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Which type of matter is composed of two or more elements that are chemically combined in a fixed proportion? (1) solution (3) homogeneous mixture (2) compound (4) heterogeneous mixture 3. Which element has the greatest density at STP? (1) scandium (3) silicon (2) selenium (4) sodium 4. Matter is classified as a (1) substance, only (2) substance or as a mixture of substances (3) homogenous mixture, only (4) homogenous mixture or as a heterogeneous mixture 5. Which substance can not be decomposed by a chemical change? (1) ammonia (3) propanol (2) copper (4) water 6. A beaker contains both alcohol and water. These liquids can be separated by distillation because the liquids have different (1) boiling points (3) particle sizes (2) densities (4) solubilities 7. Which term is defined as a measure of the average kinetic energy of the particles in a sample of matter? (1) activation energy (3) temperature (2) potential energy (4) entropy 8. Which statement describes the transfer of heat energy that occurs when an ice cube is added to an insulated container with 100 milliliters of water at 25°C? (1) Both the ice cube and the water lose heat energy. (2) Both the ice cube and the water gain heat energy. (3) The ice cube gains heat energy and the water loses heat energy. (4) The ice cube loses heat energy and the water gains heat energy. 9. Which quantity of heat is equal to 200. joules? (1) 20.0 kJ (3) 0.200 kJ (2) 2.00 kJ (4) 0.0200 kJ 10. Which equation shows conservation of atoms? (1) H2 +O2 ’! H2O (3) 2 H2 + O2 ’! 2 H2O (2) H2 +O2 ’!2 H2O (4) 2 H2 + 2 O2 ’! 2 H2O 11. Given the balanced particle-diagram equation:  SHAPE \* MERGEFORMAT  Which statement describes the type of change and the chemical properties of the product and reactants? (1) The equation represents a physical change, with the product and reactants having different chemical properties. (2) The equation represents a physical change, with the product and reactants having identical chemical properties. (3) The equation represents a chemical change, with the product and reactants having different chemical properties. (4) The equation represents a chemical change, with the product and reactants having identical chemical properties. 12. Which substance can be decomposed by chemical means? (1) aluminum (3) silicon (2) octane (4) xenon 13. At STP, which 2.0-gram sample of matter uniformly fills a 340-milliliter closed container? (1) Br2(l) (3) Fe(NO3)2(s) (2) KCl(aq) (4) Xe(g) 14. Which two particle diagrams represent mixtures of diatomic elements?  SHAPE \* MERGEFORMAT  (1) A and B (3) B and C (2) A and C (4) B and D 15. The graph below represents the relationship between temperature and time as heat is added to a sample of H2O.  SHAPE \* MERGEFORMAT  Which statement correctly describes the energy of the particles of the sample during interval BC? (1) Potential energy decreases and average kinetic energy increases. (2) Potential energy increases and average kinetic energy increases. (3) Potential energy increases and average kinetic energy remains the same. (4) Potential energy remains the same and average kinetic energy increases. 16. In which sample of water do the molecules have the highest average kinetic energy? (1) 20. mL at 100.°C (3) 60. mL at 60.°C (2) 40. mL at 80.°C (4) 80. mL at 40.°C 17. What is the total amount of heat absorbed by 100.0 grams of water when the temperature of the water is increased from 30.0°C to 45.0°C? (1) 418 J (3) 12 500 J (2) 6270 J (4) 18 800 J 18. Which process is exothermic? (1) boiling of water (3) melting of copper (2) condensation of ethanol vapor (4) sublimation of iodine 19. Which balanced equation represents a chemical change? (1) H2O(l) + energy ’! H2O(g) (2) 2 H2O(l) + energy ’! 2 H2(g) + O2(g) (3) H2O(l) ’! H2O(s) + energy (4) H2O(g) ’! H2O(l) + energy 20. Which equation represents a physical change? (1) H2O (s) + 6.01 kJ ’! H2O (l) (2) 2 H2 (g) + O2 (g) ’! 2H2O(g) + 483.6 kJ (3) H2 (g) + I2 (g) + 53.0 kJ ’! 2 HI (g) (4) N2 (g) + 2 O2 (g) + 66.4 kJ ’! 2 NO2 (g) 21. Starting as a solid, a sample of a substance is heated at a constant rate. The graph below shows the changes in temperature of this sample.  SHAPE \* MERGEFORMAT  What is the melting point of the sample and the total time required to completely melt the sample after it has reached its melting point? (1) 50°C and 3 min (3) 110°C and 4 min (2) 50°C and 5 min (4) 110°C and 14 min Base your answers to questions 22 and 23 on the following information. The boiling point of a liquid is the temperature at which the vapor pressure of the liquid is equal to the pressure on the surface of the liquid. The heat of vaporization of ethanol is 838 joules per gram. A sample of ethanol has a mass of 65.0 grams and is boiling at 1.00 atmosphere. 22. Based on Table H, what is the temperature of this sample of ethanol? [1] 23. Calculate the minimum amount of heat required to completely vaporize this sample of ethanol. Your response must include both a correct numerical setup and the calculated result. [2] Base your answer to questions 24 through 26 on the information below. A method used by ancient Egyptians to obtain copper metal from copper(I) sulfide ore was heating the ore in the presence of air. Later, copper was mixed with tin to produce a useful alloy called bronze. 24. Calculate the density of a 129.5-gram sample of bronze that has a volume of 14.8 cubic centimeters. Your response must include a correct numerical setup and the calculated result. [2] 25. Convert the melting point of the metal obtained from copper(I) sulfide ore to degrees Celsius. [1] 26. A 133.8-gram sample of bronze was 10.3% tin by mass. Determine the total mass of tin in the sample. [1] II. ATOMIC STRUCTURE 1. Which two notations represent different isotopes of the same element? (1) 64Be and 94Be (3) 147N and 146C (2) 73Li and 73Li (4) 3215P and 3215S 2. Which statement describes how an atom in the ground state becomes excited? (1) The atom absorbs energy, and one or more electrons move to a higher electron shell. (2) The atom absorbs energy, and one or more electrons move to a lower electron shell. (3) The atom releases energy, and one or more electrons move to a higher electron shell. (4) The atom releases energy, and one or more electrons move to a lower electron shell. Each diagram below represents the nucleus of a different atom.  SHAPE \* MERGEFORMAT  4. Which diagrams represent nuclei of the same element? (1) D and E, only (3) D, E, and Q (2) Q and R, only (4) Q, R, and E 5. What information is necessary to determine the atomic mass of the element chlorine? (1) the atomic mass of each artificially produced isotope of chlorine, only (2) the relative abundance of each naturally occurring isotope of chlorine, only (3) the atomic mass and the relative abundance of each naturally occurring isotope of chlorine (4) the atomic mass and the relative abundance of each naturally occurring and artificially produced isotope of chlorine 6. In an atom of argon-40, the number of protons (1) equals the number of electrons (3) is less than the number of electrons (2) equals the number of neutrons (4) is greater than the number of electrons 7. An electron in a sodium atom moves from the third shell to the fourth shell. This change is a result of the atom (1) absorbing energy (3) releasing energy (2) gaining an electron (4) losing an electron 8. Which statement describes oxygen gas, O2(g), and ozone gas, O3(g)? (1) They have different molecular structures, only. (2) They have different properties, only. (3) They have different molecular structures and different properties. (4) They have the same molecular structure and the same properties. 9. What is the overall charge on an oxygen atom that contains eight protons, seven neutrons, and eight electrons? (1) -2 (3) +8 (2) 0 (4) +15 10. Given the bright-line spectra of three elements and the spectrum of a mixture formed from at least two of these elements:  SHAPE \* MERGEFORMAT  Which elements are present in this mixture? (1) E and D, only (3) D and G, only (2) E and G, only (4) D, E, and G 11. Which electron configuration represents an atom in an excited state? (1) 2–7 (3) 2–8–1 (2) 2–6–2 (4) 2–8–8–2 12. According to the electron-cloud model of the atom, an orbital is a (1) circular path traveled by an electron around the nucleus (2) spiral path traveled by an electron toward the nucleus (3) region of the most probable proton location (4) region of the most probable electron location 13. The table below gives information about the nucleus of each of four atoms.  SHAPE \* MERGEFORMAT  How many different elements are represented by the nuclei in the table? (1) 1 (2) 2 (3) 3 (4) 4 14. What is the total number of electrons in an atom of potassium? (1) 18 (2) 19 (3) 20 (4) 39 15. The diagram below represents the nucleus of an atom.  SHAPE \* MERGEFORMAT  What are the atomic number and mass number of this atom? (1) The atomic number is 9 and the mass number is 19. (2) The atomic number is 9 and the mass number is 20. (3) The atomic number is 11 and the mass number is 19. (4) The atomic number is 11 and the mass number is 20. 16. The wave-mechanical model of the atom is required to explain the (1) mass number and atomic number of an atom (2) organization of atoms in a crystal (3) radioactive nature of some atoms (4) spectra of elements with multi-electron atoms 17. Explain, in terms of both electrons and energy, how ions emit light. Base your answers to questions 18 through 20 on the chart below.  SHAPE \* MERGEFORMAT  18. State, in terms of the number of subatomic particles, one similarity and one difference between the atoms of these isotopes of sulfur. 19. In the space below, draw a Lewis electron-dot diagram for an atom of sulfur-33. 20. In the space below, show a correct numerical setup for calculating the atomic mass of sulfur. Base your answer to questions 20 and 21 on the information below Scientists are investigating the production of energy using hydrogen-2 nuclei (deuterons) and hydrogen-3 nuclei (tritons). The balanced equation below represents one nuclear reaction between two deuterons. 21H + 21H --> 32He + 10n + 5.23 x 10-13 J 20. State, in terms of subatomic particles, how a deuteron differs from a triton 21. Identify the type of nuclear reaction represented by the equation __________________________ 22. Describe the electrons in an atom of carbon in the ground state. Your response must include:  SHAPE \* MERGEFORMAT  Base questions 23 and 24 on the following information. The nucleus of one boron atom has five protons and four neutrons. 23. Determine the total number of electrons in the boron atom. 24. Determine the total charge of the boron nucleus. Base your answers to questions 25 through 27 on the information below. Carbon has three naturally occurring isotopes, C-12, C-13, and C-14. Diamond and graphite are familiar forms of solid carbon. Diamond is one of the hardest substances known, while graphite is a very soft substance. Diamond has a rigid network of bonded atoms. Graphite has atoms bonded in thin layers that are held together by weak forces. Recent experiments have produced new forms of solid carbon called fullerenes. One fullerene, C60, is a spherical, cage-like molecule of carbon. 25. Determine both the total number of protons and the total number of neutrons in an atom of the naturally occurring carbon isotope with the largest mass number. 26. Identify the type of bonding in a fullerene molecule. 27. State, in terms of the arrangement of atoms, the difference in hardness between diamond and graphite. Base your answers to questions 28 and 29 on the information below. In 1897, J. J. Thomson demonstrated in an experiment that cathode rays were deflected by an electric field. This suggested that cathode rays were composed of negatively charged particles found in all atoms. Thomson concluded that the atom was a positively charged sphere of almost uniform density in which negatively charged particles were embedded. The total negative charge in the atom was balanced by the positive charge, making the atom electrically neutral. In the early 1900s, Ernest Rutherford bombarded a very thin sheet of gold foil with alpha particles. After interpreting the results of the gold foil experiment, Rutherford proposed a more sophisticated model of the atom. 28. State one conclusion from Rutherford’s experiment that contradicts one conclusion made by Thomson. [1] 29. State one aspect of the modern model of the atom that agrees with a conclusion made by Thomson. [1] III. NUCLEAR CHEMISTRY 1. Which type of reaction occurs when a high- energy particle collides with the nucleus of an atom, converting that atom to an atom of a different element? (1) addition (3) neutralization (2) substitution (4) transmutation 2. Which particle is emitted when an atom of 85Kr spontaneously decays? (1) an alpha particle (3) a beta particle (2) a neutron (4) a proton 3. What is a problem commonly associated with nuclear power facilities? (1) A small quantity of energy is produced. (2) Reaction products contribute to acid rain. (3) It is impossible to control nuclear fission. (4) It is difficult to dispose of wastes. 4. An original sample of K-40 has a mass of 25.00 grams. After 3.9 × 109 years, 3.125 grams of the original sample remains unchanged. What is the half-life of K-40? (1) 1.3 × 109 y (2) 2.6 × 109 y (3) 3.9 × 109 y (4) 1.2 × 109 y 5. Which nuclide has a half-life that is less than one minute? (1) cesium-137 (3) francium-220 (2) phosphorus-32 (4) strontium-90 6. Which nuclear emission has the greatest mass? (1) alpha particle (2) beta particle (2) gamma ray (4) positron 7. Which two radioisotopes have the same decay mode? (1) 37Ca and 53Fe (3) 220Fr and 60Co (2) 37K and 42K (4) 99Tc and 19Ne 8. Which list of nuclear emissions is arranged in order from the least penetrating power to the greatest penetrating power? (1) alpha particle, beta particle, gamma ray (2) alpha particle, gamma ray, beta particle (3) gamma ray, beta particle, alpha particle (4) beta particle, alpha particle, gamma ray 9. One benefit of nuclear fission reactions is (1) nuclear reactor meltdowns (3) biological exposure (2) storage of waste materials (4) production of energy 10. Which fraction of an original 20.00-gram sample of nitrogen-16 remains unchanged after 36.0 seconds? (1) 1/5 (2) 1/8 (3) 1/16 (4) 1/32 11. Which radioactive isotope is used in treating cancer? (1) carbon-14 (3) cobalt-60 (2) lead-206 (4) uranium-238 12. Which nuclide is used to investigate human thyroid gland disorders? (1) carbon-14 (2) cobalt-60 (3) potassium-37 (4) iodine-131 13. A beta particle may be spontaneously emitted from (1) a ground-state electron (3) an excited electron (2) a stable nucleus (4) an unstable nucleus The table below indicates the stability of six nuclides.  SHAPE \* MERGEFORMAT  All atoms of the unstable nuclides listed in this table have (1) an odd number of neutrons (3) more neutrons than protons (2) an odd number of protons (4) more protons than neutrons 15. Which of the following elements have no stable isotopes? (1) bismuth (3) potassium (2) boron (4) plutonium Base your answers to questions 16 through 18 on the information below. Cobalt-60 is commonly used as a source of radiation for the prevention of food spoilage. Bombarding cobalt-59 nuclei with neutrons produces the nuclide cobalt-60. A food irradiation facility replaces the cobalt-60, a source of gamma rays, when the radioactivity level falls to 1/8 of its initial level. The nuclide cesium-137 is also a source of radiation for the prevention of food spoilage. 16. Identify one emission spontaneously released by a cobalt-60 nucleus. [1] 17. Determine the total number of years that elapse before an original cobalt-60 source in an irradiation facility must be replaced. [1] 18. Complete the nuclear equation below for the decay of cesium-137. Your response must include the symbol, atomic number, and mass number of the missing particle. 137Cs --> _______ + _______ Base your answers to questions 19 and 20 on the information below. A substance known as heavy water can be obtained from ordinary water and could be a significant source of energy in the future. Heavy water contains deuterium, H-2. Instead of the two hydrogen atoms in a typical water molecule, a heavy water molecule has two deuterium atoms. In 3.78 kilograms of ordinary water, the percent composition by mass of heavy water is approximately 0.0156%. Deuterium atoms completely ionize at approximately 108 K. The result is an ionized gas consisting of electrons and deuterons (the nuclei of deuterium). A triton is the nucleus of a tritium atom, H-3. These particles react according to the equations below. In the second equation, X represents an unidentified product.  SHAPE \* MERGEFORMAT  19. Calculate the mass of heavy water in a 3.78-kilogram sample of ordinary water. Your response must include both a correct numerical setup and the calculated result. [2] 20. Identify particle X in the second nuclear equation. Your response must include the symbol, atomic number, and mass number of the particle. [1] IV. PERIODIC TABLE / BONDING 1. An element that is malleable and a good conductor of heat and electricity could have an atomic number of (1) 16 (3) 18 (2) 29 (4) 35 2. An atom in the ground state has a stable valence electron configuration. This atom could be an atom of (1) Al (3) Cl (2) Na (4) Ne 3. An atom of an element has a total of 12 electrons. An ion of the same element has a total of 10 electrons. Which statement describes the charge and radius of the ion? (1) The ion is positively charged and its radius is smaller than the radius of the atom. (2) The ion is positively charged and its radius is larger than the radius of the atom. (3) The ion is negatively charged and its radius is smaller than the radius of the atom. (4) The ion is negatively charged and its radius is larger than the radius of the atom. Which formula represents a non-polar molecule? (1) CH4 (3) H2O (2) HCl (4) NH3 5. The compound XCl is classified as ionic if X represents the element (1) H (3) I (2) Rb (4) Br 6. The chemical bonding in sodium phosphate, Na3PO4, is classified as (1) ionic, only (3) both covalent and ionic (2) metallic, only (4) both covalent and metallic 7. Which element is composed of molecules that each contain a multiple covalent bond? (1) chlorine (3) hydrogen (2) fluorine (4) nitrogen 8. Magnesium and calcium have similar chemical properties because a magnesium atom and a calcium atom have the same (1) atomic number (3) mass number (2) total number of electron shells (4) total number of valence electrons 9. As sulfur atom is formed, electrons are a bond between a hydrogen atom and a (1) shared to form an ionic bond (3) transferred to form an ionic bond (2) shared to form a covalent bond (4) transferred to form a covalent bond 10. Which type of substance can conduct electricity in the liquid phase but not in the solid phase? (1) ionic compound (3) metallic element (2) molecular compound (4) nonmetallic element 11. What can be concluded if an ion of an element is smaller than an atom of the same element? (1) The ion is negatively charged because it has fewer electrons than the atom. (2) The ion is negatively charged because it has more electrons than the atom. (3) The ion is positively charged because it has fewer electrons than the atom. (4) The ion is positively charged because it has more electrons than the atom. 12. A solid substance is an excellent conductor of electricity. The chemical bonds in this substance are most likely (1) ionic, because the valence electrons are shared between atoms (2) ionic, because the valence electrons are mobile (3) metallic, because the valence electrons are stationary (4) metallic, because the valence electrons are mobile 13. When sodium and fluorine combine to produce the compound NaF, the ions formed have the same electron configuration as atoms of (1) argon, only (3) both argon and neon (2) neon, only (4) neither argon nor neon 14. What is the net charge on an ion that has 9 protons, 11 neutrons, and 10 electrons? (1) 1+ (3) 1– (2) 2+ (4) 2– 15. How do the atomic radius and metallic properties of sodium compare to the atomic radius and metallic properties of phosphorus? (1) Sodium has a larger atomic radius and is more metallic. (2) Sodium has a larger atomic radius and is less metallic. (3) Sodium has a smaller atomic radius and is more metallic. (4) Sodium has a smaller atomic radius and is less metallic. 16. Which group on the Periodic Table of the Elements contains elements that react with oxygen to form compounds with the general formula X2O? (1) Group 1 (3) Group 14 (2) Group 2 (4) Group 18 17. At standard pressure, a certain compound has a low boiling point and is insoluble in water. At STP, this compound most likely exists as (1) ionic crystals (3) non-polar molecules (2) metallic crystals (4) polar molecules Which statement describes the process represented by this equation? Cl + Cl --> Cl2 (1) A bond is formed as energy is absorbed. (3) A bond is broken as energy is absorbed. (2) A bond is formed and energy is released. (4) A bond is broken and energy is released. 19. An oxygen molecule contains a double bond because the two atoms of oxygen share a total of (1) 1 electron (3) 3 electrons (2) 2 electrons (4) 4 electrons 20. Which two characteristics are associated with metals? (1) low first ionization energy and low electronegativity (2) low first ionization energy and high electronegativity (3) high first ionization energy and low electronegativity (4) high first ionization energy and high electronegativity 21. At STP, which element is brittle and not a conductor of electricity? (1) S (3) Na (2) K (4) Ar 22. What is the total number of electrons in a Mg2+ ion? (1) 10 (3) 14 (2) 12 (4) 24 23. At STP, fluorine is a gas and bromine is a liquid because, compared to fluorine, bromine has (1) stronger covalent bonds (3) weaker covalent bonds (2) stronger intermolecular forces (4) weaker intermolecular forces 24. Why is a molecule of CO2 non-polar even though the bonds between the carbon atom and the oxygen atoms are polar? (1) The shape of the CO2 molecule is symmetrical. (2) The shape of the CO2 molecule is asymmetrical. (3) The CO2 molecule has a deficiency of electrons. (4) The CO2 molecule has an excess of electrons. 25. Which formula represents a molecular compound? (1) HI (3) KCl (2) KI (4) LiCl 26. Samples of four Group 15 elements, antimony, arsenic, bismuth, and phosphorus, are in the gaseous phase. An atom in the ground state of which element requires the least amount of energy to remove its most loosely held electron? (1) As (3) P (2) Bi (4) Sb 27. In terms of electron configuration, explain why potassium is larger than lithium. 28. Draw electron dot diagrams a. CaCl2 b. NH3 29. Explain, in terms of electronegativity, why a P–Cl bond in a molecule of PCl5 is more polar than a P–S bond in a molecule of P2S5. [1] Base your answers to questions 30 through 32 on the information below. Elements with atomic numbers 112 and 114 have been produced and their IUPAC names are pending approval. However, an element that would be put between these two elements on the Periodic Table has not yet been produced. If produced, this element will be identified by the symbol Uut until an IUPAC name is approved. 30. In the space in your answer booklet, draw a Lewis electron-dot diagram for an atom of Uut. [1] 31. Determine the charge of an Uut nucleus. Your response must include both the numerical value and the sign of the charge. [1] 32. Identify one element that would be chemically similar to Uut. [1] MOLES, FORMULAS & REACTIONS 1. Which equation represents a decomposition reaction? (1) CaCO3 (s) ’! CaO (s) + CO2 (g) (2) Cu (s)+ 2 AgNO3 (aq)’!2 Ag (s) + Cu(NO3)2 (aq) (3) 2 H2 (g) + O2 (g) ’! 2 H2O (l) (4) KOH (aq) + HCl (aq) ’! KCl (aq) + H2O (l) 2. A compound has the empirical formula CH2O and a gram-formula mass of 60. grams per mole. What is the molecular formula of this compound? (1) CH2O (3) C3H8O (2) C2H4O2 (4) C4H8O4 3. Which formula represents strontium phosphate? (1) SrPO4 (3) Sr2(PO4)3 (2) Sr3PO8 (4) Sr3(PO4)2 4. Which Lewis electron-dot diagram represents calcium oxide?  SHAPE \* MERGEFORMAT  What is the empirical formula for a compound with the molecular formula C6H12Cl2O2? (1) CHClO (3) CH2ClO (2) C3H6ClO (4) C6H12Cl2O2 6. Given the balanced equation representing a reaction: 4 Al (s) + 3 O2 (g) ’! 2 Al2O3 (s) Which type of chemical reaction is represented by this equation? (1) double replacement (3) substitution (2) single replacement (4) synthesis 7. Which two samples of gas at STP contain the same total number of molecules? (1) 1 L of CO (g) and 0.5 L of N2 (g) (2) 2 L of CO (g) and 0.5 L of NH3 (g) (3) 1 L of H2 (g) and 2 L of Cl2 (g) (4) 2 L of H2 (g) and 2 L of Cl2 (g) 8. What is the chemical formula for iron (III) oxide? (1) FeO (3) Fe2O3 (2) Fe3O (4) Fe3O2 9. What is the percent composition by mass of hydrogen in NH4HCO3 (gram-formula mass = 79 grams/mole)? (1) 5.1% (3) 10.% (2) 6.3% (4) 50.% 10. What is the gram-formula mass of Ca3(PO4)2? (1) 248 g/mol (3) 263 g/mol (2) 279 g/mol (4) 310. g/mol Base your answers to questions 75 through 77 on the information below. A fluorescent light tube contains a noble gas and a drop of mercury. When the fluorescent light operates, the Hg is a vapor and there are free-flowing Hg ions and electrons in the tube. The electrons collide with Hg atoms that then emit ultraviolet (UV) radiation. The inside of the tube is coated with a mixture of several compounds that absorbs UV radiation. Ions in the coating emit a blend of red, green, and blue light that together appears as white light. The compound that produces red light is Y2O3. The compound that produces green light is CeMgAl11O19. The compound that produces blue light is BaMgAl10O17. 11. Write the chemical name of the compound that produces red light. [1] Calculate the percent composition by mass of aluminum in the compound that produces green light. Your response must include both a correct numerical setup and the calculated result. [2] Base your answers to 13 through 15 on the following balanced equation:  SHAPE \* MERGEFORMAT  13. Determine the total number of moles of oxygen that react completely with 8.0 moles of C2H6. 14. Determine the mass of 5.20 moles of C6H12 (gram-formula mass = 84.2 grams/mole). [1] 15. Identify the reaction _______________________________ Base your answers to questions 16 through 20 on the information below. Arsenic is often obtained by heating the ore arsenopyrite, FeAsS. The decomposition of FeAsS is represented by the balanced equation below. FeAsS (s) -----> FeS (s) + As (g) In the solid phase, arsenic occurs in two forms. One form, yellow arsenic, has a density of 1.97 g/cm3 at STP. The other form, gray arsenic, has a density of 5.78 g/cm3 at STP. When arsenic is heated rapidly in air, arsenic (III) oxide is formed. Although arsenic is toxic, it is needed by the human body in very small amounts. The body of a healthy human adult contains approximately 5 milligrams of arsenic. 16. Convert the mass of arsenic found in the body of a healthy human adult to grams. [1] 17. When heated, a 125.0-kilogram sample of arsenopyrite yields 67.5 kilograms of FeS. Determine the total mass of arsenic produced in this reaction. [1] 18. Write the formula for the compound produced when arsenic is heated rapidly in air. [1] 19. Explain, in terms of the arrangement of atoms, why the two forms of arsenic have different densities at STP. [1] 20. Calculate the percent composition by mass of arsenic in arsenopyrite. Your response must include both a correct numerical setup and the calculated result. [2] 21. A substance known as heavy water can be obtained from ordinary water and could be a significant source of energy in the future. Heavy water contains deuterium, H-2. Instead of the two hydrogen atoms in a typical water molecule, a heavy water molecule has two deuterium atoms. In ordinary water, the percent composition by mass of heavy water is approximately 0.0156%. Calculate the mass of heavy water in a 3.78-kilogram sample of ordinary water. Base your answers to questions 22 through 24 on the information below. A portable propane-fueled lantern contains a mesh silk bag coated with metal hydroxides. The primary metal hydroxide is yttrium hydroxide. When the silk bag is installed, it is ignited and burned away, leaving the metal hydroxide coating. The coating forms metal oxides that glow brightly when heated to a high temperature. During a test, a propane lantern is operated for three hours and consumes 5.0 moles of propane from the lantern s tank. The balanced equation below represents the combustion of propane. C3H8 + 5 O2 ’! 3 CO2 + 4 H2O + energy 22 At standard pressure, the boiling point of propane is 231 K. In the space below, draw a particle diagram to represent the phase of the propane as it leaves the tank at 294 K. Your response must include at least six molecules. 23. Calculate the total mass of propane consumed during the lantern test. Your response must include both a correct numerical setup and the calculated result. [2] 24. Determine the total number of moles of CO2 produced during the lantern test. 25. Write the empirical formula for the compound C8H18 Base your answers to questions 26 through 28 on the information below. The compound 1,2-ethanediol can be mixed with water. This mixture is added to automobile radiators as an engine coolant. The cooling system of a small van contains 6690 grams of 1,2-ethanediol. Some properties of water and 1,2-ethanediol are given in the table below.  SHAPE \* MERGEFORMAT  26. Identify the class of organic compounds to which 1,2-ethanediol belongs. [1] 27. State, in terms of molecular polarity, why 1,2-ethanediol is soluble in water. [1] 28. In the space below, calculate the total number of moles of 1, 2-ethanediol in the small van’s cooling system. Your response must include both a correct numerical setup and the calculated result. During the construction of the first railroad systems in the United States, sections of rail were welded together by means of the thermite reaction, in which molten iron is produced from a reaction between iron (III) oxide and aluminum metal. 29. Balance the equation for the thermite reaction, using the smallest whole-number coefficients. _____ Al + _____ Fe2O3 --> ____ Fe + ____ Al2O3 30. Classify the thermite reaction: ________________________________ GAS LAWS 1. A 1.0-mole sample of krypton gas has a mass of (1) 19 g (3) 39 g (2) 36 g (4) 84 g 2. According to the kinetic molecular theory, the molecules of an ideal gas (1) have a strong attraction for each other (2) have significant volume (3) move in random, constant, straight-line motion (4) are closely packed in a regular repeating pattern 3. At 65°C, which compound has a vapor pressure of 58 kilopascals? (1) ethanoic acid (3) ethanol (2) propanone (4) water 4. At STP, which 2.0-gram sample of matter uniformly fills a 340-milliliter closed container? (1) Br2 (l) (3) Fe(NO3)2 (s) (2) KCl (aq) (4) Xe (g) Under which conditions of temperature and pressure would a real gas behave most ideally? (1) 200. K and 50.0 kPa (3) 600. K and 50.0 kPa (2) 200. K and 200.0 kPa (4) 600. K and 200.0 kPa Base your answers to questions 6 through 8 on the information below. A sample of helium gas is in a closed system with a movable piston. The volume of the gas sample is changed when both the temperature and the pressure of the sample are increased. The table below shows the initial temperature, pressure, and volume of the gas sample, as well as the final temperature and pressure of the sample.  SHAPE \* MERGEFORMAT  6. In the space below,show a correct numerical setup for calculating the final volume of the helium gas sample. [1] 7. Convert the final temperature of the helium gas sample to degrees Celsius. [1] 8. Compare the total number of gas particles in the sample under the initial conditions to the total number of gas particles in the sample under the final conditions. [1] Base your answers to questions 9 and 10 on the following information. The balanced equation below represents the reaction between magnesium metal and hydrochloric acid to produce aqueous magnesium chloride and hydrogen gas. Mg(s) + 2 HCl (aq) ’! MgCl2 (aq) + H2 (g) A piece of Mg(s) has a volume of 0.0640 cubic centimeters. This piece of Mg(s) reacts completely with HCl(aq) to produce H2 (g). The H2(g) produced has a volume of 112 milliliters and a pressure of 1.00 atmosphere at 298 K. 9. The volume of the piece of Mg(s) is expressed to what number of significant figures? [1] 10. In the space below, show a correct numerical setup for calculating the volume of the H2(g) produced if the conditions are changed to STP. [1] Base your answers to questions 11 through 13 on the information below. Carbon forms molecular compounds with some elements from Group 16. Two of these compounds are carbon dioxide, CO2, and carbon disulfide, CS2. Carbon dioxide is a colorless, odorless gas at room temperature. At standard temperature and pressure, CO2 (s) changes directly to CO2(g). Carbon disulfide is formed by a direct reaction of carbon and sulfur. At room temperature, CS2 is a colorless liquid with an offensive odor. Carbon disulfide vapors are flammable. 11. Identify one physical property and one chemical property of CS2. [1] 12. State what happens to the potential energy of CO2 molecules during this phase change of CO2. 13. Compare the intermolecular forces in CO2 and CS2 at room temperature. [1] 14. The element oxygen can exist as diatomic molecules, O2, and as ozone, O3. At standard pressure the boiling point of ozone is 161 K. Explain, in terms of intermolecular forces, the difference in the boiling points of O2 and O3 at standard pressure. Your response must include information about both O2 and O3. [1] SOLUTIONS 1. Why did the man jump into the homogeneous mixture? (1) To keep his eye on the ions. (2) To save the solvent (3) To salute the solute (4) To be part of the solution. 2. Compared to the freezing point and boiling point of water at 1 atmosphere, a solution of a salt and water at 1 atmosphere has a (1) lower freezing point and a lower boiling point (2) lower freezing point and a higher boiling point (3) higher freezing point and a lower boiling point (4) higher freezing point and a higher boiling point 3. What is the total mass of solute in 1000. grams of a solution having a concentration of 5 parts per million? (1) 0.005 g (3) 0.5 g (2) 0.05 g (4) 5 g 4. Which compound is least soluble in water at 60.°C? (1) KClO3 (3) NaCl (2) KNO3 (4) NH4Cl 5. Which sample of HCl(aq) contains the greatest number of moles of solute particles? (1) 1.0 L of 2.0 M HCl (aq) (3) 3.0 L of 0.50 M HCl (aq) (2) 2.0 L of 2.0 M HCl (aq) (4) 4.0 L of 0.50 M HCl (aq) 6. Which phrase describes the molarity of a solution? (1) liters of solute per mole of solution (2) liters of solution per mole of solution (3) moles of solute per liter of solution (4) moles of solution per liter of solution 7. When 5 grams of KCl are dissolved in 50. grams of water at 25°C, the resulting mixture can be described as (1) heterogeneous and unsaturated (2) heterogeneous and supersaturated (3) homogeneous and unsaturated (4) homogeneous and supersaturated 8. Which aqueous solution of KI freezes at the lowest temperature? (1) 1 mol of KI in 500. g of water (2) 2 mol of KI in 500. g of water (3) 1 mol of KI in 1000. g of water (4) 2 mol of KI in 1000. g of water 9. A thermometer is in a beaker of water. Which statement best explains why the thermometer reading initially increases when LiBr (s) is dissolved in the water? (1) The entropy of the LiBr (aq) is greater than the entropy of the water. (2) The entropy of the LiBr (aq) is less than the entropy of the water. (3) The dissolving of the LiBr (s) in water is an endothermic process. (4) The dissolving of the LiBr (s) in water is an exothermic process. 10. Which barium salt is insoluble in water? (1) BaCO3 (2) Ba(ClO4)2 (3) BaCl2 (4) Ba(NO3)2 11. Which unit can be used to express solution concentration? (1) J/mol (2) L/mol (3) mol/L (4) mol/s 12. Under which conditions of temperature and pressure is a gas most soluble in water? (1) high temperature and low pressure (2) high temperature and high pressure (3) low temperature and low pressure (4) low temperature and high pressure 13. As water is added to a 0.10 M NaCl aqueous solution, the conductivity of the resulting solution (1) decreases because the concentration of ions decreases (2) decreases, but the concentration of ions remains the same (3) increases because the concentration of ions decreases (4) increases, but the concentration of ions remains the same 14. What is the concentration of O2 (g), in parts per million, in a solution that contains 0.008 gram of O2 (g) dissolved in 1000. grams of H2O(l)? (1) 0.8 ppm (2) 8 ppm (3) 80 ppm (4) 800 ppm Base your answers to questions 15 through 17 on the information below A soft-drink bottling plant makes a colorless, slightly acidic carbonated beverage called soda water. During production of the beverage, CO2 (g) is dissolved in water at a pressure greater than 1 atmosphere. The bottle containing the solution is capped to maintain that pressure above the solution. As soon as the bottle is opened, fizzing occurs due to CO2 (g) being released from the solution. 15. Explain why CO2(g) is released when a bottle of soda water is opened. [1] 16. Write the chemical name of the acid in soda water. [1] 17. State the relationship between the solubility of CO2 (g) in water and the temperature of the aqueous solution. [1] 18. When a person perspires (sweats), the body loses many sodium ions and potassium ions. The evaporation of sweat cools the skin. After a strenuous workout, people often quench their thirst with sports drinks that contain NaCl and KCl. A single 250.-gram serving of one sports drink contains 0.055 gram of sodium ions. In the space in your answer booklet, show a correct numerical setup for calculating the concentration of sodium ions in this sports drink, expressed as percent by mass. [1] Base your answers to questions 19 through 21 on the information below. An unsaturated solution is made by completely dissolving 20.0 grams of NaNO3 in 100.0 grams of water at 20.0°C. 19. In the space in your answer booklet, show a correct numerical setup for calculating the number of moles of NaNO3 (gram-formula mass = 85.0 grams per mole) used to make this unsaturated solution. [1] 20. Determine the minimum mass of NaNO3 that must be added to this unsaturated solution to make a saturated solution at 20.0°C. [1] 21. Identify one process that can be used to recover the NaNO3 from the unsaturated solution. [1] Base your answer to questions 22 and 23 on the following information. A solution is made by completely dissolving 90. grams of KNO3 (s) in 100. grams of water in a beaker. The temperature of this solution is 65°C. 22. Describe the effect on the solubility of KNO3 (s) in this solution when the pressure on the solution increases. [1] 23. Determine the total mass of KNO3 (s) that settles to the bottom of the beaker when the original solution is cooled to 15°C. [1] VIII. ACIDS, BASES & SALTS 1. Which substance is an Arrhenius acid? (1) Ba(OH)2 (2) H3PO4 (3) CH3COOCH3 (4) NaCl 2. Which compound releases hydroxide ions in an aqueous solution? (1) CH3COOH (2) CH3OH (3) HCl (4) KOH 3. What are the products of a reaction between KOH(aq) and HCl(aq)? (1) H2 and KClO (3) H2O and KCl (2) KH and HClO (4) KOH and HCl 4. Which volume of 0.10 M NaOH (aq) exactly neutralizes 15.0 milliliters of 0.20 M HNO3 (aq)? (1) 1.5 mL (3) 7.5 mL (2) 3.0 mL (4) 30. mL 5. Which indicator, when added to a solution, changes color from yellow to blue as the pH of the solution is changed from 5.5 to 8.0? (1) bromcresol green (3) bromthymol blue (2) litmus (4) methyl orange 6. The pH of an aqueous solution changes from 4 to 3 when the hydrogen ion concentration in the solution is (1) decreased by a factor of 3/4 (2) decreased by a factor of 10 (3) increased by a factor of 4/3 (4) increased by a factor of 10 7. According to the Arrhenius theory, an acid is a substance that (1) changes litmus from red to blue (2) changes phenolphthalein from colorless to pink (3) produces hydronium ions as the only positive ions in an aqueous solution (4) produces hydroxide ions as the only negative ions in an aqueous solution 8. Which compound dissolves in water to form an aqueous solution that can conduct an electric current? (1) CCl4 (2) C2H5OH (3) CH3COOH (4) CH4 9. Given the equation representing a reaction at equilibrium:  SHAPE \* MERGEFORMAT  The H+ acceptor for the forward reaction is (1) H2O (l) (2) NH4+ (aq) (3) NH3 (g) (4) OH–(aq) 10. Which statement describes an alternate theory of acids and bases? (1) Acids and bases are both H+ acceptors. (2) Acids and bases are both H+ donors. (3) Acids are H+ acceptors, and bases are H+ donors. (4) Acids are H+ donors, and bases are H+ acceptors. Which substance is an Arrhenius base? (1) CH3OH (3) LiOH (2) CH3Cl (4) LiCl 12. Which two compounds are electrolytes? (1) C6H12O6 and CH3CH2OH (2) C6H12O6 and HCl (3) NaOH and HCl (4) NaOH and CH3CH2OH 13. The only positive ion found in H2SO4 (aq) is the (1) ammonium ion (3) hydroxide ion (2) hydronium ion (4) sulfate ion 14. A 25.0-milliliter sample of HNO3 (aq) is neutralized by 32.1 milliliters of 0.150 M KOH (aq). What is the molarity of the HNO3 (aq)? (1) 0.117 M (3) 0.193 M (2) 0.150 M (4) 0.300 M Base your answers to questions 15 and 16 on the information below. In a titration, 15.65 milliliters of a KOH(aq) solution exactly neutralized 10.00 milliliters of a 1.22 M HCl(aq) solution. 15. Complete the equation below for the titration reaction by writing the formula of each product. KOH (aq) + HCl (aq) --> _______ + ________ 16. In the space in your answer booklet, show a correct numerical setup for calculating the molarity of the KOH (aq) solution. Base your answers to questions 17 through 20 on the information below. The health of fish depends on the amount of oxygen dissolved in the water. A dissolved oxygen (DO) concentration between 6 parts per million and 8 parts per million is best for fish health. A DO concentration greater than 1 part per million is necessary for fish survival. Fish health is also affected by water temperature and concentrations of dissolved ammonia, hydrogen sulfide, chloride compounds, and nitrate compounds. Most freshwater fish thrive in water with a pH between 6.5 and 8.5. A student’s fish tank contains fish, green plants, and 3800 grams of fish-tank water with 2.7 × 10–2 gram of dissolved oxygen. Phenolphthalein tests colorless and bromthymol blue tests blue in samples of the fish-tank water. 17. Based on the test results for the indicators phenolphthalein and bromthymol blue, what is the pH range of the fish-tank water? [1] 18. When the fish-tank water has a pH of 8.0, the hydronium ion concentration is 1.0 × 10–8 mole per liter. What is the hydronium ion concentration when the water has a pH of 7.0? [1] 19. State how an increase in the temperature of the fish-tank water affects the solubility of oxygen in the water. [1] 20. Determine if the DO concentration in the fish tank is healthy for fish. Your response must include: • a correct numerical setup to calculate the DO concentration in the water in parts per million [1] • the calculated result [1] • a statement using your calculated result that tells why the DO concentration in the water is or is not healthy for fish [1] Soil pH can affect the development of plants. For example, a hydrangea plant produces blue flowers when grown in acidic soil but pink flowers when grown in basic soil. Evergreen plants can show a yellowing of foliage, called chlorosis, when grown in soil that is too basic. Acidic soil can be neutralized by treating it with calcium hydroxide, Ca(OH)2, commonly called slaked lime. Slaked lime is slightly soluble in water. 21. Compare the hydrogen ion concentration to the hydroxide ion concentration in soil when a hydrangea plant produces pink flowers. [1] 22. An evergreen plant has yellowing foliage. The soil surrounding the plant is tested with methyl orange and bromthymol blue. Both indicators turn yellow in the soil tests. State, in terms of pH value, why the yellowing of the plant is not due to chlorosis. [1] 23. Write an equation, using symbols or words, for the neutralization of the ions in acidic soil by the ions released by slaked lime in water. [1] A laboratory worker filled a bottle with a hydrochloric acid solution. Another bottle was filled with methanol, while a third bottle was filled with a sodium hydroxide solution. However, the worker neglected to label each bottle. After a few days, the worker could not remember which liquid was in each bottle. The worker needed to identify the liquid in each bottle. The bottles were labeled A, B, and C. Using materials found in the lab (indicators, conductivity apparatus, and pieces of Mg metal), the worker tested samples of liquid from each bottle. The test results are shown in the table below.  SHAPE \* MERGEFORMAT  24. Using the test results, state how the worker differentiated the bottle that contained methanol from the other two bottles. [1] 25. The worker concluded that bottle C contained hydrochloric acid. Identify one test and state the corresponding test result that supports this conclusion. [1] 26. Explain, in terms of pH, why the methyl orange indicator test results were the same for each of the three liquids. [1] IX. KINETICS AND EQUILIBRIUM 1. Systems in nature tend to undergo changes toward (1) lower energy and less disorder (2) lower energy and more disorder (3) higher energy and less disorder (4) higher energy and more disorder Which 1-mole sample has the least entropy? (1) Br2 (s) at 266 K (3) Br2 (l) at 332 K (2) Br2 (l) at 266 K (4) Br2 (g) at 332 K 3. At 20.°C, a1.2- gram sample of Mg ribbon reacts rapidly with 10.0 milliliters of 1.0 M HCl(aq). Which change in conditions would have caused the reaction to proceed more slowly? (1) increasing the initial temperature to 25°C (2) decreasing the concentration of HCl(aq) to 0.1 M (3) using 1.2 g of powdered Mg (4) using 2.4 g of Mg ribbon 4. A catalyst lowers the activation energy of a reaction by (1) providing an alternate reaction pathway (2) decreasing the heat of reaction (3) increasing the mass of the reactants (4) changing the mole ratio of the reactants 5. A reaction is most likely to occur when reactant particles collide with (1) proper energy, only (2) proper orientation, only (3) both proper energy and proper orientation (4) neither proper energy nor proper orientation 6. Given the balanced equation representing a reaction: Zn(s) + 2HCl(aq) ’! H2(g) + ZnCl2(aq) Which set of reaction conditions produces H2(g) at the fastest rate? (1) a 1.0-g lump of Zn(s) in 50. mL of 0.5 M HCl(aq) at 20.°C (2) a 1.0-g lump of Zn(s) in 50. mL of 0.5 M HCl(aq) at 30.°C (3) 1.0 g of powdered Zn(s) in 50. mL of 1.0 M HCl(aq) at 20.°C (4) 1.0 g of powdered Zn(s) in 50. mL of 1.0 M HCl(aq) at 30.°C 7. Which reaction releases the greatest amount of energy per 2 moles of product? (1) 2 CO (g) + O2 (g) ( 2 CO2 (g) (2) 4 Al (s) + 3 O2(g) ( 2 Al2O3 (s) (3) 2 H2 (g) + O2 (g) ( 2 H2O(g) (4) N2 (g) + 3 H2 (g) ( 2 NH3 (g) 8. Given the equation representing a phase change at equilibrium: C2H5OH(l) <-- --> C2H5OH(g) Which statement is true? (1) The forward process proceeds faster than the reverse process. (2) The reverse process proceeds faster than the forward process. (3) The forward and reverse processes proceed at the same rate. (4) The forward and reverse processes both stop. 9. For a given reaction, adding a catalyst increases the rate of the reaction by (1) providing an alternate reaction pathway that has a higher activation energy (2) providing an alternate reaction pathway that has a lower activation energy (3) using the same reaction pathway and increasing the activation energy (4) using the same reaction pathway and decreasing the activation energy 10. Given the equation representing a reaction at equilibrium: N2 (g) + 3H2 (g) --> 2 NH3 (g) + energy Which change causes the equilibrium to shift to the right? (1) decreasing the concentration of H2(g) (2) decreasing the pressure (3) increasing the concentration of N2(g) (4) increasing the temperature Given the balanced equation representing a reaction: 2 HCl(aq) + Na2S2O3 (aq) ’! S(s) + H2SO3(aq) + 2 NaCl(aq) Decreasing the concentration of Na2S2O3(aq) decreases the rate of reaction because the (1) activation energy decreases (3) frequency of effective collisions decreases (2) activation energy increases (4) frequency of effective collisions increases 12. The potential energy diagram for a chemical reaction is shown below.  SHAPE \* MERGEFORMAT  Each interval on the axis labeled “Potential Energy (kJ)” represents 40 kilojoules. What is the heat of reaction? (1) "120 kJ (2) "40 kJ (3) +40 kJ (4) +160 kJ Base your answers to questions 13 and 14 on the information below. The equilibrium equation below is related to the manufacture of a bleaching solution. In this equation, Cl (aq) means that chloride ions are surrounded by water molecules.  SHAPE \* MERGEFORMAT  13. Below, use the key to draw two water molecules in the box, showing the correct orientation of each water molecule toward the chloride ion. [1] 14. Explain, in terms of collision theory, why increasing the concentration of Cl2(g) increases the concentration of OCl–(aq) in this equilibrium system. [1] Base your answers to questions 14 through 17 on the reaction represented by the equation below. 2 H2(g) + O2 (g) ’! 2 H2O (l) + 571.6 kJ 15. Identify the information in this equation that indicates the reaction is exothermic. [1] 16. On the axes below, complete the potential energy diagram for the reaction represented by this equation.  SHAPE \* MERGEFORMAT  17. Explain why the entropy of the system decreases as the reaction proceeds. [1] 18. How much energy is released when 4.0 moles of H2O (l) are produced? Base your answers to questions 19 through 21 on a 100.0 mL sample of a dilute aqueous solution of ethanoic acid at equilibrium. The equation below represents this system. HC2H3O2 (aq) <-- --> C2H3O2- (aq) + H+ (aq) 19. Compare the rate of the forward reaction to the rate of the reverse reaction for this system. [1] 20. Describe what happens to the concentration of H+(aq) when 10 drops of concentrated HC2H3O2(aq) are added to this system. 21. In the space below, draw a structural formula for ethanoic acid. Base your answers to questions 22 through 24 on the information below. Nitrogen gas, hydrogen gas, and ammonia gas are in equilibrium in a closed container at constant temperature and pressure. The equation below represents this equilibrium. N2 (g) + 3 H2(g) <-- --> 2 NH3 (g) The graph below shows the initial concentration of each gas, the changes that occur as a result of adding H2 (g) to the system, and the final concentrations when equilibrium is reestablished.  SHAPE \* MERGEFORMAT  22. What information on the graph indicates that the system was initially at equilibrium? [1] 23. Explain, in terms of Le Chatelier’s principle, why the final concentration of NH3 (g) is greater than the initial concentration of NH3 (g). [1] 24. Explain, in terms of collision theory, why the concentration of H2 (g) begins to decrease immediately after more H2 (g) is added to the system. [1] Base your answers to questions 25 through 28 on the following information In a laboratory, 0.100 mole of colorless hydrogen iodide gas at room temperature is placed in a 1.00-liter flask. The flask is sealed and warmed, causing the HI (g) to start decomposing to H2 (g) and I2 (g). Then the temperature of the contents of the flask is kept constant. During this reaction, the contents of the flask change to a pale purple-colored mixture of HI(g), H2(g), and I2 (g). When the color of the mixture in the flask stops changing, the concentration of I2 (g) is determined to be 0.013 mole per liter. The relationship between concentration and time for the reactant and products is shown in the graph below.  SHAPE \* MERGEFORMAT  25. Write a balanced equation to represent the decomposition reaction occurring in the flask. [1] 26. State, in terms of concentration, evidence that indicates the system in the flask has reached equilibrium. [1] 27. Calculate the mass of I2(g) in the flask at equilibrium. Your response must include both a correct numerical setup and the calculated result. [2] 28. Explain, in terms of collision theory, why the rate of a chemical reaction increases with an increase in temperature. [1] X. REDOX 1. Which balanced equation represents a redox reaction? (1) CuCO3 (s) ’! CuO(s) + CO2 (g) (2) 2KClO3 (s) ’! 2KCl(s) + 3O2 (g) (3) AgNO3 (aq) + KCl (aq) ’! AgCl (s) + KNO3 (aq) (4) H2SO4 (aq) + 2 KOH (aq) ’! K2SO4 (aq) + 2 H2O (l) 2. Which energy conversion occurs in a voltaic cell? (1) chemical energy to electrical energy (3) electrical energy to chemical energy (2) chemical energy to nuclear energy (4) nuclear energy to electrical energy 3. Which metal is more active than Ni and less active than Zn? (1) Cu (2) Cr (3) Mg (4) Pb 4. Given the unbalanced ionic equation: 3Mg + ____ Fe3+ ’! 3 Mg2+ + ____ Fe When this equation is balanced, both Fe3+ and Fe have a coefficient of (1) 1, because a total of 6 electrons is transferred (3) 1, because a total of 3 electrons is transferred (2) 2, because a total of 6 electrons is transferred (4) 2, because a total of 3 electrons is transferred 5. Reduction occurs at the cathode in (1) electrolytic cells, only (3) both electrolytic cells and voltaic cells (2) voltaic cells, only (4) neither electrolytic cells nor voltaic cells 6. Given the balanced ionic equation representing a reaction: 2Al3+ (aq) + 3 Mg (s) (3Mg2+ (aq) + 2Al (s) In this reaction, electrons are transferred from (1) Al to Mg2+ (3) Mg to Al3+ (2) Al3+ to Mg (4) Mg2+ to Al 7. Which reaction occurs spontaneously? (1) Cl2 (g) + 2NaBr (aq) ’! Br2 (l) + 2 NaCl (aq) (2) Cl2 (g) + 2NaF (aq) ’! F2 (g) + 2 NaCl (aq) (3) I2 (s) + 2NaBr (aq) ’! Br2 (l) + 2 NaI (aq) (4) I2 (s) + 2NaF( aq) ’! F2 (g) + 2 NaI (aq) Given the balanced equation representing a reaction:  SHAPE \* MERGEFORMAT  During this reaction, the oxidation number of Fe changes from (1) +2 to 0 as electrons are transferred (2) +2 to 0 as protons are transferred (3) +3 to 0 as electrons are transferred (4) +3 to 0 as protons are transferred 9. Which half-reaction correctly represents reduction?  SHAPE \* MERGEFORMAT  Base your answers to questions 10 and 11 on the information below. Electroplating is an electrolytic process used to coat metal objects with a more expensive and less reactive metal. The diagram below shows an electroplating cell that includes a battery connected to a silver bar and a metal spoon. The bar and spoon are submerged in AgNO3(aq).  SHAPE \* MERGEFORMAT  10. Explain why AgNO3 is a better choice than AgCl for use in this electrolytic process. [1] 11. Explain the purpose of the battery in this cell. [1] Base your answers to questions 12 through 14 on the information below. A flashlight can be powered by a rechargeable nickel-cadmium battery. In the battery, the anode is Cd(s) and the cathode is NiO2(s). The unbalanced equation below represents the reaction that occurs as the battery produces electricity. When a nickel-cadmium battery is recharged, the reverse reaction occurs.  SHAPE \* MERGEFORMAT  12. Balance the equation below for the reaction that produces electricity, using the smallest whole-number coefficients. [1] ___ Cd + ___ NiO2 + ___ H2O --> ___ Cd(OH)2 + ___ Ni(OH)2 13. Determine the change in oxidation number for the element that makes up the anode in the reaction that produces electricity. [1] 14. Explain why Cd would be above Ni if placed on Table J. [1] Base your answers to questions 15 through 19 on the information below. A voltaic cell with magnesium and copper electrodes is shown in the diagram below. The copper electrode has a mass of 15.0 grams. When the switch is closed, the reaction in the cell begins. The balanced ionic equation for the reaction in the cell is shown below the cell diagram. After several hours, the copper electrode is removed, rinsed with water, and dried. At this time, the mass of the copper electrode is greater than 15.0 grams.  SHAPE \* MERGEFORMAT  15. State the direction of electron flow through the wire between the electrodes when the switch is closed. [1] 16. State the purpose of the salt bridge in this cell. [1] 17. Write a balanced half-reaction for oxidation. 18. Explain, in terms of copper ions and copper atoms, why the mass of the copper electrode increases as the cell operates. Your response must include information about both copper ions and copper atoms. [1] 19. Give one difference between a voltaic cell and an electrolytic cell. Your answer must mention both the voltaic cell and the electrolytic cell. 20. What is the oxidation number of nitrogen in KNO3 (g)? XI. ORGANIC CHEMISTRY 1. What is the name of the reaction between a fat and a base? 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Given the formulas for two compounds:  SHAPE \* MERGEFORMAT  These compounds differ in (1) gram-formula mass (3) percent composition by mass (2) molecular formula (4) physical properties at STP 3. The organic compound represented by the condensed structural formula CH3CH2CH2CHO is classified as an (1) alcohol (2) aldehyde (3) ester (4) ether 4. Which compound is an unsaturated hydrocarbon? (1) hexanal (2) hexanoic acid (3) hexane (4) hexyne 5. Which general formula represents the compound CH3CH2CCH?  SHAPE \* MERGEFORMAT  6. The isomers butane and methylpropane differ in their (1) molecular formulas (3) total number of atoms per molecule (2) structural formulas (4) total number of bonds per molecule 7. Which formula represents 2-butene?  SHAPE \* MERGEFORMAT  8. Which two compounds are isomers of each other?  SHAPE \* MERGEFORMAT  9. Given the formula of a substance:  SHAPE \* MERGEFORMAT  What is the total number of shared electrons in a molecule of this substance? (1) 22 (2) 11 (3) 9 (4) 6 Base your answers to questions 10 through 12 on the information below. The hydrocarbon 2-methylpropane reacts with iodine as represented by the balanced equation below. At standard pressure, the boiling point of 2-methylpropane is lower than the boiling point of 2-iodo-2-methylpropane.  SHAPE \* MERGEFORMAT  10. To which class of organic compounds does this organic product belong? [1] 11. Explain, in terms of bonding, why the hydrocarbon 2-methylpropane is saturated. [1] 12. Explain the difference in the boiling points of 2-methylpropane and 2-iodo-2-methylpropane in terms of both molecular polarity and intermolecular forces. [2] Base your answers to questions 13 through 16 on the information below. The incomplete equation below represents an esterification reaction. The alcohol reactant is represented by X. 13. On the structural formula above,circle the acid functional group, only. [1]  SHAPE \* MERGEFORMAT  14. Write an IUPAC name for the reactant represented by its structural formula in this equation. [1] 15. In the space below, draw the structural formula for the alcohol represented by X. [1] 16. Write an IUPAC name for the organic compound produced. Base your answers to questions 17 through 19 on the information below. Ozone gas, O3, can be used to kill adult insects in storage bins for grain without damaging the grain. The ozone is produced from oxygen gas, O2, in portable ozone generators located near the storage bins. The concentrations of ozone used are so low that they do not cause any environmental damage. This use of ozone is safer and more environmentally friendly than a method that used bromomethane, CH3Br. However, bromomethane was more effective than ozone because CH3Br killed immature insects as well as adult insects. Adapted From: The Sunday Gazette (Schenectady, NY) 3/9/03 17. Determine the total number of moles of CH3Br in 19 grams of CH3Br (gram-formula mass = 95 grams/mol). [1] 18. Given the balanced equation for producing bromomethane: Br2 + CH4 ’! CH3Br + HBr Identify the type of organic reaction shown. _________________________________ 19. Based on the information in the passage, state one advantage of using ozone instead of bromomethane for insect control in grain storage bins. [1] Base your answers to questions 20 through 23 on the information below. During a bread-making process, glucose is converted to ethanol and carbon dioxide, causing the bread dough to rise. Zymase, an enzyme produced by yeast, is a catalyst needed for this reaction. 20. Balance the equation in your answer booklet for the reaction that causes bread dough to rise, using the smallest whole-number coefficients. [1]  SHAPE \* MERGEFORMAT  21. In the space below draw a structural formula for the alcohol formed in this reaction. [1] 23. State the effect of zymase on the activation energy for this reaction. [1] Identify the type of organic reaction shown. ____________________________ Base your answers to questions 23 and 24 on the information below.  SHAPE \* MERGEFORMAT  23. Octane has a molar mass of 114 grams per mole. According to this graph, what is the boiling point of octane at standard pressure? [1] 24. State the relationship between molar mass and the strength of intermolecular forces for the selected alkanes. 25. 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