1 mile = 5280 feet 1 gallon = 4 quarts 1 hour = 60 min 1 kg = lbs. 1 mile = km 1 quart = 2 pints 1 min = 60 s 1 lb = 16 oz

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1 Fall Academic Chemistry Semester Practice Test ANSWER KEY I. MEASUREMENT AND CALCULATION 1. List the metric system prefixes, abbreviations, and meanings for kilo- through milli-. kilo hecto deka BASE deci centi milli k h da BASE d c m Give the appropriate metric unit for measuring each of the following: 2. Mass of a German Shepherd kg 3. Length of your arm cm 4. Volume of a liquid medicine dispenser ml 5. Volume of a bathtub L 6. Distance from Houston to Austin km 7. Convert ml to liters L 8. Convert the mass of your pet cat from 4500 g to kg. 4.5 kg 9. How many dm in 14.5 mm? dm 10. What is the formula for density? d = m/v 11. What is the density of an object with a mass of 5.72 g and a volume of 8.6 ml? 0.67 g/ml 12. What is the volume of a necklace if the density is 21.9 g/ml and the mass is 0.99 g? ml 13. If a metal peg has a density of 18.2 g/ml and it is cut in half and then in half again, what is the density of the smallest remaining piece of the peg? 18.2 g/ml; density does not change with amount! 14. Describe how you would experimentally determine the density of an irregularly shaped object, like a stone. Find the mass of the object on a balance, then drop the object into a graduated cylinder with water in it. Record the volume before and after you add the object to the water and use that change in volume as the volume of the object. D = m/v 15. Define accuracy and precision as related to scientific measurement. Accuracy how close to the actual value a measurement is. Precision how close a set of values are to each other 16. The true volume of a liquid is 45.5 ml. The volume is experimentally measured three times, with the following results: 35.5 ml, 35.6 ml, 35.5 ml. Are these results accurate, precise, both, or neither? Precise, but not accurate 17. How many significant figures are in ? How many significant figures are in 900.? How many significant figures are in ? How many significant figures are in ? How many significant figures are in x 10 5? Calculate and round your answer to the correct # of significant figures: 65.7 g 13 g g 9 g =? 133 g 23. Calculate and round your answer to the correct # of significant figures: ml 12.3 ml 8.41 ml 0.5 ml =? 22.2 ml 24. Calculate and round your answer to the correct # of significant figures: 15 g / 4 ml =? 4 g/ml 25. Calculate and round your answer to the correct # of significant figures: 12.5 cm x 3.5 cm x 8.5 cm 370 cm 26. Convert to scientific notation: (a) x 10-4 (b) x 10 6 (c) x 10 3 (d) x Convert to regular (standard) notation: (a) 4.45 x (b) 6.00 x (c) 1.19 x (d) 7.2 x Use the following conversion factors to perform calculations this section: 1 mile = 5280 feet 1 gallon = 4 quarts 1 hour = 60 min 1 kg = lbs 1 mile = km 1 quart = 2 pints 1 min = 60 s 1 lb = 16 oz 1 inch = 2.54 cm 1 pint = 2 cups 1 cup = 16 T 1 T = 3 t

2 28. How many quarts are in 10. gallons? 40. quarts 29. How many teaspoons (t) are in 8 cups? 384 t 400 t 30. How many kilometers are in inches? km 31. How many ounces are in 50 kilograms? 1764 oz 2000 oz 32. How many miles are in 10. kilometers? mi 6.2 mi 33. How many seconds are in 1.56 days? s s II. MATTER AND CHANGE 34. Contrast the 3 states of matter in terms of shape and volume (definite or indefinite?) Solid definite shape and volume; Liquid definite volume, no definite shape; Gas no definite shape, no definite volume 35. Contrast the 3 states of matter in terms of how closely packed the particles are. Solids particles touching and fixed in place but vibrating; Liquids particles are touching but can slide past each other (not fixed in place); Gases particles are far apart and free to move around rapidly 36. What are the two major types of mixtures? Homogeneous and heterogeneous Define each and tell how you recognize them. Homogeneous no particles visible/looks uniform throughout, particles do not settle out of solution; Heterogeneous particles are visible/non-uniform appearance, particles can settle out/separate Give 2 examples of each. Homogeneous saline solution, kool-aid, etc.; Heterogeneous salad dressing, orange juice with pulp, oil-based paint, etc. 37. What are the two types of pure substances? compounds and elements Define each and give 2 examples of each. Compound substance made up of two or more elements chemically bonded, has a fixed chemical formula and a specific set of properties. Examples: CO 2, NaCl. Element simplest form of matter, composed of 1 type of atom, found on the periodic table. Examples: Fe, S 38. How can a mixture of salt and water be separated? Evaporation (the water becomes a gas and the salt is left behind). Can a compound be separated? yes How? By a chemical reaction What about an element? No, unless it undergoes a nuclear reaction 39. Consider the following equation: CH 4 2 O 2 CO 2 2 H 2O If the mass of the reactants is 80 g, and the mass of the CO 2 is 44 g, what is the mass of the water? 36 g What law does this illustrate? Law of Conservation of Mass 40. What is the difference between a chemical property and a physical property? Chemical property determined by the reactivity of a substance; physical property can be determined without any chemical reaction taking place 41. What are the five signs of a chemical reaction? Precipitate formed, gas formed, light produced, energy (temp) change, permanent color change 42. What is an intensive property? Independent of amount (eg. density, color, etc) An extensive property? Dependent on amount (eg. mass, volume, etc.) 43. How can exothermic and endothermic reactions be identified? Exothermic releases energy (energy is exiting the system); endothermic absorbs energy (energy is entering the system.) Determine whether each of the following is a physical or chemical change: 44. iron rusting chemical 45. wood burning chemical 46. punching holes in a piece of paper physical 47. a leaf changing color in the fall chemical 48. meat rotting chemical 49. water freezing physical Determine whether each of the following is an intensive or extensive property: 50. Mass extensive 51. Density intensive 52. Temperature intensive 53. Color intensive

3 Which type of thermal energy transfer is occurring in each of the following: 54. A flask on a hot plate becomes warm after the hot plate is turned on. conduction 55. Hot steam from the shower warms the air in the bathroom. convection 56. Sunlight travels through space to the earth. radiation III. ATOMIC STRUCTURE/ NUCLEAR CHEMISTRY 57. What is the definition of an atom? The smallest particle of an element that retains the properties of that element. 58. How does Dalton's atomic theory help to explain the Law of Conservation of Mass? Part of Dalton s theory states that atoms cannot be created or destroyed in ordinary chemical reactions, just rearranged. 59. Explain Rutherford s gold foil experiment. What did Rutherford discover about the structure of the atom? In the experiment, alpha particles (helium nuclei) produced from the radioactive decay of polonium streamed toward a sheet of gold foil. To Rutherford s surprise, some of the alpha particles bounced off of the gold foil. This meant that they were hitting a dense, relatively large object, which Rutherford called the nucleus. 60. Where are each of subatomic particles located in the atom? Protons and neutrons: nucleus; electron: electron cloud What is the charge on each particle? Proton 1; neutron 0; electron Why is an atom electrically neutral? Because the number of protons and electrons are the same and they have the opposite charge, so the positive and negative charges cancel out. 62. How do isotopes of an element differ? They have different number of neutrons (which makes their masses different) 63. Define: atomic number the number of protons (what defines an element) and what the periodic table is organized by; mass number the sum of the number of protons and neutrons for an isotope; atomic mass the weighted average mass of the element, including all isotopes and their %abundance. 64. Calculate the number of protons, neutrons and electrons in the following isotopes: (a) Oxygen-18 8 protons, 8 electrons, 10 neutrons (b) Cadmium protons, 48 electrons, 64 neutrons (c) Tellurium protons, 52 electrons, 76 neutrons (d) Neon protons, 10 neutrons, 10 electrons 65. An element is defined by the number of protons it has, also known as the atomic number. 66. Find the average atomic mass for oxygen, which has 3 isotopes: oxygen-16 (mass amu and 99.8%), oxygen-17 (mass amu and 0.04%), and oxygen-18 (mass amu and 0.2%). Check on Periodic Table amu 67. Where are the following located on the periodic table? a. metals to the left of the stairstep line c. nonmetals to the right of the stairstep line ( H) b. metalloids on the stairstep line 68. Classify the following elements as metal (M), nonmetal(nm), or metalloid (MD): carbon-nm hydrogen-nm tin-m germanium-md 69. What are cations? Positively charged ions Anions? Negatively charged ions 70. Give the charge (oxidation number) for the following ions: Na 1 PO 4 3- Cl 1- NH 4 1 NO 3 1- O 2- iron-m potassium-m silicon-md boron-md Which are formed from metals? Cations From nonmetals? Anions SO Write the complete reaction for the following radioactive decay processes: Alpha: Mo He Zr Li He H Md He Es Beta: U e Np Pa e U Ac e Th

4 72. Define half-life. The time during which a given atom of a radioactive isotope has a chance of decaying OR the time it takes for half of a sample of a radioactive isotope to decay. IV. CHEMICAL NAMES AND FORMULAS 73. How can ionic compounds be recognized? Compounds formed between a M NM, M polyatomic ion, polyatomic ion NM, or polyatomic ion polyatomic ion. What type of bond do they have? Ionic bond (attraction between a positive and a negative charge) 74. Which type of ion comes first in an ionic compound formula? The positive ions (cations), which are the metals, also NH How can molecular compounds be recognized? Compounds with two nonmetals (or metalloid) What type of bond do they have? They have covalent bonds (bonds formed by sharing of electrons) 76. The elements of which group/family do not typically form bonds? Noble gases Why not? Because they already have 8 valence electrons Classify each of the following as ionic, molecular, acid, or base: 77. Na 2S ionic 78. H 2CO 3 acid 79. P 4O 10 molecular 80. KOH base 81. (NH 4) 3PO 4 ionic Ionic Compounds: write the name. 82. Na 2O sodium oxide 83. Cu 3 PO 4 copper (I) phosphate 84. Sn(CO 3 ) 2 tin (IV) carbonate Molecular Compounds: write the name. 85. As2S3 diarsenic trisulfide 86. XeF4 xenon tetrafluoride Acids: write the name. 87. HBr hydrobromic acid 88. HC 2H 3O 2 acetic acid Ionic Compounds: write the formula. 89. aluminum chlorate Al(ClO 3) iron(iii) nitrate Fe(NO 3) 3 Molecular Compounds: write the formula. 91. phosphorus pentachloride PCl dichlorine heptoxide Cl 2O 7 Acids: write the formula. 93. hydrofluoric acid HF 94. nitric acid HNO 3

5 V. CHEMICAL QUANTITIES (MOLES, EMP/MOLEC FORMULAS, % COMP) 95. Find the molar mass for K 2SO g/mol 96. Find the molar mass for Ca 3(PO 4) g/mol 97. Find the molar mass for NH 4NO g/mol 98. Convert 65.4 mol Na 3PO 4 to molecules x molecules 99. Convert 1.23 moles KCl to grams g 100. Convert 5.97 g BF 3 to molecules x molecules 101. Convert 9.5 L CO 2 gas at STP to grams g 19 g 102. Calculate the percent composition of phosphorus in AlPO % 103. Determine the mass of aluminum in a 150. g sample of AlPO g 104. Determine the empirical formula for a compound containing 40.0% C, 6.7% H, 53.3% O. CH2O 105. The empirical formula for vitamin C is C 3H 4O 3. Experimental data indicates that the molar mass of vitamin C is g/mol. What is the molecular formula of vitamin C? C6H8O6 VI. ELECTRONS IN ATOMS 106. What is the frequency of radiation with a wavelength of 3.82 x 10-7 m? 7.85 x Hz 107. Gamma rays have frequencies around 1.0 x Hz. What wavelength does this correspond to? 3.00 x m 108. It takes 6.63 x J of energy to eject an electron from a certain atom.what frequency of light is this? 1.00 x10 16 Hz 109. List the number of orbitals contained in each of the s, p, d, and f sublevels. s: 1, p: 3, d: 5, f: How many electrons can occupy any s sublevel? 2 p? 6 d? 10 f? Write electron configurations and orbital (arrow) diagrams for these elements: B, N, Ne, K B: 1s 2 2s 2 2p 1 N: 1s 2 2s 2 2p 3 Ne: 1s 2 2s 2 2p 6 K: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 (ask teacher for orbital diagrams or see class notes for examples) 112. The element with electron configuration 1s 2 2s 2 2p 6 3s 2 3p 4 is? sulfur 113. Define valence electrons. Electrons located in the outermost energy level, they are the electrons that are involved in chemical reactions. Explain how you would find the valence electrons for an element and give an example. The number of valence electrons equals the number of electrons in the highest energy level (it also equals the group number for the s block and the group number -10 for the p block). Examples: Na has 1, Ca has 2, Al has 3, C has 4, P has 5, O has 6, Cl has 7, and Ar has 8 valence electrons Draw electron dot diagrams for calcium, oxygen, argon, and phosphorus. Your diagrams should show: Ca: 2 dots, singles O: 6 dots, 2 paired, 2 singles Ar: 8 dots, 4 paired P: 5 dots, 1 paired, 3 singles 115. Write the electron configuration for both a sodium atom and a sodium ion (charge?) Na: 1s 2 2s 2 2p 6 3s 1 Na : 1s 2 2s 2 2p 6 (lost 1 electron) 116. Write the electron configuration for both an oxygen atom and an oxide ion (charge?) O: 1s 2 2s 2 2p 4 O 2- : 1s 2 2s 2 2p 6 (gained 2 electrons) 117. Draw a diagram using dots and arrows to show the ionic compound that forms when sodium and oxygen react. VII. THE PERIODIC TABLE 118. Who is credited with developing the first periodic table? Dmitri Mendeleev 119. Mendeleev s periodic table was arranged in order of increasing atomic mass Discuss the power of Mendeleev s periodic table; what was he able to predict? He was able to predict the existence of undiscovered elements Moseley made an adjustment to Mendeleev s table by realizing that the elements are arranged in order of increasing atomic number Why is Mendeleev's table called periodic? The properties of the elements repeat in a regular pattern.

6 123. What can be said about the properties of elements in the same group or column in the periodic table? They will have similar properties 124. Where are the following families found on the periodic table? (list their group #): a. alkali metals: group 1 b. alkaline earth metals: group 2 c. halogens: group 17 d. noble gases: group Where are the transition metals on the periodic table? In what block are they found? middle, d block 126. Where are the lanthanides and actinides on the periodic table? In what block are they found? bottom, f block 127. Define atomic radius. A measure of the size of the atom; distance from the center of the nucleus to the edge of the electron cloud How does atomic radius change: a. going down a group? Increases b. across a period left to right? Decreases 129. Define ionization energy. The amount of energy required to remove an electron from an atom How does ionization energy change: a. going down a group? Decreases b. across a period left to right? Increases 131. Define electronegativity. The tendency of an atom to attract electrons to itself when it is chemically bonded with another element How does electronegativity change: a. going down a group? Decreases b. across a period left to right? Increases c. Considering this information, what is the most electronegative element on the periodic table? Do not consider noble gases, because they generally do not participate in bonding! Fluorine 133. On what side of the Periodic Table do elements have more: a. metallic character? Left side b. nonmetallic character? Right side 134. What type of element tends to: a. gain electrons to become stable? Nonmetals b. lose electrons to become stable? Metals Part II: FREE RESPONSE KEY x molecules x molecules L % g 11. 1s 2 2s 2 2p 6 3s [Xe] 6s 2 4f 14 5d 10 6p calcium phosphate 14. MgBr dinitrogen pentoxide 16. BF hydrochloric acid 18. H 2SO HClO 2 20.

Academic Chemistry Fall 2016 Semester Practice Test (Clarke, Eller, Hessel, and White)

Name: Academic Chemistry Fall 2016 Semester Practice Test (Clarke, Eller, Hessel, and White) Period: Test Format: 100 multiple choice questions at 1 point each I. MEASUREMENT AND CALCULATION 1. List the