WRITE BALANCED CHEMICAL EQUATIONS AND LABEL EACH WITH ONE OF THE “TYPES
OF REACTIONS” DESCRIBED IN THE TEXT :
1. Na
+ O2
Na2O
P4 + Cl2
2. PCl3
3. CaCl2 + AlF3
CaF2 + AlCl3
WRITE MOLECULAR FORMULAS FOR THESE COMPOUNDS:
4. pentacarbontetrahydrogenhexabrominemonochloride
5. californium (III) carbonate
DRAW A STRUCTURAL FORMULA FOR:
6. aluminum hydroxide
WRITE BALANCED EQUATIONS FOR THESE:
7. the decomposition of liquid iron (III) bromide
8. a single replacement reaction between elemental sodium and water
9. the formation of a precipitate by mixing of lead nitrate and potassium chloride solutions
CALCULATE TO THE PROPER SIGNIFICANT FIGURES:
10. 6.00 x 4.000=
11. 25 x 25.0=
12. 1000.0 x .010 x 10.0=
13. What is the density of a mysterious blue liquid given the data below?
MASS OF EMPTY FLASK
56.68 g
INITIAL PIPET READING
1.24 ml
FINAL PIPET READING
9.12 ml
MASS OF FLASK & LIQUID
72.89 g
MASS OF EMPTY PIPET
6.02 g
CALCULATE TO 3 SIGNIFICANT FIGURES:
14. What is the molar mass of potassium nitrate?
15. What is the mass of 2.00 moles of aluminum chloride?
CALCULATE THESE:
16. The mass of calcium oxide from the reaction of one mole of calcium and excess oxygen gas.
17. The volume of hydrogen bromide gas required to react completely with 2.5 moles of aluminum.
18. The number of chlorine molecules produced when 1.0 grams of aluminum chloride decompose.
19. The mass of solid product from the double replacement reaction between 3.45 grams of silver
nitrate and a solution of excess sodium sulfide.
ANSWER THE FOLLOWING ON THE ATOM AND NUCLEAR CHEMISTRY:
20. Imagine that you are James Chadwick in 1932 and are publishing a chemistry textbook for high
school students. How would you sketch atomic models of two ions of iron that are different
isotopes? You do not have enough space to show each particle, but would be able to indicate the
numbers of each in your sketches.
21. Write nuclear disintegration equations showing the atomic and mass numbers as symbolically
shown in the text:
a. radium-226 decaying by alpha emission
b. potassium-40 decaying by beta emission
22. How much energy in kilojoules is released when 1 mole of oxygen-14 decays by b+ emission to
form nitrogen-14? The molar masses are: oxygen-14 = 14.0086g, b+ = 0.0005 g, and nitrogen-14
= 14.0031g
23. The half-life of chromium-51 is 28 days. If a sample contains 510 grams, how much chromium
would remain after 56 days? How much of the original would remain after exactly 1 year?
24. In terms of unity and diversity, please describe fission and fusion.
25. A photon of red light has a a wavelength of 6.56 x 10-7 m. How much energy does that photon
carry? [HINT: Planck’s constant is 6.6262 x 10-34 Joules/sec.
26. Draw a transverse wave and label the following:
a. Amplitude
b. Wavelength
c. Crest
d. Trough
e. Node
27. List each and every bond in each of the following chemicals:
a. solid magnesium fluoride
b. liquid methanol
28. Indicate the chemical having the higher boiling point in each pair and explain why in terms of
bonding:
a. liquid H2 or liquid H2O
b. liquid NH3 or liquid PH3
29. How is it possible for a molecule to be nonpolar when its individual bonds are polar?
30. Please write the electron configuration, orbital diagram, and noble gas configuration for the
chromium atom and the chromium (I) ion.
31. Which atom has a larger atomic radius, chlorine or bromine?
32. Which has a larger atomic radius chlorine or magnesium?
33. What element has the lowest first ionization energy?
34. Please list the names of groups 1, 2, 11, 15, 16, 17, and 18.
35.The chemical analysis of citric acid shows that it contains 37.51% C, 4.20% H, and 58.29% O.
What is its empirical formula?
36. Identify the element that has the following electron configuration: 1s22s22p63s23p64s23d5
37. Assuming all volume measurements are made at the same T and P, what V of H2 gas is needed to
react completely with 5.44 L of O2 gas to produce water vapor?
38. Why does the ionization energy decrease as you go down a group?
39. At STP, how many liters does 3.7 moles of N2 gas occupy?
40. What mass of Cl2, in grams, is contained in a 10.0 L tank at 27 oC and 3.5 atm?
41. As you heat a sample of gas, while keeping pressure constant, the sample of gas will ________.
42. In terms of unity and diversity, describe ionic and covalent bonds.
43. A sample of oxygen gas has a volume of 150 mL when its pressure is 0.95 atm. What will the
volume be if the pressure is raised to 0.99 atm (constant T)?
44. Iron is usually produced from iron ore through the following process in a blast furnace:
Fe2O3
+ CO
Fe
+ CO2
If 4.00 kg of iron (III) oxide are available to react, how many moles of CO are needed? How many
grams of CO are needed? How many molecules of CO are needed?
45. Zinc and sulfur react to form zinc sulfide according to the following reaction:
Zn (s) +
S8 (s)
ZnS (s)
If 84 grams of Zn are heated with 84 grams of S8, identify the limiting reactant.
46. If a reaction were to take place, at STP, between 4.00 grams of hydrogen gas and 22.4 liters of
oxygen gas, what volume of liquid water would be produced? What would be the volume of the
excess gas?
47. Please make a graph of temperature versus heat extracted as a gas is cooled well into the solid
phase. Also describe what is taking place in terms of intermolecular forces and molecular motion.
SOLUTIONS:
48. What is the molarity of a solution in which you dissolve 59 grams of sodium chloride in enough
distilled water to create 250.0 ml of solution?
49. Explain how solubility and temperature are related for all solutes and solvents.
50. What is the molality of a solution made by dissolving 2 grams of NaOH in 200 ml of solvent?
51. Please calculate the parts per million and parts per billion for the solution in number 50.
52. How many grams of lead iodide will precipitate out of solution if an excess of sodium iodide were
added to 2.2 ml of a tap-water solution with a lead (II) ion concentration of 2.0 ppb.
53. What will the freezing point be of a solution made by dissolving 5.0 grams of solid calcium
chloride in 50.0 ml of water?
54. Describe in detail what happens when a crystal of salt is dropped into each of the following then
stirred:
a] a saturated salt solution
b] an unsaturated salt solution
c] a supersaturated salt solution
55. Identify the solution that has the lowest freezing point and explain your reasoning:
1m C2H5OH
1m NaOH
1m Ca(OH)2
THERMO:
56. Calculate the ∆H in kJ/mole NH3 for the combustion of ammonia, NH3, in pure oxygen to produce liquid
water and nitrogen dioxide gas. HINT: use Hess's Law
57. Calculate the ∆H produced by burning octane, C8H18, in kcal/mole of fuel burned using the data below:
wt. of empty beaker
193g
initial wt. of beaker & octane
290g
final wt. of beaker & octane
213g
wt of empty 600ml flask
150g
volume of water in flask
220 ml
initial water temp in flask
15oC
final water temp in flask
56oC
tap water temp.
30oC
room temp.
20oC
time of heating
10 min
58. If 35.5g a given food serving of food has 1.5 grams of protein and 6.7 g of carbohydrates and 5.5
grams of water, how many Calories are from fat? Provided that protein supplies 4.0 Calories per
gram, carbohydrates supply 3.7 Calories per gram, and fats supply 9.0 Calories per gram.
59. Where do you write the energy value for an endothermic change, when you include it in a
balanced equation, and why is it written there?
60. Use Hess’s Law to calculate ∆H for the combustion of one mole of ethylene, C2H4, assuming all
products are gases. Show the mechanism of the reaction by listing each step and label each step as
exothermic or endothermic.
61. Use Hess’s Law to calculate ∆H for this reaction: CuO (s) + Cu (s)
Cu2O (s)
62. Calculate, if you will please (using standard heats of formation), the amount of heat energy
generated when 175 grams of propane C3H8 gas is burned in a gas grill.
63. If, on the lab practical portion of the final exam, you were asked to test your answer to the
question above, Describe the procedure you would employ in the laboratory.
64. Predict whether or not the following reactions are spontaneous or nonspontaneous at room
temperature and 1 atmosphere of pressure – and explain for each in 1 sentence or less.
a. 2 KClO3(s) + 3 C(s)
3 CO2(g) + 2 KCl(s)
exothermic
b. 2 NO(g) + CO2(g)
2 NO2(g)
exothermic
c. H2(g)
+
Zn(ClO3)2(s)
2 HClO3(l) + Zn(s)
endothermic
EQUILIBRIUM
65. What is equal about equilibrium?
66. What is K for the equilibrium 2 NO(g) + O2 (g)
2 NO2 (g)
given that, 1 liter at equilibrium there exists 8 moles of NO2, 3 moles of NO and 2 moles of O2.
67. At 1273°C, Kc for 2 CO(g) + O2(g)
2 CO2(g) is 2.24 × 1022. What is Kp for the reaction at the
same temperature? If enough CO2 were injected into a rigid sealed container to create an initial
pressure of 2.5 atm, what would be the partial pressures of each gas when equilibrium is reached?
68. What is the solubility, in moles per liter, of Zn(OH)2 if the Ksp = 4.5 x 10-17?
69. If the solubility of Li2CO3 = 0.15 moles per liter, what is the Ksp at this temperature?
70. If a lead iodide solution has a lead ion concentration of .00049, and an iodide ion concentration of
0.0001, given that PbI2 if the Ksp = 8.5 x 10-9, Is the solution unsaturated saturated or supersaturated?
71. Imagine that you a flask containing 2.0 liters of saturated lead chloride solution with no excess
solid. What would the flask weigh if all of the water were evaporated out? [The empty flask weighs
172.87 grams and the Ksp for lead chloride is 1.6 x 10-5].
72. Describe what happens to the amount of each chemical in the equilibrium state below:
2NH3(g)
+ 92 kJ
N2(g) + 3H2(g)
a. Decrease N2
b. Increase H2
c. Decrease temperature
d. Decrease pressure
ACID/BASE
73. Please calculate the hydrogen ion concentration in mol/L for solutions with the following pH values:
(a) 2.42
(b) 11.21
(c) 15.00.
74. Please calculate the pH of each of the following solutions:
(a) 5.2 X 10-4 M Ba(OH)2, (b) 2.8 X 10-4 M HNO3.
75. A 35.7 ml amount of 1.4 M acid was titrated to the end point with 10.0 ml of base of unknown
molarity, then 12.9 ml of this base was titrated with 6.41 grams of an unknown solid acid. Calculate
the concentration of the base and the molecular wt. of the solid acid.
76. Calculate the pH of 3.5 M acetic acid, Ka = 1.8 x 10-5. What would be the pH if 50.0 ml of this
acid were titrated with 25.0 ml of 3.5 M NaOH? What would the pH be after being titrated with 50.0
ml of NaOH?
22. 4.5 x 108kJ, 23. 127.5 g, and 0.060 g, 25. 3.03 x 10-19 J, 35. C6H8O7, 37. 10.9 L H2,
39. 83 L, 40. 1.0 x 102 g, 43. 140 ml, 44. 75.1 mol, 2100 g, 4.52 x 1025 molecules,
45. Zinc, 46. 35.7 mL H2O, 0.2 L O2, 48. 4.0 M, 50. 0.3 m, 51. 10,000 ppm, 10,000,000 ppb,
52. 4.4 x 10-9 g PbI2, 53. -5.0 oC, 56. -349.46 kJ/mol, 57. 13.381 kCal 58. 196.2 Cal,
60. -1410.91 kJ/mol 61. -11.3 kJ/mol, 62. If you considered water to be a liquid, -8807.700 kJ. If
you considered water to be a gas, -8109.059 kJ, 64. a) yes, DH < 0, DS > 0, therefore always
spontaneous. b) depends, DH < 0, DS < 0, therefore is spontaneous at high temperatures.
c)
20
-7
no, DH > 0, DS < 0, therefore never spontaneous. 66. 4, 67. 1.76 x 10 , CO = 3.1 x 10 atm, O2 =
3.1 x 10-7 atm, 2.5 atm, 68. 2.2 x 10-6 M, 69. 0.014, 70. Q = 4.9 x 10-12 > Ksp, therefore
unsaturated. 71. 181.70 grams. 73. a) 0.0038 M, b) 6.2 x 10-12 M, c) 1.0 x 10-15 M,
74. a) 11.02, b) 3.55, 75. 99.4 g/mol = 99.4 amu/molecule, 1.65 x 10-22 g/molecule, 76. pH = 4.74,
pH = 9.5