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Chemistry 12/12AP Review
Thermochemistry
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Law of Conservation of Energy
Distinguish between endothermic and exothermic reactions, enthalpy diagrams
Types of systems
Heat calculations – First Law of Thermodynamics
Determine the heat of reaction –Hess’ Law, Heats of Formation, bond energies
Calorimetry and associated calculations
1. a) Write a balanced chemical equation for the combustion of octane, including the heat of combustion.
The combustion products are CO2(g) and H2O(l). (Value: 3)
b) Sketch an enthalpy diagram for the combustion of octane. Your sketch should indicate the ΔH comb
value.
2. A 0.50 g sample of NH4NO3(s) is dissolved in 35.0 g of water in a Styrofoam cup. The temperature of
the solution drops from 22.7ºC to 21.6ºC.
i) Is the process exothermic or endothermic?
ii) What is the heat of solution of ammonium nitrate, expressed in kJ/mol NH 4NO3?
3. Given the thermochemical equations:
2B(s) + 3/2 O2(g)  B2O3(s)
ΔH = -1273 kJ
B2H6(g) + 3O2(g)  B2O3(s) + 3H2O(g) ΔH = -2035 kJ
H2(g) + ½O2(g) H2O(g)
ΔH = -242 kJ
Calculate the ΔH for the reaction:
2B(s) + 3H2(g)  B2H6(g)
4. Using the data below, calculate the amount of energy required to melt a 16.2 g aluminum pop can.
Assume that the aluminum is at an initial temperature of 20°C.
• Molar heat of fusion of Al(s) = 10.7 kJ/mol
• Melting point of Al(s) = 660°C
• Heat capacity of Al(s) = 0.900 J/g°C
5. Use bond energies to calculate the enthalpy of formation of CH3CH2OH(g) from gaseous atoms. Show
your calculations.
6. 0.844 g of formaldehyde (methanal), HCHO, reacted with oxygen in a styrofoam calorimeter to produce
carbon dioxide and water. The mass of water in the calorimeter was 150.0 g and the temperature
increased from 21.8°C to 45.1°C. Calculate the heat of reaction per mole of formaldehyde burned.
7. Calculate the ΔH of the following reaction; Na(s) + ½Cl2(g)  NaCl(s)
2Na(s) + 2HCl(g)  2NaCl(s) + H2(g)
ΔH = –637.4 kJ
H2(g) + Cl2(g)  2HCl(g)
ΔH = –184.6 kJ
8. Draw a heating curve for ammonium chloride which has a melting point of 330oC and a boiling point of
520oC. Properly label the axes.
*9. Calculate the standard-state entropy of reaction for the following reactions and explain the sign of S
for each reaction.
(a) Hg(l)
Hg(g)
(b) 2NO2(g)
N2O4(g)
(c) N2(g) + O2(g)
2NO(g)
*10. Calculate ΔG at 25°C for the following reaction, by first calculating ΔH and ΔS. Once you've found
ΔH and ΔS, solve for ΔG using the formula:
ΔG = ΔH - T ΔS
Also - will this reaction be spontaneous at this temperature?
CH3CO2H (l) + 2 O2 (g) → 2 CO2 (g) + 2 H2O (g)
Solutions
 Properties, types of solutions
 Solubility Curves
 Concentration calculations, including dilution
 Solution preparation
 Predicting products of reactions – net ionic equations
 *Ksp
11. a) Calculate the concentration of glucose (C6H12O6) in a solution in which 10.0 g of glucose is
dissolved in water and diluted to a final volume of 200.0 mL.
b) Determine the volume of commercial grade nitric acid (15.8 mol/L HNO 3) that is needed to prepare
100.0 mL of 3.00 mol/L HNO3.
c) Rubbing alcohol is commonly used as an antiseptic for small cuts. It is sold as a 23% (v/v) solution of
isopropyl alcohol in water. What volume of isopropyl alcohol is used to make 570mL of solution?
12. When silver nitrate is added to sodium chromate, a brick-red precipitate forms. Calculate the mass of
precipitate that forms when 50.0 mL of 0.100 mol/L silver nitrate is added to excess sodium chromate.
13. For each of the following:
i) determine the products formed and write a balanced chemical equation
ii) write a total ionic equation
iii) write a net ionic equation if possible
A) Na2 SO4 (aq)  Fe( NO3 ) 2 (aq) 
B)
CoCl2 (aq)  NaNO3 (aq) 
C)
BaCl 2 (aq)  ( NH 4 ) 2 CO3 (aq) 
D)
MgS (aq)  Na2 CO3 (aq) 
*14. Solid barium phosphate dissolves into its respective ions at 25oC. Given that [Ba2+] = 1.5 x 10-2 M,
find the value of Ksp.
Kinetics and Equilibrium
 Collision Theory
 Rates of reaction, Reaction Mechanisms
 Factors affecting rate of reaction
 Keq
 Reaction Quotient
 Le Chatelier’s Principle
15. Based on the following reaction profile,
a) How many intermediates are formed in the reaction A  C?
b) How many activated complexes are there?
c) Which step is the fastest?
d) Is the reaction exothermic or endothermic?
Energy
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2
*16. For the chemical reaction: 2 IO3  5SO3  2 H
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 I 2  5SO42  H 2 O , the following data were
recorded:
Experiment
[IO3-]
[SO32-]
1
0.001
0.01
2
0.001
0.02
3
0.002
0.01
4
0.001
0.01
Calculate the rate law equation and the rate constant.
[H+]
0.1
0.1
0.1
0.2
Initial rate
0.0001
0.0001
0.0002
0.0004
17. Calculate the equilibrium constant Kc at 25 oC for the reaction
2 NOCl(g) ↔ 2 NO(g) + Cl2(g)
using the following information. In one experiment 2.00 mol of NOCl is placed in a 1.00 -L flask, and the
concentration of NO after equilibrium is achieved is 0.66 mol/L.
18. For the gas phase reaction
H2(g) + I2(g)  2 HI(g)
Kc = 50.3 at 731 K. 0.100 mol of HI is introduced to a 0.500L container, and then the temperature is
raised to 731 K. Calculate the concentration of each when the system is at equilibrium.
19. A mixture of 9.22 moles of A, 10.11 moles of B, and 27.83 moles of C is placed in a one-liter
container at a certain temperature. The reaction is allowed to reach equilibrium. At equilibrium the number
of moles of B is 18.32. Calculate the equilibrium constant for the reaction: A (g) + 2 B (g) ↔ 3 C (g)
20. Use the following equation for parts i) and ii)
4 NH 3( g )  5O2( g )  4 NO( g )  6 H 2 O( g )
i) How will the following changes affect the above equilibrium? (shift right, shift left, no change)
a) decrease [O2] ________________________
b) increase [H2O] _________________
c) increase [NH3] ________________________
d) increase volume ___________________
e) decrease the pressure __________________________
ii) How will the following changes affect the concentration of O 2 above? (increase, decrease, no change)
a) decreasing [NH3] will cause [O2] to ________________________________________
b) increasing [NO] will cause [O2] to ________________________________________
c) increasing [H2O] will cause [O2] to ________________________________________
21. Hydrogen can be produced industrially by a steam-hydrocarbon process. One step in the process
involves the following equilibrium system:
H 2 O( g )  CH 4( g )  CO( g )  3H 2( g )
(a) A mixture containing 0.725 mol H2O(g) and 0.2 mol of CH4 is put in a 1.50L container at 725oC.
When equilibrium is reached there is 0.075 mol of CO. What is the value of Keq at 725 oC?
(b) A chemist contacted a company who produced H2 using the above process. The chemist
claimed he discovered a catalyst that would increase the profits by shifting the equilibrium to
produce more H2 under the same conditions. Should the company invest in this technology?
Why or why not?
Acids and Bases
 Properties/Characteristics
 indicators
 pH/pOH calculations
 titrations
 Ka/Kb
22. Summarize (in your own words) each acid-base theory (Arrhenius, Brønsted-Lowry).
23. Determine the pH/pOH/[H+] and [OH-] of each of the following:
a) 0.05M H2SO4
b) 0.08M HNO3
c) 0.15M NaOH
d) 4.3x10-3M KOH
24. Predict whether or not each of the following oxides would be acidic or basic.
a) Cl2O
b) MgO
c) K2O
d) SO3
25. If the pH of a solution is 6.8, what is the colour of each of the following indicators in this solution?
a) methyl red
b) chlorophenol red
c) bromothymol blue
d) phenolphthalein
e) methyl orange
26. 17.50 mL of hydrochloric acid was neutralized by 23.50 mL of 0.250 M sodium hydroxide. Calculate
the concentration of the hydrochloric acid
27. Explain the difference between the equivalence point and the end point of a titration.
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28. Codeine has a Kb of 1.73 x 10 . Calculate the pH of a 0.020 mol/L codeine solution.
29. A 0.10 mol/L solution of lactic acid, found in sour mild, has a pH of 2.43. Calculate the Ka of lactic
acid.
*30. Calculate the pH of the solution that results when 40.0 mL of 0.100 M NH 3 is: (weak base)
a) Diluted with 20.0 mL of distilled water.
b) Mixed with 20.0 mL of 0.200 M HCl solution
Electrochemistry
 Oxidation Numbers
 Writing Half-reactions
 Balancing in acidic and basic conditions
 Cell Potential and parts of an electrochemical cell
31. Assign oxidation numbers to each element in the following:
a) H2SO4
b) CO32c) MnO4d) O3 e) AsO432. Balance the following equations using the half-reaction method:
a) NiO2  S 2 O3  Ni(OH ) 2  SO3 (acidic )
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f) Pb(NO3)4
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b) Ag 2 S  CN  O2  S  Ag (CN ) 2 (basic)
33. a) Calculate the cell potentials for each of the following reactions.
b. Indicate if the reaction is spontaneous or non-spontaneous.
(i)
Cr  3 Ag   Cr 3  3 Ag
(ii)
2Cr 3  7 H 2 O  3Pb 2  Cr2 O7
2
 14H   3Pb
34. 21.9mL of 0.04705M KMnO4 is used to neutralize a 1.43g of H2O2. Calculate the % by mass of H2O2
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in the solution. ( 2 MnO4  5 H 2 O2  6 H

 5O2  2Mn 2  8H 2 O )
35 Refer to this diagram of an electrochemical cell to answer i) to v).
i) What is the oxidation half-reaction?
ii) What is the reduction half-reaction?
iii) What is the balanced net oxidation-reduction
reaction?
iv) Which ions migrate into the porous cup?
v) In which direction do the electrons flow?