level Diagrams (for atoms, cations, anions)
 Electron Configurations (2)
 Electron Configuration Periodic Table Trends
 Magnetism
 Intermolecular and Intramolecular Bonding
 Lewis Structures and Structural Formulas
 Polarity of Molecules
 VSEPR Theory (shapes of molecules)
 The Structure and Properties of Solids (ionic crystals, metallic crystals, molecular crystals, covalent network
crystals)
Practice:
1. Draw lewis diagrams, name the shape, determine the bond angles, determine the polarity and determine the type of
bonding percent for each of the following: NH3, CH4, BeF2, SO22- and MgO.
2. Draw the energy level diagrams, and write both electron configurations for oxygen atom, nickel 3+ ion, and
3.
molybdenum atom.
Identify the following atoms or ions from their electron configurations:
A) A: 1s22s22p63s23p64s13d10
B) B2+: 1s22s22p63s23p63d5
C) C3+: 1s22s22p63s23p64s23d104p64d5
4. What causes line spectra?
5. State a similarity and difference between the terms “orbit” and “orbital”. Which atomic models that you
have studied would use each of these terms?.
6. Given the following electron configurations, state the values of the four quantum numbers for the electron
indicated.
1s2 2s2 2p5 (second e– in p orbital)
B)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 (tenth e– in d orbital)
n =
n =
l =
l =
ml =
ml =
ms =
ms =
7.
Distinguish between the terms "photon" and "quantum".
8.
Explain why CH3F is a polar molecule while CF4 is not.
9.
Explain intermolecular forces and identify what type of molecules they will be found in.
10. Make a chart of the properties of solids.
Unit #2 Organic Chemistry (p6-116)
Important Summaries; Organic Families (p93), Reactions (p83)
 Nomenclature (names and structures of hydrocarbons, alcohols, aldehydes, ketones, carboxylic acids, esters, ethers,
amines, amides, aromatics)
 Chemical Reactions (addition, substitution, elimination, oxidation, hydrolysis, esterification)
 Physical Properties (solubility, odour, melting point, boiling point, unsaturation, hardness)
 Polymers
 Polymerization (addition and condensation)
 Applications
Practice:
1. Make a summary table of organic functional groups using the heading: Name, Functional Group, IUPAC
Suffix
2. Draw formulas for the following:
(a) 3-ethylhexanoic acid
(b) 2-methyl-1-butene
(c) 3,5-dibromo-4-pentanol
(d) 2-butylpropanoate
(e) propanone
(f) cis-3,4-difluoro-1-amino-3-octene
(g) propoxybutane
(h) 1,3-pentandiol
(i) butanamide
(j) cyclohexanol
(k) 1,3-dimethylbenzene
(l) 2-methylbutanal
(m) 2-aminoethanoic acid
3. Write the name and formula for:
(a) a five carbon aldehyde
(b) a four carbon ketone
(c) a three carbon acid
(d) a three carbon alkyne
(e) an amide made from a three carbon acid and a three carbon secondary amine
(f) an ether with four carbons
(g) a tertiary alcohol with four carbons
4. Show equations for the following syntheses
(a) 2-propanol from an alkene
(b) ethanoic acid from an alcohol
(c) 1-butyl propanoate from an acid and an alcohol
(d) ethyl amine from an alkene
(e) propanal from and alcohol
(f) 1,2-dichloroethane from ethene
5. Complete the following reaction equations & name the product in each case.
i) CH3–CH2–CH3
+
Cl2 →
(halogenation via substitution)
ii) CH3–(CH2)4–CH=CH2
+ Cl2
→
(halogenation via addition)
iii) CH3–CH(CH3)–CH=CH2
+
HCl →
(Markovnikov addition)
iv) CH3–CH2–CH2OH
→
(oxidation of primary alcohol)
v) CH3–CH2–CH2OH +
HOOC–CH2–CH2–CH3
→
(Esterification)
vi) C6H5–CH2-CHOH-CH3
benzene
→
(oxidation of secondary alcohol)
ring
vii)
6. a) Draw the formula for the polymer made from a two carbon double acid and a three carbon double
alcohol.
b) Draw the structural formula for the polymer formed from monomer units of 1,1-fluoro ethene.
7. Order the substance form the lowest boiling point to the highest for the following compounds: propanol, propanoic acid, prpane, and propanal. Explain the order using bonding.
Unit #3 Energy Changes and Rates of Reaction (p296-417)
 Thermochemistry Terms (exothermic, endothermic, open and closed system, heat, temperature, enthalpy)
 Energy Changes (physical, chemical, nuclear)
 Heat Transfer Calculations (q= mcΔT)
 Molar Enthalpy Calculations
 Thermochemical Equations
 Potential Energy Diagrams (activation energy, activated complex, transition state, intermediates, ΔH, catalyst)
 Hess’s Law Calculations
 Standard Enthalpy of Formation Calculations
 Energy Production Methods (solar, hydroelectric, geothermal, tidal, nuclear, fossil fuel, etc.)
 Rates of Reaction Calculations
 Factors affecting Reaction Rates (concentration, nature of reactants, temperature, catalyst, surface area)
 Rate Laws Calculations and Order of the Reaction
 Collision Theory
 Mechanisms
Practice:
1. In ____________________ changes, energy is released from a system to the surroundings and in
____________________ changes, energy is absorbed by a system.
2. The combustion of methanol is shown by the following equation:
i. Given the data which follows:
a. Find the heat of reaction for the equation above.
b. State the molar heat of combustion of methanol.
c. State whether the reaction is endothermic or exothermic.
ii. What mass of water could be heated from 20.00 C to 35.00 C by the burning of 2.57 mol of methanol?
(
)
3. Calculate the Hf for the target reaction: 3 Fe2O3 (s)  2 Fe3O4 (s) + ½ O2 (g) given the data below. Explain
what you needed to do to each equation to get your answer.
2 Fe (s) + 3/2 O2 (g)  Fe2O3 (s)
Hf = -823.41 kJ
3 Fe (s) + 2 O2 (g)  Fe3O4 (s)
Hf = -1120.48 kJ
4. What is the Ho in kilojoules for the combustion of 1 mol of ethanol, C2H5OH(l), to form gaseous carbon dioxide and
gaseous water?
5. If an iron ring with a mass of 5.5 g changes temperature from 25.0ºC to 28.0ºC, how much energy has it absorbed?
6. Consider the following reaction pathway in which S2O82- (aq) reacts with I- (aq) ions
Cu2+ (aq)
+
I- (aq)
Cu+ (aq) +
I (aq)
—>
2 I (aq)
I2 (aq)
—>
Cu+ (aq)
+
S2O82- (aq)
CuSO4+ (aq)
+
SO42- (aq)
—>
+
+
2+
Cu (aq)
+
CuSO4 (aq)
2 Cu (aq)
+
SO42- (aq)
—>
a) Identify the catalyst and the reaction intermediates in this reaction.
CATALYST:
[FAST]
[FAST]
[SLOW]
[FAST]
INTERMEDIATES:
b) Identify the reactants and products, by writing the overall reaction equation.
c) What is the rate law equation for the above mechanism?
d) Explain what effect increasing [I- (aq)] and [S2O82- (aq)] would have on the overall rate.
7. The initial rate of the reaction: BrO3-(aq) + 5 Br-(aq) + 8 H+(aq)  3 Br2(l) + H2O(l)
has been measured at the reactant concentrations shown (in mol/L)
Experiment
[BrO3-]
[Br-]
[H+]
Initial rate (mol/(L·s))
1
0.10
0.10
0.10
8.0  10-4
2
0.20
0.10
0.10
1.6  10-3
3
0.10
0.20
0.10
1.6  10-3
4
0.10
0.10
0.20
3.2  10-3
a) Determine the rate law.
b) According to these results, what would be the initial rate (in mol/(L·s)) if all three concentrations are:
[BrO3-]=[Br-]=[H+]=0.20 mol/L?
c) State the overall order of the reaction
8. Sketch a potential energy diagram for a system in which Ea = + 50kJ, and H= 20 kJ. Label the axes, reactants,
products, the activation energy and the activated complex.
9. What are the five factors the can affect the rate of a reaction? Explain each one on a molecular level.
Unit #4 Chemical Systems and Equilibrium (p422-493, p526-645)
 Dynamic Equilibrium (explain, graphs)
 Equilibrium Constant (K) Calculations
 Le Chatelier’s Principle
 The Reaction Quotient (Q) Calculations
 Solubility Product (Ksp)
 Predicting a whether there is a Precipitate
 Common Ion Effect Calculations
 Bronsted-Lowry Acid-Base Theory
 Ion Product Constant for Water (Kw) Calculations
 pH, pOH, [H+], [OH-] Calculations for Strong Acids and Bases
 Weak Acids and Bases Percent Ionization
 Acid (Ka) and Base (Kb) Ionization Constant Calculations
 Acid-Base Titration Calculations (strong acid and strong base, weak acid or base and a strong acid or a base)
 Titration Curves
 Selecting an Indicator (endpoint = equivalence point)
 Buffers (what is it, how does it work, applications)
Practice:
1. If the solubility of AgBr is 8.8  10-7 mol/L, what is its Ksp?
2.
If 145 mL of a 0.0078 mol/L solution of CuNO3 was mixed with 195 mL of a 1.48  10-3 mol/L solution of KI,
would a precipitate form? Calculate the ion product for the potential precipitate. The Ksp of CuI is
1  10-12.
3.
If the Kb of a 0.34 mol/L solution of weak base is 1.3  10-12 what is its percent ionization?
4.
What is the pH of a 1.47 mol/L solution of HCN(aq) if its Ka = 3.5  10-11?
5.
28 mL of standardized 0.43 mol/L NaOH is titrated with 23 mL of 0.36 mol/L Hydrochloric acid. Calculate the pH
of the solution. Sketch a potential energy diagram. State an indicator.
6. Use Le Chatelier’s Principle to explain in which direction the equilibrium will shift if the following
stresses are applied to the following system:
CH4(g) + 3O2(g) ⇌CO2(g) + H2O(g) H = 56 kJ
i)
ii)
iii)
iv)
increasing the volume of the system
decreasing the temperature of the system
adding a catalyst
increasing the concentration of the products
7.
Consider the equilibrium below:
If 1.6 mol of HI was placed in a 1.0 L container and allowed to reach equilibrium, what would the equilibrium
concentrations be for H2(g), I2(g) and HI(g) if the K = 36?
2HI(g) <=====> H2(g) + I2(g)
8.
If the pH of an acid solution at 25oC is 4.53, what is the pOH and the [H1+], [OH1-] in mol/L?
9.
The system F2(g) + Cl2(g)  2 ClF(g) is at equilibrium when the concentrations of the materials are as follows:
[F2] = 0.40 mol/L, [Cl2] = 0.20 mol/L and [ClF] =1.00 mol/L.
a) Calculate the equilibrium constant for the reaction.
10. What is a buffer? How does it work? What are some applications of a buffer?
Unit #5 Electrochemistry (p651-769)
 Redox Reactions (oxidation, reduction, reducing agent, oxidizing agent)
 Oxidation numbers
 Balancing Redox Reactions (oxidation, half-reaction method)
 Predicting Redox Reactions (spontaneous or non-spontaneous)
 Electrochemical Cells (drawing galvanic cells and calculating Eocell)
 Batteries
 Corrosion (Chemistry, prevention)
 Electrolytic Cells
 Applications of Electrolytic Cells (production of Al, Na, refining of metals, electroplating)
 Faraday’s Calculations
Practice:
1. What is the oxidation number of Cl in
?
2. Will a reaction take place if a strip of zinc metal is placed in a solution of lead (II) sulfate? Justify your answer.
3. Calculate
for the following equation:
Decide whether the equation represents a spontaneous or a non-spontaneous reaction at standard conditions.
4. When sodium sulfite solution is added to acidified potassium permanganate solution, the purple permanganate
solution turns colourless. The skeleton net ionic half-cell reactions are:
reduction half-cell reaction:
oxidation half-cell reaction:
Construct and balance the equation of the overall redox reaction (in acidic medium) using the half-reaction method.
5. An electrochemical cell consists of silver and copper electrodes immersed in 1.0 mol solutions of
Silver nitrate and copper (II) nitrate respectively. Potassium chloride solution is used as a salt bridge.
a) Write the half cell reactions that occur at each electrode.
b) Write the cell reaction and calculate the cell potential.
c) Draw a labeled diagram of the cell. Include labeling the anode, cathode and direction of flow of
electrons.
6. Balance the following using oxidation numbers and half reactions.
i)
Mn2+ + HBiO3 +
H+ → Bi3+ + MnO4– +
ii)
H2O
Cr2(SO4)3 + Na2CO3 + KNO3 → Na2CrO4 + KNO2 + Na2SO4 + CO2
7. What is corrosion? Explain it in terms of the following terms: oxidation, reduction, reducing agents, oxidizing
agents and redox. How is corrosion prevented?
8. In a school chemistry experiment, an electrolytic cell containing aqueous copper(II)suphate is constructed to
electroplate electroplate an iron bolt and copper metal.
a) At which electrode should the bolt be attached? Why/
b) What are the reaction and products at the other electrode running for 30 minutes
c) Determine the mass of copper plated out 1.5A current running for 30 min.