SCH 4U
Exam Review
Unit 1 - Electrochemistry
1.
If Ag+ (aq) + e-  Ag (s) was chosen as the standard for half-cell potentials, what would be the
half-cell potential for Ni2+ (aq) + 2e-  Ni (s)?
2.
3Ca (s) + 2Al(NO3)3 (aq)  3Ca(NO3)2 (aq) + 2Al (s)
a) Predict the voltage for an electrochemical cell based on the above reaction. Assume
standard conditions.
3.
Use the Eo values to determine if the following reaction is spontaneous or not:
Mg2+ (aq) + Cu (s)  Mg (s) + Cu2+ (aq)
4.
balance the following reaction
a) Zn (s) + NO3- (aq)  Zn2+ (aq) + NH4+ (aq)
b) Cl2 (aq)  Cl- (aq) + ClO3- (aq)
c) MnO4-(aq) + Sn2+ (aq)  Mn2+ (aq) + Sn4+ (aq)
d) I- (aq) + Cr2O72- (aq)  I2 (s) Cr3+ (aq)
(acidic)
(basic)
(acidic)
(basic)
5.
Draw a sketch of the following cell. Calculate the cell voltage:
Zn (s) | Zn2+ || Ag+ | Ag (s)
6.
Consider an electrochemical cell consisting of a nickel half-cell and a tin half-cell.
a) Draw and fully label the above cell.
b) Calculate the cell potential, assuming standard conditions.
c) Calculate the cell potential if a zinc-copper cell was connected to this cell in series.
7.
Explain, with the use of a diagram, how a copper can be plated onto a spoon in an electrolytic
cell.
8.
Consider an electrolytic cell that contains aqueous lithium bromide with carbon electrodes.
a) Write the half-reaction that would occur at the cathode.
b) Write the half-reaction that would occur at the anode.
c) What is the minimum voltage that will need to be applied to this cell?
Unit 2 - Structure and Properties
1.
Write the electron configuration for
a)Sr
b)Ta
c)Gd
d)Br
2. Write the short form electron configuration for 1 a) and 1 b).
3. Draw the energy level diagram for 1 a) and 1 b).
4. How many valence electrons does each of the following have?
a) silicon
b) phosphorus
5. For Li and Be
a) compare the size of the 1st ionization energies.
b) compare the size of the 2nd ionization energies.
c) compare the size of the 3rd ionization energies.
6. Explain why an Al atom forms three bonds instead of one.
e)Cl-
f)Mg2+
7. Draw the Lewis structure for
a) BCl3
b) N2O2
8. Draw the structure and the orbital representation for each of the following:
a) LiH
b) BeCl2
9. Name the shape, draw the structure and state the bond angle for each of the following:
a) CaCl2
b) NH3
c) SF6
10.
What type of bond would hold each of the following pairs of atoms together?
a) N-S
b) Si-I
c) N-Br
d) K-F
11.
Which of the following is the most polar bond?
a) H-I
b) P-I
c) Si-F
12.
What shape is most likely about each N atom in N2F2?
13.
Which of the following has a molecular dipole?
a) LiCl
b) BeF2
c) BI3
d) CaCl2
14.
d) Mg-N
e) NH3
f) SF6
Explain what was wrong with Rutherford’s model of the atom. What changes did Bohr make to the
Atomic Theory?
15.
What are London Dispersion Forces and how do they occur?
16.
What are dipole-dipole forces? What are hydrogen bonds? When do they occur?
17.
Consider two isomers C2F2H2 (cis) and C2F2H2 (trans). Which one would you expect to have a
higher melting point? Which one would dissolve in water?
18.
Define each of the following, give an example and three properties.
a) metallic solid
b)network solid
b) ionic solid
d) molecular solid
Unit 3 – Energy and Reaction Rates
1.
Write each of the following using the ∆H notation. Then state whether the reaction is exothermic or
endothermic.
a. SO2(g) + ½ O2(g)  SO3(g) + 88kJ
b. Ca(OH)2(s) + 64.9kJ  H2O(l) + CaO(s)
c. 6C(s) + 3 H2O(l) + 902kJ  C6H6(l) + 3/2 O2(g)
d. N2(g) + 3H2(g)  2NH3(g) + 92kJ
2.
When heptane, C7H16, undergoes combustion, carbon dioxide and water vapour are formed.
a. Write a balanced chemical equation for this reaction.
b. Given the following information, determine ∆H for the reaction in kJ/mol of C7H16.
C(s) + O2(g)  CO2(g)
∆H = -394 kJ
H2(g) + ½ O2(g)  H2O(g)
∆H = -242 kJ
7C(s) + 8H2(g)  C7H16(l)
∆H = +121 kJ
c. How much energy (in kJ) will be released if 5.00 x 102 g of heptane undergoes combustion?
3.
When 40.8g of potassium hydroxide were dissolved in 1.30 L of water, the temperature of the water
rose 46.0ºC. Calculate ∆H in kJ/mol of potassium hydroxide.
4.
2C(s) + O2(g)  2CO(g)
74.0g of oxygen gas reacts in 20.0 minutes. Calculate the rate of the reaction in moles of carbon
monoxide gas produced per hour.
5.
Sketch a potential energy curve for an exothermic reaction and label it completely.
6.
What effect does a catalyst have on the rate of reaction? Explain.
7.
A + B  products
Trial
1
2
3
4
5
6
[A] (mol/L)
0.001
0.002
0.003
0.004
0.004
0.004
[B] (mol/L)
0.004
0.004
0.004
0.001
0.002
0.003
Rate (mol/sec)
0.003
0.012
0.027
0.002
0.032
0.162
a) Determine a rate law expression for the above reaction.
b) What would happen to the reaction rate if [A] was tripled and [B] was doubled?
Unit 4 - Equilibrium
1.
What effect would the changes below have on this equilibrium?
2SO2 (g) + O2 (g) ↔ 2SO3 (g) + heat
a) [SO2] is increased.
b) [SO3] is decreased.
c) Temperature is increased.
d) Volume of container is increased.
2.
Write the equilibrium expression for the reaction in question 1.
3.
Fe (s) + ½ N2 (g) ↔ FeN (s)
-41.8 kJ
a) Would the tendency towards minimum energy favour the reactants or products?
b) Would the tendency towards maximum entropy favour reactants or products?
4.
A (g) + B (g) ↔ AB (g)
K = 4.0 x 10-2
0.50 mol of A and 0.60 mol of B were placed in a 2.0 L container. Calculate the concentration of
AB at equilibrium.
5.
At 773oC, a mixture of CO (g), H2 (g) and CH3OH was allowed to come to equilibrium. The following
equilibrium concentrations were then measured: [CO] = 0.105 M, [H2] = 0.250 M, [CH3OH] = 0.00261
M. Calculate K for the reaction CO (g) + 2H2 (g) ↔ CH3OH (g)
6.
At a high temperature, 0.500 mol HBr was placed in a 1.00L container and allowed to decompose
according to the reaction
2HBr (g) ↔ H2 (g) + Br2 (g)
At equilibrium, the concentration of Br2 was measured to be 0.130 M. What is the K for the
reaction?
7.
At a certain temperature, the reaction
CO (g) + H2O (g) ↔ CO2 (g) + H2 (g)
has K = 0.400. Exactly 1.00 mol of each reactant was placed in a 100 L vessel and the mixture
was allowed to react. What were the equilibrium concentrations of each gas?
8.
Write the Ksp expression for the dissolving of each of the following in water:
a) Silver chromate
b) Calcium phosphate
9.
Calculate the solubility of cuprous chloride in grams per litre. Ksp = 3.2 x 10-7.
10.
Calculate the solubility, in moles per litre, of calcium sulphate. Ksp = 2.4 x 10-5.
11.
Calculate the Ksp value for copper (II) sulphide if 2.0 x 10-17g of it dissolves in 1.0L of water.
12.
Will a precipitate form if 0.002 moles of lead (II) nitrate is added to 2.0L of 0.001 M sulphuric acid?
Ksp (PbSO4) = 1.3 x 10-8.
13.
Write an equation for the reaction of hydrochloric acid with ammonia. What would be the effect of
adding NH4Cl to the solution?
14.
Calculate the [H3O+] for 0.50 M solution of acetic acid. KA = 1.8 x 10-5.
15.
Calculate the pH for the solution in question 14.
16.
Calculate the [OH-] for the solution in question 14.
Unit 5 – Organic Chemistry
1.
Give the general formulas for
a. Alkanes
b. Alkenes
c. Alkynes
d. Cycloalkanes
2.
Copy and complete the following chart:
Organic
Compound
Functional
Group
Diagram of General Formula
Alkyl Halides
(Organic Halides)
Halogen atom
R–X
Alcohols
Ethers
Aldehydes
Ketones
Carboxylic Acids
Esters
Amines
Amides
Aromatic
Naming Prefix
“Bromo”
“Chloro”
“Fluoro”
“Iodo”
“Nitro”
3.
a.
b.
c.
d.
e.
f.
4.
a.
Draw structural diagrams for:
Butanoic acid
2, 3 – dimethylbutane
2 – ethyl – 1 – pentene
methoxyheptane
3 – pentanone
hexylpropanoate
g.
h.
i.
j.
k.
2 – methyl propanamide
Ethanol
1, 3 – dichlorobenzene
3 – bromo – 4 – fluoro cycloheptene
1 – chloropentane
Give the name for each of the following:
Cl
Cl
|
|
Cl – CHCl
–
–
CH
CH
–
–
CH
CH
–
–
CH
CH
2
22
2
3 – CH3
c.
b.
CH3
CH3 CH3
|
|
|
CH3 – CH – CH2 – CH2 – CH – CH – CH2
|
CH3 – CH2 – CH
d.
NH2 – CH3
e.
f.
O
||
CH3 – C – CH3
CH3 – CH2 – OH
g.
i.
CH3 – O – CH3
O
||
HO – C – CH3
h.
O
||
CH3 – CH2 – C - H
j.
CH3 – CH – CH3
5. Draw the structural isomers of C2H2Br2. Identify the cis and trans isomers.