AP CHEMISTRY ASSIGNMENT- DUE FIRST DAY BACK : Answer on a separate sheet of paper numbers 1 through 151.
1. Distinguish between quantative and qualitive measurements. Give an example of each.
2. Distinguish between the accuracy and precision of a measurement.
3. Distinguish between physical and chemical changes in matter. Give an example of each.
4. Distinguish between a homogeneous and heterogenous mixture. Give an example of each.
5. Name the five major divisions of Chemistry.
6. What are the SI base units for each of the following: length, mass, time, temperature, amount of substance?
7. What are the steps involved in the scientific method?
8. Name three physical properties and distinguish between extensive and intensive.
9. Name three chemical properties.
10. Rewrite the following in scientific notation: .000007345, 4 304 000, .00230
11. Give the number of significant figures in the following: 1603, 3.28 x 10-3, 1500, .00240
12. Round off the measurement 417.01 g to three significant figures: _____________
13. Round off the measurement .0030955 to three significant figures: _____________
14. What is the sum of 6.210 L and 4 L expressed in the correct number of significant figures?
15. A rock has a mass of 127 g and displaces 32.4 mL of water. What is the density of the rock?
16. A piece of copper has a volume of 28.6 cm3. The density of copper is 8.92 g/ cm3. What is the mass of the copper?
17. The accepted value for the density of sodium chloride is 2.165 g/mL. A student measured the density as 2.075 g/mL. Calculate the percent error of the measurement.
18. What is the volume of an object with a density of 5.20 g/mL and a mass of 600 g?
19. What is the temperature 195 K expressed in degrees Celsius?
20. How many joules are there in 165.2 calories?
21. Make the following conversions:
a. 382.2 mg to g
b. 5.0 x 103 mm to meters
c. .12 L to mL
d. 13.5 km to meters
e. 18.7 x 103 mg to decigrams
22. List five forms of energy.
23. Distinguish between potential energy and kinetic energy and give an example of each.
24. A 55.0 g piece of copper wire is heated and the temperature of the wire changes from 19.0 C to 86.0 C. The amount of heat absorbed in 343 cal. What is the specific heat of copper?
25. Define isotope.
26. What is the charge and location of the following subatomic particles in a neutral atoms?
a. Proton:
b. Neutron:
c. Electron:
27. What is the charge of the nucleus of a neutral atom? What is the charge of a neutral atom?
28. How can you find the number of neutrons an atom has?
29. The atomic number is always equal to the number of ________________ in an atom.
30. What are the two different ways to express an isotope? _____________________ ____________________________________________________________________
31. How do the isotopes protium, deutrium, and tritium differ?
32. Average relative atomic masses are measured in ___________________.
33. There are four naturally occurring isotopes of the element chromium. The relative abundance of each is:
Chromium – 50 4.31%
Chromium – 52 83.76%
Chromium – 53 9.55%
Chromium – 48 2.38%
34. Explain how the atoms of one element differ from those of another element.
35. If the atomic number of an element is 8, then how many neutrons, protons, and electrons are there in a neutral atom?
36. Define quantum.
37. Define orbital.
38. What are the colors of the visible spectrum?
39. Name six types of electromagnetic waves.
40. Draw a diagram of a wave and label the following terms: wavelength, amplitude, origin, crest, and trough.
41. Who discovered the electron?
42. Proton?
43. What are the shapes of the s and p orbitals?
44. What is the maximum number of electrons each of the following orbtals hold?
a. s b. p c. d d. f
45. How many orbitals are in the following?
a. s b. p c. d d. f
46. If three electrons are available to fill three empty 2p atomic orbitals, how will the electrons be distributed in the three orbitals?
47. If the spin of one electron in an orbital is clockwise, what is the spin of the other electron in that orbital? Why?
48. What is the next atomic orbital in the series 1s, 2s, 2p, 3s, 3p, 4s? __________
49. Write the electron configurations for the following elements and their ions:
a. K K+
b. P P-
c. O O-2
d. Ca Ca+2
50. Which color of visible light has the shortest wavelength?
51. The amplitude of a wave is the measure of _____________________________.
52. Define stable electron configuration.
53. Which electromagnetic wave has the highest frequency?
54. What is the frequency of light with a wavelength of 1.0 x 10-6 cm?
[f = c/λ] [c = 3.0 x 1010 cm/s]
55. What is the approximate energy of a photon having a frequency of 4.0 x 107 s-1?
[E = h x f] [h = 6.626 x 10-34 J s ]
56. Which of the following are representative elements: Mg, Li, Fe, Se, K, Pb, Ni, Br, Na, S, O?
57. Which of the following are transition elements: Sr, Ru, As, W, Ag, Al, Ni, Mg, Zn, Sn, Sb?
58. Which of the following are metals: Ta, Nd, Se, F, C, Br, Ba, Na, Cl, Ag, Sr?
59. Which elements constitute the halogens?
60. Which elements constitute the noble gases?
61. Which elements constitute the alkali metals?
62. Which elements constitute the alkaline earth metals?
63. List four different things you can find out about lithium by using the periodic table.
1.
2.
3.
4.
64. What are the typical charges of the following monatomic ions?
F ___ O ___ Al ___ Mg ___ Cl ___ P ___ Be ___ Na ____ Ba ____
65. Name all of the elements that form diatomic elements.
66. What are the formulas for the following polyatomic ions?
a. acetate _____________ d. nitrate ______________
b. phosphate _____________ e. sulfate ______________
c. carbonate _____________ f. sulfite ______________
67. Distinguish between iconic compounds and molecular compounds.
68. Define empirical formula.
69. Write the chemical formulas for each of the following:
a. magnesium chloride b. tin (IV) oxide
c. diphosphorous trioxide
d. potassium nitrate e. calcium hydroxide
f.carbon tetrachloride g. carbonic acid
h.hydrochloric acid i. sulfurous acid
j. ammonium chloride k. dinitrogen pentoxide l. iron (III) acetate
70. State the number of electrons lost or gained in forming each of these ions:
a. Mg+2 b. Br- c. Ag+ d. O-2 e. Al+3
71. Determine the empirical formula of the compound with the percent composition of 29.1% Na, 40.5% S, and 30.2 O.
72. How many kilograms of iron can be recovered from 643 kilograms of the ore Fe2O3?
73. Avogadro’s number of particles is equal to one ____________.
74. The representative particle of an element is a(n) ___________________.
75. The representative particle of a molecular compound is a(n) _______________.
76. The representative particle of an ionic compound is a(n) __________________.
77. Convert 50.4 grams CaBr2 to moles.
78. How many grams are there in 1.26 moles NaI?
79. How many moles are there in 86.2 grams C2H4?
80. How many molecules are there in 0.943 moles H20?
81. How many grams are there in 2.54 liters of O2 gas?
Complete and write balanced equations for each of the following:
82. calcium chloride + nitric acid à
83. aluminum + iron (III) oxide à
84. Mg(OH)2 + HNO3 à
85. PbO2 à
86. Pb(NO3)2 + H2SO4 à
87. Water and diphosphorous pentoxide yields phosphoric acid.
88. Hydrogen and chlorine yields hydrochloric acid.
89. Calcium oxide and water yields calcium hydroxide.
90. What are the products of a combustion reaction?
91. What are the five different types of chemical reactions? Give an example of each.
92. Write a balanced equation for the following: water decomposes into hydrogen and oxygen gas when electricity is added.
93. How many moles of oxygen are produced if you have 2.5 moles of water?
94. How many liters of hydrogen gas are produced if 7.2 moles of water are used?
95. How many grams of oxygen will be produced if 10.0 g of water are used?
96. Distinguish between limiting and excess reagents.
97. The percentage yield of the reaction 2 Al + 2 H3PO4 à 2 Al PO4 + 3H2 is 73.7%. How many grams of Al must be used to react with an excess of H3PO4 to give 30.0g of Al PO4?
98. If a 200.0g sample of Al is reacted with 175.0 L of O2 at STP according to the reaction 4 Al + 3O2 à 2 Al2O3, what is the limiting reagent?
99. Calculate the molarity of a solution made by adding 34.8 g of K2SO4 to 800 g of water.
100. Calculate the concentration of a solution composed of 13.1 g of Ba(NO3) 2 dissolved in 750 g of water.
101. If 100.0 mL of a 12.0 M HCl solution is diluted to 2.00 L, what is the molarity of the final solution?
102. How many milliliters of 16.0 M HNO3 would be required to prepare 750. mL of a .500 M solution?
103. You wish to prepare an 8.00% (m/v) solution containing 4.50 g of NaOH. How many milliliters of water would you use?
104. Calculate the percent (m/v) of a solution that was prepared by dissolving 5.29 g of silver nitrate in 81.0 mL of solution.
105. What is the molarity (moles/1 L) of a solution that contains 212.5 g of sodium in 3.0 L solution?
106. You must prepare 300 mL of .75 M NaBr stock solution. How many milliliters of stock solution should you use?
107. What is the mole fraction(moles:moles) of ethanol in a solution of 5.45 moles of ethanol and 25.5 moles of water?
108. What is the boiling point of a solution that has 5 moles of KBr in 2500 g of water? (Kb = 0.512 oC/m constant)( see boiling point depression)
109. What is the freezing point of a solution that has 5 moles of NaI in 2250 g of water? ((KF = 1.86 oC/m constant)
110. Balance the following equation: CuCl (aq) + H2S (s) à Cu2S (s) + HCl (aq)
How many moles of Cu2S could be produced from 9.5 moles of CuCl reacting with an excess of H2S gas?
How many grams of Cu2S could be produced from 9.90 g of CuCl reacting with an excess of H2S gas?
How many liters of gas would be used to produce 2.50 moles of HCl? (22.4 L = I mole at STP)
How many grams of CuCl would you need to produce 1.75 moles of Cu2S?
111. What mass of silver can be produced from 3.00 moles of copper and 3.85 moles of silver nitrate? What is the limiting reagent? What is the excess reagent? How much reagent is left?
112. From the following empirical formulas and the formula weight for each compound, determine the correct molecular formulas.
a. Vitamin C, C3H4O3, molar mass = 176.0 g/mol)
b. Nicotine, C5H7N, molar mass = 81.0 g/mol)
113. A Freon gas sample with a molecular weight of 121.0 g/mol was found to be 9.92% carbon, 58.68% chlorine, and 31.40% fluorine. What is the molecular formula of the Freon gas?
114. Determine the boiling point and freezing point of a solution of 10.0 g of urea ((NH2) 2CO) in 200 g of water.
115. Write the formulas for the following compounds: calcium hydroxide, barium sulfate, calcium selenide, iron (II) oxide, zinc oxide, sodium peroxide, copper (!!) sulfate, sulfurous acid, nitric acid.
116. Write the dot structure for the following: K+, Sr, N, Br, O-2
117. Write the outer electron configuration for the following: K, Na, Cl, Br, N, O, I, Ca+2, H
118. A diatomic molecule with a triple covalent bond is ______________.
119. A diatomic molecule with a double covalent bond is _______________.
120. How many unshared pairs of electrons does the nitrogen atom in ammonia possess?
121. How many unshared pairs of electrons are there in a molecule of hydrogen iodide?
122. Draw the structural formulas, label their shapes, and name the following compounds:
Cl4, C2Cl2, CH4, C2Cl4, NCl3, PBr3, ClF, OF2, NH3, CCl4, O2, N2
123. How much heat will be absorbed by 320 g of water when its temperature is raised by 35oC? (The specific heat for water is 4.18 J/g x C)
124. Calculate the specific heath for aluminum if 16 500 J ofheat are absorbed in raising the temperature of 1.50 x 10 2 g of aluminum by 125 oC.
125. Calculate the number of moles of oxygenin a 12.5 L tank at a pressure of 25 325 kPa, measured at 22 oC.
126. The volume of a balloon is 340 mL at a pressure of 142 kPa and a temperature of 25 oC. Calculate the new volume if the temperature is raised to 85 oC and the pressure changes to 182 kPa. (See gas laws)
127. A volume of 3.0 L of air is warmed from 50 oC to 100 oC. What is the new volume if the pressure remains constant?
128. What is the temperature of the gas inside a 750 mL balloon filled with 0.030 g H2 gas? The pressure of the balloon is 120 kPa.
129. Nitrogen gas in a steel cylinder is under a pressure of 15 200 kPa at 27 oC. What will the pressure in the tank be if the tank is left in the sun and the internal temperature rises to 55 oC?
130. A sample of gas occupies a volume of 80 mL at a pressure of 0.50 atm and a temperature of 0 oC. What will be it’s volume at a pressure of 1.50 atm and a temperature of 50 oC?
131. Explain why a magnesium atom is smaller than the atoms of both sodium and calcium.
132. Would you expect a Cl- ion to be larger or smaller than an Mg +2 ion? Explain.
133. Which family of elements is characterized by an s2p3 configuration?
134. Name the element that matches the following description:
a. one that has 5 outer electrons on the third row of the periodic table
b. one with a 4s24p5 electron configuration