INTRODUCTION TO CHEMISTRY

1 INTRODUCTION TO CHEMISTRY Chapter Test A A. Multiple Choice Choose the best answer and write its letter on the line. 1. Identify the activity that belongs in the field of chemistry. a. developing medicines c. production of a new plastic b. analysis of a compound d. all of the above 2. Which of these chemicals is definitely inorganic? a. one that is made of carbon and hydrogen b. one that is made of nitrogen and carbon c. one that is made of nitrogen and hydrogen d. one that is made of carbon and oxygen 3. Which of the following is not a step for solving a numeric word problem? a. calculate c. analyze b. conclude d. evaluate 4. Identify the false statement. a. Chemistry plays an important role in efforts to increase the world's food supply and to protect crops. b. Biodiesel is a fossil fuel. c. Potato plants with a jellyfish gene will glow when they need to be watered. d. Chemists are working to develop more pest-resistant and disease-resistant plants. 5. A hypothesis is a. an observation recorded from an experiment. b. a proposed explanation for what is observed. c. a summary of the results of many experiments. d. a well-tested explanation for many observations. 6. Which of the following is not a part of the scientific method? a. experimenting c. proving b. observing d. hypothesizing 7. Identify the false statement. a. A scientific law fully explains a set of observations. b. The scientific method is a logical, systematic approach to the solution of a problem. c. For the results of an experiment to be accepted, the experiment must produce the same results no matter how many times it is repeated. d. The scientific method is repeated until a hypothesis either fits all the observed experimental results or the hypothesis is discarded. Chapter 1 Introduction to Chemistry 21

8. Which of these steps of the scientific method involves stating a relationship but not proposing an explanation for the relationship? a. observation c. theory b. scientific law d. hypothesis B. Questions Write your answer in the space provided. 9. Define chemistry. 10. List the major steps in the scientific method and describe each briefly. 22 Core Teaching Resources

C. Essay Write a short essay for the following. 11. What are the three steps that you can use to help solve numeric word problems? Describe what you would do in each step. Chapter 1 Introduction to Chemistry 23

2 MATTER AND CHANGE Chapter Test A A. Matching Match each description in Column B to the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 1. 2. 3. 4. 5. 6. product phase physical change liquid mass element a. matter that flows but has a definite volume b. amount of matter that an object contains c. starting substance in a chemical reaction d. homogeneous mixture e. the simplest form of matter with a unique set of properties f. a part of a sample with uniform composition and properties 7. 8. 9. solid solution compound 10. reactant B. Multiple Choice Choose the best answer and write its letter on the line provided. g. alters a substance without changing its composition h. matter with a definite shape and volume i. substance formed in a chemical reaction j. contains two or more elements chemically combined in a fixed proportion 11. All of the following are general characteristics of a substance in the liquid state except a. definite volume. c. not easily compressed. b. able to flow. d. definite shape. 12. In the chemical reaction iron plus oxygen y iron oxide, a. iron oxide is a reactant. c. oxygen is a product. b. iron is a reactant. d. iron is a product. 13. Which term does not fit with the others listed? a. solid c. gas b. reactant d. liquid Chapter 2 Matter and Change 45

14. The chemical symbol for sodium is a. NA. c. SO. b. Na. d. So. 15. A basketball has more mass than a golf ball because: a. the basketball takes up more space. b. the basketball contains more matter. c. the golf ball contains a different kind of matter. d. the golf ball has an indefinite composition. 16. A gas is a form of matter that a. has a definite volume. b. is generally a liquid or solid at room temperature. c. takes the shape and volume of its container. d. is difficult to compress. 17. Homogeneous mixtures a. are always liquids. b. consist of two or more phases. c. have a composition that is fixed. d. are known as solutions. 18. A compound a. is a pure substance. b. has a composition that varies. c. can be physically separated into its elements. d. has properties similar to those of its elements. 19. Physical properties of a substance include a. color and odor. c. malleability. b. melting and boiling points. d. all of the above. 20. When iron and oxygen combine to form iron oxide, a. a physical change occurs. c. a change in mass occurs. b. a change of state occurs. d. a chemical change occurs. C. True-False Classify each of these statements as always true, AT; sometimes true, ST; or never true NT. 21. Any part of a sample with uniform composition and properties is called a phase. 22. A substance does not have a fixed composition. 23. A heterogeneous mixture consists of two or more phases. 24. A vapor is a gaseous substance that is generally a liquid at room temperature. 25. A compound can be physically separated into its elements. 46 Core Teaching Resources

D. Completion Fill in the word(s) that will make each statement true. 26. During chemical and physical changes, substances can 26. absorb or give off 26. 27. During 27, a liquid is boiled to produce a vapor that 27. is then condensed again to a liquid. 28. Mixtures differ from substances because the composition 28. of a mixture can 28. 29. When iron and sulfur combine to form iron sulfide, a 29. 29 change takes place. 30. A 30 property is a quality of a substance that can 30. be observed or measured without changing the composition of the substance. 31. Elements combine chemically to form 31. 31. 32. Fe is the chemical symbol for the element 32. 32. 33. A 33 mixture has a uniform composition throughout. 33. 34. In any physical or chemical change, mass is 34. 34. 35. In a chemical reaction, the new substances formed are 35. called 35. E. Essay Write a short essay for the following. 36. Distinguish between physical changes and chemical changes and give two examples of each. Then, list three clues that are often indications of chemical changes. Chapter 2 Matter and Change 47

4 ATOMIC STRUCTURE Chapter Test A A. Matching Match each description in Column B with the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 1. 2. 3. 4. 5. 6. proton atom mass number atomic mass unit electron isotopes a. the total number of protons and neutrons in the nucleus of an atom b. the weighted average mass of the atoms in a naturally occurring sample of an element 1 c. 1 2 the mass of a carbon-12 atom d. the number of protons in the nucleus of an element e. atoms with the same number of protons but different numbers of neutrons f. negatively charged subatomic particle 7. 8. 9. atomic number atomic mass period 10. neutron B. Multiple Choice Choose the best answer and write its letter on the line. 11. Which of the following is not a part of Dalton s atomic theory? a. All elements are composed of atoms. b. Atoms of the same element are alike. c. Atoms are always in motion. d. Atoms that combine do so in simple whole-number ratios. 12. The nucleus of an atom is a. negatively charged and has a low density. b. negatively charged and has a high density. c. positively charged and has a low density. d. positively charged and has a high density. g. the smallest particle of an element that retains its identity in a chemical reaction h. a horizontal row of the periodic table i. subatomic particle with no charge j. positively charged subatomic particle Chapter 4 Atomic Structure 93

13. Dalton theorized that atoms are indivisible and that all atoms of an element are identical. Scientists now know that a. Dalton s theories are completely correct. b. atoms of an element can have different numbers of protons. c. atoms are all divisible. d. all atoms of an element are not identical but they all have the same mass. 14. The number of neutrons in the nucleus of an atom can be calculated by a. adding together the numbers of electrons and protons. b. subtracting the number of protons from the number of electrons. c. subtracting the number of protons from the mass number. d. adding the mass number to the number of protons. 15. The sum of the protons and neutrons in an atom equals the a. atomic number. c. atomic mass. b. number of electrons. d. mass number. 16. All atoms of the same element have the same: a. number of protons. c. mass number. b. number of neutrons. d. mass. 17. Which of these statements is false? a. Electrons have a negative charge. b. Electrons have a mass of 1 amu. c. The nucleus of an atom is positively charged. d. The neutron is found in the nucleus of an atom. 18. An atom of an element with atomic number 48 and mass number 120 contains a. 48 protons, 48 electrons, and 72 neutrons. b. 72 protons, 48 electrons, and 48 neutrons. c. 120 protons, 48 electrons, and 72 neutrons. d. 72 protons, 72 electrons, and 48 neutrons. 19. How do the isotopes hydrogen-2 and hydrogen-3 differ? a. Hydrogen-3 has one more electron than hydrogen-2. b. Hydrogen-3 has two neutrons. c. Hydrogen-2 has three protons. d. Hydrogen-2 has no protons. 20. The number 80 in the name bromine-80 represents a. the atomic number. b. the mass number. c. the sum of protons and electrons. d. none of the above 21. Which of these statements is not true? a. Atoms of the same elements can have different masses. b. The nucleus of an atom has a positive charge. c. Atoms of isotopes of an element have different numbers of protons. d. Atoms are mostly empty space. 94 Core Teaching Resources

22. Relative atomic masses are measured in a. nanograms. c. angstroms. b. grams. d. amus. 23. If E is the symbol for an element, which two of the following symbols represent isotopes of the same element? 1. 24 12E 2. 24 13E 3. a. 1 and 2 c. 1 and 4 b. 3 and 4 d. 2 and 3 25 11E 4. 25 12E C. Problems Solve the following problem in the space provided. Show your work. 24. There are five naturally occurring isotopes of the element zinc. The relative abundance and mass of each are as follows. 64 30Zn 48.89%, 63.929 amu 66 30Zn 27.81%, 65.926 amu 67 30Zn 4.11%, 66.927 amu 68 30Zn 18.57%, 67.925 amu 70 30Zn 0.62%, 69.925 amu Calculate the average atomic mass of zinc. Chapter 4 Atomic Structure 95

25. Complete this table. Atomic Mass Number of Number of Number of Number Number Protons Neutrons Electrons 9 10 14 7 21 20 13 27 56 26 26. List the number of protons, neutrons, and electrons in each of the following atoms. Protons Neutrons Electrons 13 6C 10 4Be 20 10 Ne 11 5B D. Essay 33 16 S Write a short essay to answer the following. 27. Explain how the atoms of one element differ from those of another element. Then explain how the atoms of one isotope differ from those of other isotopes of the same element. 96 Core Teaching Resources

5 ELECTRONS IN ATOMS Chapter Test A A. Matching Match each description in Column B with the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 1. 2. 3. 4. 5. 6. hertz atomic orbital spectrum atomic emission spectrum quantum of energy amplitude a. the number of wave cycles that pass a given point per unit of time b. the SI unit of frequency c. lines of colored light obtained by passing the light emitted by an element through a prism d. the region around an atomic nucleus where an electron is likely to be moving e. a range of colors seen when light passes through a prism f. Electrons enter orbitals of lowest energy first. 7. 8. 9. wavelength energy level Aufbau principle 10. frequency B. Multiple Choice Choose the best answer and write its letter on the line. g. a region outside the nucleus where there is high probability of finding an electron 11. The fourth principal energy level has a. 4 orbitals. c. 32 orbitals. b. 16 orbitals. d. 9 orbitals. h. the amount of energy required to move an electron from one energy level to the next higher one i. the distance between crests of waves j. the height of a wave from the origin to the crest 12. If the electron configuration of an element is 1s 2 2s 2 2p 6 3s 2 3p 5, the element is a. iron. c. chlorine. b. bromine. d. phosphorus. 116 Core Teaching Resources

13. The quantum mechanical model of the atom a. is concerned with the probability of finding an electron in a certain position. b. was proposed by Neils Bohr. c. defines the exact path of an electron around the nucleus. d. has many analogies in the visible world. 14. The electron configuration of calcium is a. 1s 2 2s 2 2p 2 3s 2 3p 3 4s 2. c. 1s 2 2s 2 3s 2 3p 6 3d 8. b. 1s 2 2s 2 2p 10 3s 2 3p 4. d. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2. 15. The maximum number of electrons that can occupy the third principal energy level is a. 18. c. 2. b. 32. d. 8. 16. As the frequency of light increases, the wavelength a. increases. c. decreases. b. remains the same. d. approaches the speed of light. 17. The formula 2n 2 represents a. the number of sublevels in any energy level. b. the maximum number of electrons that can occupy an energy level. c. the number of orbitals in a sublevel. d. none of the above 18. In order to occupy the same orbital, two electrons must have a. the same direction of spin. c. opposite charge. b. low energy. d. opposite spin. 19. Stable electron configurations are likely to contain a. high-energy electrons. b. unfilled s orbitals. c. fewer electrons than unstable configurations. d. filled energy sublevels. 20. According to Hund s rule, when electrons occupy orbitals of equal energy, one electron enters each orbit until a. all the orbitals contain one electron, with spins parallel. b. all the orbitals contain one electron, with opposite spins. c. there are two electrons in each orbital. d. electron velocities become constant. C. Problems Solve the following problems in the space provided. Show your work. 21. Write electron configurations for these atoms, using arrows to represent electrons. Then, use the shorthand method to write the configurations. a. S b. Na Chapter 5 Electrons in Atoms 117

22. What is the frequency of radiation whose wavelength is 2.40 10 5 cm? 23. Identify the elements that have the following electron configurations. a. 1s 2 2s 2 2p 6 3s 2 3p 1 a. b. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 b. c. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 7 4s 2 c. 24. Consider the elements neon, bromine, and phosphorus. Which has a. three electrons in its 3p sublevel? a. b. its highest energy level completely filled? b. c. the highest occupied energy level? c. 25. What is the maximum number of electrons that can occupy each of the first five principal energy levels? D. Essay Write a short essay for the following. 26. Explain why the 5s sublevel fills before any electrons are added to the 4d sublevel. 118 Core Teaching Resources

E. Additional Matching Match each description in Column B with the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 27. 28. 29. 30. 31. 32. ground state photoelectric effect photons Heisenberg uncertainty principle de Broglie s equation Planck s constant a. a process in which electrons are ejected by metals when certain wavelengths of light shine on them b. the lowest energy level for a given electron c. predicts that all matter exhibits wavelike motions d. 6.6262 10 34 Js e. It is impossible to know both the velocity and the position of a particle at the same time. f. quanta of light F. True-False Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT. 33. The energy of a body can change only in small discrete units. 34. The position and velocity of an electron in an atom can be determined with great certainty. 35. Above the threshold frequency, increasing the intensity of the light striking a metal increases the number of electrons ejected. 36. Electrons moving from one energy level to another emit photons. 37. The photoelectric effect will occur no matter what frequency of light strikes a metal. 38. In the Bohr model of the hydrogen atom, when the electron is in the ground state, n 1. 39. The lines in the emission spectrum of hydrogen can be correlated to transitions of electrons from higher energy levels to lower energy levels. 40. The outermost s orbital of an atom contains two electrons. 41. According to quantum mechanics, the motions of subatomic particles may be described as wavelike. Chapter 5 Electrons in Atoms 119

6 THE PERIODIC TABLE Chapter Test A A. Matching Match each description in Column B with the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 1. 2. 3. 4. 5. 6. metal halogen inner transition metal representative element ionization energy alkaline earth metal a. positively charged ion b. an element whose highest occupied s or p sublevels are partially filled c. the energy needed to remove an electron from an atom in the gaseous state d. good conductor of heat and electric current e. an element whose highest occupied s sublevel and a nearby d sublevel contain electrons f. any nonmetal in Group 7A 7. 8. 9. 10. cation electronegativity transition metal noble gas B. Multiple Choice Choose the best answer and write its letter on the line. g. the tendency of an atom to attract electrons when the atom is in a compound h. an element in which the highest occupied s and p sublevels are filled i. an element in Group 2A j. an element whose highest occupied s sublevel and a nearby f sublevel contain electrons 11. In the periodic table, there is a periodic pattern in the physical and chemical properties of elements when they are arranged in order of a. increasing atomic mass. b. increasing electronegativity. c. increasing atomic radius. d. increasing atomic number. Chapter 6 The Periodic Table 143

12. Which sublevel corresponds to the transition metals in the periodic table? a. s c. d b. p d. f 13. The representative elements are a. inner transition metals. c. Group B elements. b. transition metals. d. Group A elements. 14. Which of the following elements is a metalloid? a. As c. Br b. Se d. Kr 15. When a strontium atom loses two electrons to form an Sr 2 ion, the electrons are lost from the a. 5s orbital. c. 3d orbital. b. 5p orbital. d. 4f orbital. 16. The element iodine, I, is a a. period 5 alkali metal. c. period 5 halogen. b. period 4 halogen. d. period 5 transition metal. 17. The subatomic particle that plays the greatest role in determining the physical and chemical properties of an element is the a. proton. c. electron. b. neutron. d. photon. 18. Which of the following atoms would you expect to have the largest atomic radius? a. I c. Ca b. K d. Rb 19. From left to right across the second period of the periodic table, a. first ionization energy increases. c. electronegativity decreases. b. atomic radii increase. d. atomic mass decreases. 20. For which element would you expect a large jump between the first and second ionization energies? a. F c. Fe b. Ca d. Na 21. The category of elements that is characterized by the filling of f orbitals is the a. inner transition metals. c. alkali earth metals. b. alkali metals. d. transition metals. 22. Electronegativity a. generally decreases from left to right across a period. b. is the energy change that accompanies the loss of an electron from a gaseous atom. c. generally decreases from top to bottom within a group. d. is generally higher for metals than for nonmetals. 144 Core Teaching Resources

23. Atomic size generally a. increases from left to right across a period. b. decreases from top to bottom within a group. c. remains constant within a period. d. decreases from left to right across a period. 24. Of the following atoms, which one has the smallest first ionization energy? a. boron c. nitrogen b. aluminum d. silicon 25. The alkali metals do not include a. Li. c. Na. b. Ca. d. Rb. C. Questions Answer the following in the space provided. 26. For the elements whose electron configurations end as follows, state the period and the group to which each element belongs. a. 5s 2 b. 4s 2 4p 4 c. 6s 2 6p 2 27. Write the configurations for the highest occupied energy level for each of the following elements. a. He b. K c. Al d. Kr e. O 28. Arrange the following elements in order of decreasing ionization energy: a. Cs, Li, K b. Cl, Si, P, Ar c. Ca, Ba, Be, Sr Chapter 6 The Periodic Table 145

29. Tell whether each of the following elements is an inner transition metal, a noble gas, an alkali metal, an alkaline earth metal, or a halogen. Then give its period and group numbers. a. calcium b. cesium c. fluorine d. chromium e. neon f. silver 30. Which element in each pair has the higher electronegativity value? a. Na, Mg b. Rb, I c. Cl, Br D. Essay Write a short essay for the following. 31. Why does atomic size generally increase as you move down a group of the periodic table and decrease as you move from left to right across a period? 146 Core Teaching Resources

7 IONIC AND METALLIC BONDING Chapter Test A A. Matching Match each description in Column B with the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 1. 2. 3. 4. 5. electron dot structure ionic compound valence electron ionic bond chemical formula a. a mixture of two or more elements, at least one of which is a metal b. the number of ions of opposite charge surrounding each ion in a crystal c. the force of attraction binding oppositely charged ions together d. the attraction of valence electrons for positive metal ions e. a depiction of valence electrons around the symbol of an element 6. 7. 8. 9. 10. 11. halide ion alloy octet rule formula unit coordination number metallic bond B. Multiple Choice Choose the best answer and write its letter on the line. f. compound of cations and anions g. an anion of a halogen h. an electron in the highest occupied energy level of an atom i. Atoms in most compounds tend to achieve the electron configuration of a noble gas. j. shows the kinds and numbers of atoms in the smallest representative unit of a substance k. lowest whole-number ratio of ions in an ionic compound 12. How many valence electrons does an atom of any element in Group 6A have? a. 2 c. 6 b. 4 d. 8 168 Core Teaching Resources

13. The electron dot structure for an atom of phosphorus is a. =P= c. = b. =P= d. 14. When an aluminum atom loses its valence electrons, what is the charge on the resulting ion? a. 2 c. 3 b. 2 d. 1 15. The electron configuration of a fluoride ion, F, is a. 1s 2 2s 2 2p 5. b. the same as that of the neon atom. c. 1s 2 2s 2 2p 6 3s 1. d. the same as that of a potassium ion. 16. Metals are good conductors of electricity because they a. form crystal lattices. b. contain positive ions. c. contain mobile valence electrons. d. form ionic bonds. 17. In forming chemical bonds, atoms tend to attain a. a state of higher energy. b. the electron configuration of noble gas atoms. c. the electron configuration of halogen atoms. d. all of the above =P= P 18. An ionic compound is a. electrically neutral. c. composed of anions and cations. b. held together by ionic bonds. d. all of the above 19. Which of these is not a characteristic of most ionic compounds? a. solid at room temperature b. has a low melting point c. conducts an electric current when melted d. produced by reaction between metallic and nonmetallic elements 20. A metallic bond is a bond between a. valence electrons and positively charged metal ions. b. the ions of two different metals. c. a metal and nonmetal. d. none of the above 21. Which element when combined with chlorine would most likely form an ionic compound? a. lithium c. phosphorus b. carbon d. bromine 22. A cation is any atom or group of atoms with a. a positive charge. b. no charge. c. a negative charge. d. more electrons than the corresponding atoms. Chapter 7 Ionic and Metallic Bonding 169

23. The cation Fe 3 is formed when a. an atom of iron loses two electrons. b. an atom of zinc loses two electrons. c. an atom of iron loses three electrons. d. an atom of iron gains three electrons. C. True-False Classify each of these statements as always true, AT; sometimes true, ST; or never true, NT. 24. The chemical properties of an element are largely determined by the number of valence electrons the element has. 25. Fluorine and chlorine each have one valence electron. 26. The coordination number gives the total number of ions in a crystal. 27. Atoms acquire the stable electron structure of a noble gas by losing electrons. 28. An alloy is a mixture of two or more elements, of which at least one is a metal. 29. The crystal structure of ionic compounds such as sodium chloride is very unstable. 30. When melted, ionic compounds conduct electricity. 31. Metals are ductile because the cations in a piece of pure metal are insulated from one another by a sea of electrons. 32. Metal atoms are arranged in a face-centered cubic structure. 33. During the formation of ionic compounds, electrons are transferred from one atom to another. D. Questions Answer the following in the space provided. 34. Write electron dot structures for the atoms and ions of each of the following elements. Atoms Ions a. Ca b. Br c. Al 35. Write the formulas obtained when each of these atoms loses or gains valence electrons and becomes an ion. Tell whether each is a cation or an anion. a. Cl c. Na b. Be d. O 170 Core Teaching Resources

36. Write the complete electron configurations for the ions in problem 35. a. b. c. d. 37. Use electron dot structures to predict the structure of the ionic compound composed of aluminum and chlorine. 38. Write the electron configuration diagram that shows the transfer of electrons that takes place to form the compound sodium fluoride. Include the electron configurations of the ions formed. Which noble gas configuration does each ion have? E. Essay Write a short essay for the following. 39. Explain how scientists have used the concept of metallic bonding to account for many of the physical properties of metals, such as electrical conductivity and malleability. Chapter 7 Ionic and Metallic Bonding 171

9 CHEMICAL NAMES AND FORMULAS Chapter Test A A. Matching Match each description in Column B to the correct term in Column A. Write the letter of the correct description in the blank. Column A Column B 1. 2. 3. 4. 5. cation anion law of definite proportions acid base a. any atom or group of atoms with a negative charge b. When two elements combine to form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers. c. tightly bound group of atoms that behaves as a unit and carries a charge d. a compound that produces hydroxide ions when dissolved in water e. a compound composed of two different elements 6. 7. 8. 9. law of multiple proportions polyatomic ion monatomic ion binary compound B. Multiple Choice Choose the best answer and write its letter on the line. f. any atom or group of atoms with a positive charge g. In all samples of the same chemical compound, the elements are always combined in the same proportion by mass. h. a compound that produces hydrogen ions when dissolved in water i. consists of a single atom with a positive or negative charge 10. The correct name for the N 3 ion is the: a. nitrate ion. c. nitride ion. b. nitric ion. d. nitrite ion. 228 Core Teaching Resources

11. Elements of Group 4A: a. generally form positive ions. b. generally form negative ions. c. do not commonly form ions. d. do not combine with other elements. 12. What is the ionic charge on the chromium ion in the ionic compound that has the formula Cr 2 O 3? a. 3 c. 5 b. 2 d. 5 13. Which element when combined with chlorine would most likely form an ionic compound? a. lithium c. phosphorus b. carbon d. bromine 14. What is the formula for calcium hydrogen phosphate? a. CaHPO 4 c. Ca(H 2 PO 4 ) 2 b. Ca 2 HPO 4 d. Ca(HPO 4 ) 2 15. A cation is any atom or group of atoms with: a. a positive charge. b. no charge. c. a negative charge. d. more electrons than the corresponding atoms. 16. The cation Fe 3 is formed when: a. an atom of iron loses two electrons. b. an atom of zinc loses two electrons. c. an atom of iron loses three electrons. d. an atom of iron gains three electrons. 17. A molecular formula: a. gives information about molecular geometry. b. can be written for ionic compounds. c. shows the number and kinds of atoms in a molecule of a compound. d. uses superscripts to show the number of atoms of each kind. 18. The metals in Groups 1A, 2A, and 3A: a. gain electrons when they form ions. b. form ions with a charge found by subtracting 8 from the group number. c. all form ions with a 1 charge. d. lose electrons when they form ions. 19. When naming an ion of a transition metal that has more than one common ionic charge, the numerical value of the charge is indicated by a: a. prefix. b. suffix. c. Roman numeral following the name. d. superscript after the name. Chapter 9 Chemical Names and Formulas 229

20. In naming a binary molecular compound, the number of atoms of each element present in the molecule is indicated by: a. Roman numerals. c. prefixes. b. superscripts. d. suffixes. 21. An -ite or -ate ending on the name of a compound indicates that the compound: a. is a binary ionic compound. b. is a binary molecular compound. c. contains a polyatomic anion. d. contains a polyatomic cation. 22. What is the formula for sulfuric acid? a. H 2 S 2 c. H 2 SO 3 b. H 2 SO 4 d. H 2 S C. Completion Fill in the word(s) that will make each statement true. 23. The ionic charge of lead, Pb, in the compound PbS 2 is 23. 23. 24. The typical ionic charge of an ion formed by an element in 24. Group 7A is 24. 25. Atoms that have a positive or negative charge are called 25. 25. 26. The name of a monatomic anion ends in 26. 26. 27. The ionic charge of chlorine, Cl, in the compound MgCl 2 is 27. 27. 28. In a polyatomic ion, the -ite ending indicates one fewer 28 28. atom than the -ate ending. 29. Binary molecular compounds are composed of two 29 29. elements. 30. Acids are compounds that produce 30 ions when dissolved 30. in water. D. Problems Write the answers in the space provided. 31. Write the formulas for these compounds. a. magnesium cyanide a. b. mercury(ii) bromide b. c. sulfur hexafluoride c. 230 Core Teaching Resources

32. Name these compounds. a. CuCl a. b. N 2 O 3 b. c. KC 2 H 3 O 2 c. E. Essay Write a short essay for the following. 33. Why was it necessary for chemists to develop a system for naming chemical compounds? Chapter 9 Chemical Names and Formulas 231

10 CHEMICAL QUANTITIES Chapter Test A A. Matching Match each description in Column B to the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 1. 2. 3. 4. 5. 6. standard temperature and pressure molar volume representative particle molar mass atomic mass Avogadro s number a. the volume occupied by a mole of any gas at STP (22.4 L) b. the mass in amus of an atom of an element c. the number of representative particles in a mole of a substance d. the percent by mass of each element in a compound e. the mass of a mole of a substance f. the formula that gives the lowest whole-number ratio of the elements in a compound 7. 8. 9. 10. mole (mol) empirical formula percent composition molar mass B. Multiple Choice Choose the best answer and write its letter on the line. g. an atom, a formula unit, or a molecule h. the mass in grams of a mole of any pure substance i. 6.02 10 23 representative particles of a substance j. 0 C and 101.3 kpa 11. How many hydrogen atoms are in 4 molecules of isopropyl alcohol, C 3 H 8 O? a. 4 (6.02 10 23 ) c. 32 b. 4 d. 32 (6.02 10 23 ) 12. The mass of a mole of NaCl is its: a. molar mass. c. gram atomic mass. b. empirical formula. d. atomic mass. Chapter 10 Chemical Quantities 253

13. The chemical formula of aspirin is C 9 H 8 O 4. What is the mass of 0.200 mol of aspirin? a. 22.5 g c. 80 g b. 5.4 g d. 36.0 g 14. How many moles of SO 3 are in 2.4 10 24 molecules of SO 3? a. 0.25 c. 4.0 b. 3.4 10 22 d. 2.9 10 23 15. What is the volume (in liters at STP) of 2.50 mol of carbon monoxide? a. 0.112 L c. 56.0 L b. 3.10 L d. 8.96 L 16. The molar mass of molecular oxygen is: a. equal to the mass of one mole of oxygen atoms. b. 16.0 g. c. 32.0 g. d. none of the above 17. At STP, 1 mol each of hydrogen, oxygen, nitrogen, and fluorine: a. have different densities. b. occupy different volumes. c. are monatomic elements. d. contain twice Avogadro s number of representative particles. 18. A large weather balloon filled with helium has a volume of 7.00 10 2 L at STP. Which expression should be used to find the mass of helium in the balloon? a. 2 2.4 m L ol 4 g He 22. 4 L/ m l m c. ol 7. 00 10 ol 2 4 g He mol b. 7.00 10 2 4 g L m He ol d. 7. 00 10 2 2 2. 4 L/ mol 4 g He l mol 19. The molar mass of a gas can be determined from: a. the density of the gas at STP. b. the volume of a mole of the gas. c. Avogadro s number. d. none of the above 20. What is the percent by mass of carbon in acetone, C 3 H 6 O? a. 1.60% c. 62% b. 30.0% d. 20.7% 21. Which of the following is not an empirical formula? a. Na 2 SO 4 c. N 2 H 4 b. C 6 H 5 Cl d. Sn 3 (PO 4 ) 4 22. Which expression represents the percent by mass of nitrogen in NH 4 NO 3? a. 1 4. 0 g 100 c. 8 0. 0 g 100 80. 0 g 14. 0 g b. 2 8. 0 g 100 d. 8 0. 0 g 100 80. 0 g 28. 0 g 254 Core Teaching Resources

23. A compound has the empirical formula C 2 H 3 O and a molar mass of 172 g. What is its molecular formula? a. C 6 H 9 O 3 c. C 9 H 16 O 3 b. C 8 H 12 O 4 d. C 7 H 8 O 5 24. The empirical formula of a compound is CH 2 F. The molar mass of this compound is 66.0 g. The molecular formula of the compound is: a. C 4 H 8 F 4. c. C 2 H 4 F 2. b. C 4 H 4 F 4. d. CH 2 F. C. Problems Solve the following problems in the space provided. Show your work. 25. Find the mass, in grams, of 5.00 10 23 molecules of F 2. 26. The polymer used for the nonstick surface of cooking utensils is 24.0% C and 76.0% F by mass. What is the empirical formula of this polymer? 27. Find the number of moles of argon in 364 g of argon. D. Essay Write a short essay for the following. 28. The adjective molar in the phrase molar mass means of or pertaining to a mole. How does this definition support your understanding of molar mass as it is applied to atoms, molecules, ions, formula units, and other particles? Chapter 10 Chemical Quantities 255

11 CHEMICAL REACTIONS Chapter Test A A. Matching Match each description in Column B with the correct term in Column A. Write the letter of the correct description on the line. Column A Column B 1. 2. 3. 4. 5. 6. coefficient spectator ion combustion reaction reactant skeleton equation balanced equation a. a chemical equation that does not indicate relative amounts of reactants and products b. a starting substance in a chemical reaction c. a list of metals in order of decreasing reactivity d. a reaction in which a single compound is broken down into simpler substances e. a whole number that appears before a formula in an equation f. a new substance formed in a chemical reaction 7. 8. 9. activity series product decomposition reaction 10. single-replacement reaction B. Multiple Choice Choose the best answer and write its letter on the line. 11. The symbol in a chemical equation means: a. heat is supplied to the reaction. b. a catalyst is needed. c. yields. d. precipitate. g. a particle not directly involved in a chemical reaction h. an equation in which each side has the same number of atoms of each element i. a reaction in which the atoms of one element replace the atoms of a second element j. a reaction in which oxygen reacts with another substance, often producing heat or light 12. A catalyst is: a. a solid product of a reaction. b. one of the reactants in a single-replacement reaction. c. not used up in a reaction. d. the product of a combustion reaction. Chapter 11 Chemical Reactions 277

13. When the equation Fe O 2 y Fe 2 O 3 is balanced, the coefficient for O 2 is: a. 4 c. 2 b. 3 d. 1 14. The reaction in question 13 is an example of a(n): a. aqueous reaction. c. combination reaction. b. single-replacement reaction. d. decomposition reaction. 15. The equation H 3 PO 4 3KOH y K 3 PO 4 3H 2 O is an example of which type of reaction? a. double-replacement c. decomposition b. combination d. single-replacement 16. In a double-replacement reaction: a. the reactants are usually a metal and a nonmetal. b. the reactants are generally two ionic compounds in aqueous solution. c. one of the reactants is often water. d. energy in the form of heat or light is always produced. 17. In order for the reaction 2Al(s) 6HCl(aq) y 2AlCl 3 (aq) 3H 2 (g) to occur, which of the following must be true? a. Al must be above Cl on the activity series. b. Al must be above H on the activity series. c. Heat must be supplied for the reaction. d. AlCl 3 must be a precipitate. 18. In a double-replacement reaction: a. one of the products is always a gas. b. one of the products must be an element. c. positive ions are exchanged between two compounds. d. all of the above 19. When the following equation is balanced, the coefficient in front of HCl is: Ba(s) HCl(aq) y BaCl 2 (aq) H 2 (g) a. 6 c. 1 b. 3 d. 2 20. The equation in question 19 is an example of which type of reaction? a. combustion c. decomposition b. single-replacement d. double-replacement 21. This symbol 1 indicates: a. that heat must be applied. b. an incomplete combustion reaction. c. that a gas is formed by the reaction. d. that the reaction is reversible. 278 Core Teaching Resources

22. Which of the following statements is incorrect? a. Complete combustion has occurred when all the carbon in the product is in the form of carbon dioxide. b. A single reactant is the identifying characteristic of a decomposition reaction. c. The only way to determine the products of a reaction is to perform the reaction. d. All chemical reactions can be classified as one of four general types. 23. A chemical formula written above or below the yield sign indicates: a. that a gas is formed. b. that the substance is used as a catalyst. c. that heat must be supplied. d. a reversible reaction. 24. The equation 2C 3 H 7 OH(g) 9O 2 (g) y 6CO 2 (g) 8H 2 O(g) is an example of which type of reaction? a. combustion c. double-replacement b. combination d. decomposition 25. A double-replacement reaction takes place when aqueous K 2 SO 4 reacts with aqueous Pb(NO 3 ) 2. You would expect one of the products of this reaction to be: a. K 2 S. c. PbSO 4. b. NaPb. d. CNO 3. C. Problems Solve the following problems in the space provided. Show your work. 26. Balance the following equations: a. Ca(s) H 3 PO 4 (aq) y Ca 3 (PO 4 ) 2 (s) H 2 (g) b. KBrO 3 (s) y KBr(s) O 2 (g) c. (NH 4 ) 2 CO 3 (aq) NaOH(aq) y Na 2 CO 3 (aq) NH 3 (g) H 2 O(l) 27. For each equation in question 26, identify the type of reaction. a. b. c. Chapter 11 Chemical Reactions 279

28. Balance the following equations. Indicate whether combustion is complete or incomplete. a. C 5 H 10 (g) O 2 (g) y CO(g) H 2 O(g) b. C 3 H 7 OH(l) O 2 (g) y CO 2 (g) H 2 O(g) 29. Complete and balance the following equations. Then write each as a net ionic equation. a. K 3 PO 4 (aq) MgCl 2 (aq) y b. Fe(NO 3 ) 3 (aq) Na 2 CO 3 (aq) y D. Essay Write a short essay for the following. 30. What determines whether one metal will replace another metal from a compound in a single-replacement reaction? 280 Core Teaching Resources