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Page 68
Chapter Summary
Section 2.1
  • Dalton's atomic theory states that all matter is made up of tiny indivisible, immutable particles called atoms. Compounds form, moreover, when atoms of different elements combine in fixed ratios. According to the law of definite proportions, any sample of a given compound will always contain the same elements in the same mass ratio.

  • The law of multiple proportions states that if two elements can form more than one compound with one another, the mass ratio of one will be related to the mass ratio of the other by a small whole number.

  • The law of conservation of mass states that matter can be neither created nor destroyed.

Section 2.2
  • On the basis of Dalton's atomic theory, the atom is the basic unit of an element. Studies with radiation indicated that atoms contained subatomic particles, one of which was the electron.

  • Experiments with radioactivity have shown that some atoms give off different types of radiation, called alpha (α) rays, beta (β) rays, and gamma (γ) rays. Alpha rays are composed of α particles, which are actually helium nuclei. Beta rays are composed of β particles, which are actually electrons. Gamma rays are high-energy radiation.

  • Most of the mass of an atom resides in a tiny, dense region known as the nucleus. The nucleus contains positively charged particles called protons and electrically neutral particles called neutrons. The charge on a proton is equal in magnitude but opposite in sign to the charge on an electron. The electrons occupy the relatively large volume around the nucleus. A neutron has a slightly greater mass than a proton, but each is almost 2000 times as massive as an electron.

Section 2.3
  • The atomic number (Z) is the number of protons in the nucleus of an atom. The atomic number determines the identity of the atom. The mass number (A) is the sum of the protons and neutrons in the nucleus.

  • Protons and neutrons are referred to collectively as nucleons.

  • Atoms with the same atomic number but different mass numbers are called isotopes.

Section 2.4
  • The periodic table arranges the elements in rows (periods) and columns (groups or families). Elements in the same group exhibit similar properties.

  • All elements fall into one of three categories: metal, nonmetal, or metalloid.

  • Some of the groups have special names including alkali metals (Group 1A, except hydrogen), alkaline earth metals (Group 2A), chalcogens (Group 6A), halogens (Group 7A), noble gases (Group 8A), and transition elements or transition metals (Group 1B and Groups 3B–8B).

Section 2.5
  • Atomic mass is the mass of an atom in atomic mass units. One atomic mass unit (amu), is exactly one-twelfth the mass of a carbon-12 atom.

  • The periodic table contains the average atomic mass (sometimes called the atomic weight) of each element.

Section 2.6
  • A molecule is an electrically neutral group of two or more atoms. Molecules consisting of just two atoms are called diatomic. Diatomic molecules may be homonuclear (just one kind of atom) or heteronuclear (two kinds of atoms). In general, molecules containing more than two atoms are called polyatomic.

  • A chemical formula denotes the composition of a substance. A molecular formula specifies the exact numbers of atoms in a molecule of a compound. A structural formula shows the arrangement of atoms in a substance.

  • An allotrope is one of two or more different forms of an element.

  • Molecular compounds are named according to a set of rules, including the use of Greek prefixes to specify the number of each kind of atom in the molecule.

  • Binary compounds are those that consist of two elements. An acid is a substance that generates hydrogen ions when it dissolves in water. An ionizable hydrogen atom is one that can be removed in water to become a hydrogen ion (H+).

  • Inorganic compounds are generally those that do not contain carbon. Organic compounds contain carbon and hydrogen, sometimes in combination with other elements. Hydrocarbons contain only carbon and hydrogen. The simplest hydrocarbons are the alkanes. A functional group is a group of atoms that determines the chemical properties of an organic compound.

  • Empirical formulas express, in the smallest possible whole numbers, the ratio of the combination of atoms of the elements in a compound. The empirical and molecular formulas of a compound may or may not be identical.

Section 2.7
  • An ion is an atom or group of atoms with a net charge. An atomic ion or a monatomic ion consists of just one atom.

  • An ion with a net positive charge is a cation. An ion with a net negative charge is an anion. An ionic compound is one that consists of cations and anions in an electrically neutral combination. A three-dimensional array of alternating cations and anions is called a lattice.

  • Ionic compounds are named using rules similar to those for molecular compounds. In general, prefixes are not used to denote the number of ions in the names of ionic compounds.

  • Polyatomic ions are those that contain more than one atom chemically bonded together. Oxoanions are polyatomic ions that contain one or more oxygen atoms.

  • Oxoacids are acids based on oxoanions. Acids with more than one ionizable hydrogen atom are called polyprotic.

  • Hydrates are compounds whose formulas include a specific number of water molecules.

Page 69
Key Words
Questions and Problems
Section 2.1: The Atomic Theory
Review Questions
  • 2.1

    What are the hypotheses on which Dalton's atomic theory is based?

  • 2.2

    State the laws of definite proportions and multiple proportions. Illustrate each with an example.

  • 2.3

    The elements nitrogen and oxygen can form a variety of different compounds. Two such compounds, NO and N2O4, were decomposed into their constituent elements. One produced 0.8756 g N for every gram of O; the other produced 0.4378 g N for every gram of O. Show that these results are consistent with the law of multiple proportions.

  • 2.4

    Two different compounds, each containing only phosphorus and chlorine, were decomposed into their constituent elements. One produced 0.2912 g P for every gram of Cl; the other produced 0.1747 g P for every gram of Cl. Show that these results are consistent with the law of multiple proportions.

  • 2.5

    Sulfur reacts with fluorine to produce three different compounds. The mass ratio of fluorine to sulfur for each compound is given in the following table:

    Compound mass F : mass S
    S2F10 2.962
    SF4 2.370
    SF6 3.555

    Show that these data are consistent with the law of multiple proportions.

  • 2.6

    Both FeO and Fe2O3 contain only iron and oxygen. The mass ratio of oxygen to iron for each compound is given in the following table:

    Compound mass O : mass Fe
    FeO 0.2865
    Fe2O3 0.4297

    Show that these data are consistent with the law of multiple proportions.

Problems
  • 2.7

    For the two compounds pictured, evaluate the following ratio:

  • 2.8

    For the two compounds pictured, evaluate the following ratio:

Section 2.2: The Structure of the Atom
Review Questions
  • 2.9

    Define the following terms: (a) α particle, (b) β particle, (c) γ ray, (d) X ray.

  • 2.10

    Name the types of radiation known to be emitted by radioactive elements.

  • 2.11

    Compare the properties of the following: α particles, cathode rays, protons, neutrons, and electrons.

  • 2.12

    Describe the contributions of the following scientists to our knowledge of atomic structure: J. J. Thomson, R. A. Millikan, Ernest Rutherford, and James Chadwick.

  • 2.13

    Describe the experimental basis for believing that the nucleus occupies a very small fraction of the volume of the atom.

Problems
    Page 70
  • 2.14

    The diameter of a neutral helium atom is about 1 × 102 pm. Suppose that we could line up helium atoms side by side in contact with one another. Approximately how many atoms would it take to make the distance 1 cm from end to end?

  • 2.15

    Roughly speaking, the radius of an atom is about 10,000 times greater than that of its nucleus. If an atom were magnified so that the radius of its nucleus became 2.0 cm, about the size of a marble, what would be the radius of the atom in miles (1 mi = 1609 m)?

    Answer

Section 2.3: Atomic Number, Mass Number, and Isotopes
Review Questions
  • 2.16

    Use the helium-4 isotope to define atomic number and mass number. Why does knowledge of the atomic number enable us to deduce the number of electrons present in an atom?

  • 2.17

    Why do all atoms of an element have the same atomic number, although they may have different mass numbers?

  • 2.18

    What do we call atoms of the same elements with different mass numbers?

  • 2.19

    Explain the meaning of each term in the symbol .

Problems
  • 2.20

    What is the mass number of an iron atom that has 28 neutrons?

  • 2.21

    Calculate the number of neutrons of 239Pu.

    Answer

  • 2.22

    For each of the following species, determine the number of protons and the number of neutrons in the nucleus: , , , , , , .

  • 2.23

    Indicate the number of protons, neutrons, and electrons in each of the following species: , , , , , , .

    Answer

  • 2.24

    Write the appropriate symbol for each of the following isotopes: (a) Z = 11, A = 23; (b) Z = 28, A = 64; (c) Z = 50, A = 115; (d) Z = 20, A = 42.

  • 2.25

    Write the appropriate symbol for each of the following isotopes: (a) Z = 74, A = 186; (b) Z = 80, A = 201; (c) Z = 34, A = 76; (d) Z = 94, A = 239.

    Answer

  • 2.26

    Determine the mass number of (a) a boron atom with 5 neutrons, (b) a magnesium atom with 14 neutrons, (c) a bromine atom with 46 neutrons, and (d) a mercury atom with 116 neutrons.

  • 2.27

    Determine the mass number of (a) a fluorine atom with 10 neutrons, (b) a sulfur atom with 18 neutrons, (c) an arsenic atom with 42 neutrons, and (d) a platinum atom with 114 neutrons.

    Answer

  • 2.28

    The following radioactive isotopes are used in medicine for imaging organs, studying blood circulation, treating cancer, and so on. Give the number of neutrons present in each isotope: 198Au, 47Ca, 60Co, 18F, 125I, 131I, 42K, 43K, 24Na, 32P, 85Sr, 99Tc.

Section 2.4: The Periodic Table
Review Questions
  • 2.29

    What is the periodic table, and what is its significance in the study of chemistry?

  • 2.30

    State two differences between a metal and a nonmetal.

  • 2.31

    Write the names and symbols for four elements in each of the following categories: (a) nonmetal, (b) metal, (c) metalloid.

  • 2.32

    Give two examples of each of the following: (a) alkali metals, (b) alkaline earth metals, (c) halogens, (d) noble gases, (e) chalcogens, (f) transition metals.

  • 2.33

    The explosion of an atomic bomb in the atmosphere releases many radioactive isotopes into the environment. One of the isotopes is 90Sr. Via a relatively short food chain, it can enter the human body. Considering the position of strontium in the periodic table, explain why it is particularly harmful to humans.

Problems
  • 2.34

    Elements whose names end with –ium are usually metals; sodium is one example. Identify a nonmetal whose name also ends with –ium.

  • 2.35

    Describe the changes in properties (from metals to nonmetals or from nonmetals to metals) as we move (a) down a periodic group and (b) across the periodic table from left to right.

    Answer

  • 2.36

    Consult a handbook of chemical and physical data (ask your instructor where you can locate a copy of the handbook) to find (a) two metals less dense than water, (b) two metals more dense than mercury, (c) the densest known solid metallic element, and (d) the densest known solid nonmetallic element.

  • 2.37

    Group the following elements in pairs that you would expect to show similar chemical properties: K, F, P, Na, Cl, and N.

    Answer

  • 2.38

    Group the following elements in pairs that you would expect to show similar chemical properties: I, Ba, O, Br, S, and Ca.

  • 2.39

    Write the symbol for each of the following biologically important elements in the given periodic table: iron (present in hemoglobin for transporting oxygen), iodine (present in the thyroid gland), sodium (present in intracellular and extracellular fluids), phosphorus (present in bones and teeth), sulfur (present in proteins), and magnesium (present in chlorophyll molecules).

    Answer

Section 2.5: The Atomic Mass Scale and Average Atomic Mass
Review Questions
  • 2.40

    What is an atomic mass unit? Why is it necessary to introduce such a unit?

  • 2.41

    What is the mass (in amu) of a carbon-12 atom? Why is the atomic mass of carbon listed as 12.01 amu in the table on the inside front cover of this book?

  • 2.42

    Explain clearly what is meant by the statement “The atomic mass of gold is 197.0 amu.”

  • 2.43

    What information would you need to calculate the average atomic mass of an element?

Page 71
Problems
  • 2.44

    The atomic masses of (75.53 percent) and (24.47 percent) are 34.968 and 36.956 amu, respectively. Calculate the average atomic mass of chlorine. The percentages in parentheses denote the relative abundances.

  • 2.45

    The atomic masses of 204Pb (1.4 percent), 206Pb (24.1 percent), 207Pb (22.1 percent), and 208Pb (52.4 percent) are 203.973020, 205.974440, 206.975872, and 207.976627 amu, respectively. Calculate the average atomic mass of lead. The percentages in parentheses denote the relative abundances.

    Answer

  • 2.46

    The atomic masses of 203Tl and 205Tl are 202.972320 and 204.974401 amu, respectively. Calculate the natural abundances of these two isotopes. The average atomic mass of thallium is 204.4 amu.

  • 2.47

    The atomic masses of 6Li and 7Li are 6.0151 amu and 7.0160 amu, respectively. Calculate the natural abundances of these two isotopes. The average atomic mass of Li is 6.941 amu.

    Answer

  • 2.48

    What is the mass in grams of 13.2 amu?

  • 2.49

    How many atomic mass units are there in 8.4 g?

    Answer

Section 2.6: Molecules and Molecular Compounds
Review Questions
  • 2.50

    What is the difference between an atom and a molecule?

  • 2.51

    What are allotropes? Give an example. How are allotropes different from isotopes?

  • 2.52

    Describe the two commonly used molecular models.

  • 2.53

    What does a chemical formula represent? Determine the ratio of the atoms in the following molecular formulas: (a) NO, (b) NCl3, (c) N2O4, (d) P4O6.

  • 2.54

    Define molecular formula and empirical formula. What are the similarities and differences between the empirical formula and molecular formula of a compound?

  • 2.55

    Give an example of a case in which two molecules have different molecular formulas but the same empirical formula.

  • 2.56

    What is the difference between inorganic compounds and organic compounds?

  • 2.57

    Give one example each for a binary compound and a ternary compound. (A ternary compound is one that contains three different elements.)

  • 2.58

    Explain why the formula HCl can represent two different chemical systems.

Problems
  • 2.59

    For each of the following diagrams, determine whether it represents diatomic molecules, polyatomic molecules, molecules that are not compounds, molecules that are compounds, or an elemental form of the substance.

    Answer

  • 2.60

    For each of the following diagrams, determine whether it represents diatomic molecules, polyatomic molecules, molecules that are not compounds, molecules that are compounds, or an elemental form of the substance.

  • 2.61

    Identify the following as elements or compounds: NH3, N2, S8, NO, CO, CO2, H2, SO2.

    Answer

  • 2.62

    Give two examples of each of the following: (a) a diatomic molecule containing atoms of the same element, (b) a diatomic molecule containing atoms of different elements, (c) a polyatomic molecule containing atoms of the same element, (d) a polyatomic molecule containing atoms of different elements.

  • 2.63

    Write the empirical formulas of the following compounds: (a) C2N2, (b) C6H6, (c) C9H20, (d) P4O10, (e) B2H6.

    Answer

  • 2.64

    Write the empirical formulas of the following compounds: (a) Al2Br6, (b) Na2S2O4, (c) N2O5, (d) K2Cr2O7, (e) H2C2O4.

  • 2.65

    Write the molecular formula of alanine, an amino acid used in protein synthesis. The color codes are black (carbon), blue (nitrogen), red (oxygen), and white (hydrogen).

    Answer

  • 2.66

    Write the molecular formula of ethanol. The color codes are black (carbon), red (oxygen), and white (hydrogen).

  • 2.67

    Name the following binary molecular compounds: (a) NCl3, (b) IF7, (c) P4O6, (d) S2Cl2.

    Answer

  • 2.68

    Write chemical formulas for the following molecular compounds: (a) phosphorus tribromide, (b) dinitrogen tetrafluoride, (c) xenon tetroxide, (d) selenium trioxide.

  • Page 72
  • 2.69

    Write the molecular formulas and names of the following compounds.

    Answer

  • 2.70

    Write the molecular formulas and names of the following compounds.

Section 2.7: Ions and Ionic Compounds
Review Questions
  • 2.71

    Give an example of each of the following: (a) a monatomic cation, (b) a monatomic anion, (c) a polyatomic cation, (d) a polyatomic anion.

  • 2.72

    What is an ionic compound? How is electrical neutrality maintained in an ionic compound?

  • 2.73

    Explain why the chemical formulas of ionic compounds are usually the same as their empirical formulas.

  • 2.74

    What is the Stock system? What are its advantages over the older system of naming cations?

Problems
  • 2.75

    Give the number of protons and electrons in each of the following common ions: Na+, Ca2+, Al3+, Fe2+, I, F, S2–, O2–, N3–.

    Answer

  • 2.76

    Give the number of protons and electrons in each of the following common ions: K+, Mg2+, Fe3+, Br, Mn2+, C4–, Cu2+.

  • 2.77

    Write the formulas for the following ionic compounds: (a) sodium oxide, (b) iron sulfide (containing the Fe2+ ion), (c) cobalt sulfate (containing the Co3+ and ions), (d) barium fluoride.

    Answer

  • 2.78

    Write the formulas for the following ionic compounds: (a) copper bromide (containing the Cu+ ion), (b) manganese oxide (containing the Mn3+ ion), (c) mercury iodide (containing the ion), (d) magnesium phosphate (containing the ion).

  • 2.79

    Which of the following compounds are likely to be ionic? Which are likely to be molecular? SiCl4, LiF, BaCl2, B2H6, KCl, C2H4

    Answer

  • 2.80

    Which of the following compounds are likely to be ionic? Which are likely to be molecular? CH4, NaBr, BaF2, CCl4, ICl, CsCl, NF3

  • 2.81

    Name the following compounds: (a) KH2PO4, (b) K2HPO4, (c) HBr (gas), (d) HBr (in water), (e) Li2CO3, (f) K2Cr2O7, (g) NH4NO2, (h) HIO3, (i) PF5, (j) P4O6, (k) CdI2, (l) SrSO4, (m) Al(OH)3.

    Answer

  • 2.82

    Name the following compounds: (a) KClO, (b) Ag2CO3, (c) HNO2, (d) KMnO4, (e) CsClO3, (f) KNH4SO4, (g) FeO, (h) Fe2O3, (i) TiCl4, (j) NaH, (k) Li3N, (l) Na2O, (m) Na2O2.

  • 2.83

    Write the formulas for the following compounds: (a) rubidium nitrite, (b) potassium sulfide, (c) sodium hydrogen sulfide, (d) magnesium phosphate, (e) calcium hydrogen phosphate, (f) lead(II) carbonate, (g) tin(II) fluoride, (h) ammonium sulfate, (i) silver perchlorate, (j) boron trichloride.

    Answer

  • 2.84

    Write the formulas for the following compounds: (a) copper(I) cyanide, (b) strontium chlorite, (c) perbromic acid, (d) hydroiodic acid, (e) disodium ammonium phosphate, (f) potassium dihydrogen phosphate, (g) iodine heptafluoride, (h) tetraphosphorus decasulfide, (i) mercury(II) oxide, (j) mercury(I) iodide, (k) selenium hexafluoride.

  • 2.85

    In the diagrams shown here, match each of the drawings with the following ionic compounds: Al2O3, LiH, Na2S, Mg(NO3)2. (Green spheres represent cations and red spheres represent anions.)

    Answer

  • 2.86

    Given the formulas for the ionic compounds, draw the correct ratio of cations to anions as shown in Problem 2.85: (a) BaSO4, (b) CaF2, (c) Mg3N2, (d) K2O.

Additional Problems
  • 2.87

    Define the following terms: acids, bases, oxoacids, oxoanions, and hydrates.

    Answer

  • 2.88

    A sample of a uranium compound is found to be losing mass gradually. Explain what is happening to the sample.

  • 2.89

    In which one of the following pairs do the two species resemble each other most closely in chemical properties: (a) and , (b) and , (c) and Explain.

    Answer

  • 2.90

    One isotope of a metallic element has mass number 65 and 35 neutrons in the nucleus. The cation derived from the isotope has 28 electrons. Write the symbol for this cation.

  • 2.91

    One isotope of a nonmetallic element has mass number 127 and 74 neutrons in the nucleus. The anion derived from the isotope has 54 electrons. Write the symbol for this anion.

    Answer

  • 2.92

    The following table gives numbers of electrons, protons, and neutrons in atoms or ions of a number of elements. Answer the following: (a) Which of the species are neutral? (b) Which are negatively charged? (c) Which are positively charged? (d) What are the conventional symbols for all the species?

        Atom or Ion of Element  
      A B C D E F G
    Number of electrons 5 10 18 28 36 5 9
    Number of protons 5 7 19 30 35 5 9
    Number of neutrons 5 7 20 36 46 6 10
  • 2.93

    What is wrong with or ambiguous about the phrase “four molecules of NaCl”?

    Answer

  • 2.94

    The following phosphorus sulfides are known: P4S3, P4S7, and P4S10. Do these compounds obey the law of multiple proportions?

  • 2.95

    Which of the following are elements, which are molecules but not compounds, which are compounds but not molecules, and which are both compounds and molecules? (a) SO2, (b) S8, (c) Cs, (d) N2O5, (e) O, (f) O2, (g) O3, (h) CH4, (i) KBr, (j) S, (k) P4, (l) LiF.

    Answer

  • 2.96

    What is wrong with the name (given in parentheses or brackets) for each of the following compounds: (a) BaCl2 (barium dichloride), (b) Fe2O3 [iron(II) oxide], (c) CsNO2(cesium nitrate), (d) Mg(HCO3)2 [magnesium(II) bicarbonate]?

  • 2.97

    Discuss the significance of assigning an atomic mass of exactly 12 amu to the carbon-12 isotope.

    Answer

  • Page 73
  • 2.98

    Determine what is wrong with the chemical formula and write the correct chemical formula for each of the following compounds: (a) (NH3) 2CO3 (ammonium carbonate), (b) CaOH (calcium hydroxide), (c) CdSO3 (cadmium sulfide), (d) ZnCrO4 (zinc dichromate).

  • 2.99

    Fill in the blanks in the table:

    Symbol        
    Protons 5     79 86
    Neutrons 6   16 117 136
    Electrons 5   18 79  
    Net charge     +3   0

    Answer

  • 2.100

    (a) Which elements are most likely to form ionic compounds? (b) Which metallic elements are most likely to form cations with different charges?

  • 2.101

    Write the formula of the common ion derived from each of the following: (a) Li, (b) S, (c) I, (d) N, (e) Al, (f) Cs, (g) Mg.

    Answer

  • 2.102

    Which of the following symbols provides more information about the atom: 23Na or 11Na? Explain.

  • 2.103

    Write the chemical formulas and names of the binary acids and oxoacids that contain Group 7A elements. Do the same for elements in Groups 3A, 4A, 5A, and 6A.

    Answer

  • 2.104

    Determine the molecular and empirical formulas of the compounds shown here. (Black spheres are carbon, and white spheres are hydrogen.)

  • 2.105

    For the noble gases (the Group 8A elements) , , , , and (a) determine the number of protons and neutrons in the nucleus of each atom, and (b) determine the ratio of neutrons to protons in the nucleus of each atom. Describe any general trend you discover in the way this ratio changes with increasing atomic number.

    Answer

  • 2.106

    List the elements that exist as gases at room temperature. (Hint: Most of these elements can be found in Groups 5A, 6A, 7A, and 8A.)

  • 2.107

    The Group 1B metals, Cu, Ag, and Au, are called coinage metals. What chemical properties make them especially suitable for making coins and jewelry?

    Answer

  • 2.108

    The elements in Group 8A of the periodic table are called noble gases. Can you suggest what “noble” means in this context?

  • 2.109

    The formula for calcium oxide is CaO. What are the formulas for magnesium oxide and strontium oxide?

    Answer

  • 2.110

    A common mineral of barium is barytes, or barium sulfate (BaSO4). Because elements in the same periodic group have similar chemical properties, we might expect to find some radium sulfate (RaSO4) mixed with barytes since radium is the last member of Group 2A. However, the only source of radium compounds in nature is in uranium minerals. Why?

  • 2.111

    List five elements each that are (a) named after places, (b) named after people, (c) named after a color. (Consult www.Google.com, www.Wikipedia.com, or www.Webelements.com.)

    Answer

  • 2.112

    Name the only country that is named after an element. (Hint: This country is in South America.)

  • 2.113

    Fluorine reacts with hydrogen (H) and deuterium (D) to form hydrogen fluoride (HF) and deuterium fluoride (DF), where deuterium is an isotope of hydrogen. Would a given amount of fluorine react with different masses of the two hydrogen isotopes? Does this violate the law of definite proportion? Explain.

    Answer

  • 2.114

    Predict the formula and name of a binary compound formed from the following elements: (a) Na and H, (b) B and O, (c) Na and S, (d) Al and F, (e) F and O, (f) Sr and Cl.

  • 2.115

    Identify each of the following elements: (a) a halogen whose anion contains 36 electrons, (b) a radioactive noble gas with 86 protons, (c) a Group 6A element whose anion contains 36 electrons, (d) an alkali metal cation that contains 36 electrons, (e) a Group 4A cation that contains 80 electrons.

    Answer

  • 2.116

    Show the locations of (a) alkali metals, (b) alkaline earth metals, (c) the halogens, and (d) the noble gases in the given outline of a periodic table. Also draw dividing lines between metals and metalloids and between metalloids and nonmetals.

  • 2.117

    Fill in the blanks in the table.

    Answer

    Cation Anion Formula Name
          Magnesium bicarbonate
        SrCl2  
    Fe3+    
          Manganese(II) chlorate
        SnBr4  
    Co2+    
    I    
        Cu2CO3  
          Lithium nitride
    Al3+ S2–    
  • 2.118

    Some compounds are better known by their common names than by their systematic chemical names. Give the chemical formulas of the following substances: (a) dry ice, (b) salt, (c) laughing gas, (d) marble (chalk, limestone), (e) baking soda, (f) ammonia, (g) water, (h) milk of magnesia, (i) epsom salt.

  • 2.119

    In the footnote on page 37 it was pointed out that mass and energy are alternate aspects of a single entity called mass-energy. The relationship between these two physical quantities is Einstein's equation, E = mc2, where E is energy, m is mass, and c is the speed of light. In a combustion experiment, it was found that 12.096 g of hydrogen molecules combined with 96.000 g of oxygen molecules to form water and released 1.715 × 103 kJ of heat. Use Einstein's equation to calculate the corresponding mass change in this process, and comment on whether or not the law of conservation of mass holds for ordinary chemical processes.

    Answer

  • Page 74
  • 2.120

    (a) Describe Rutherford's experiment and how the results revealed the nuclear structure of the atom. (b) Consider the 23Na atom. Given that the radius and mass of the nucleus are 3.04 × 10−15 m and 3.82 × 10−23 g, respectively, calculate the density of the nucleus in g/cm3. The radius of a 23Na atom is 186 pm. Calculate the density of the space occupied by the electrons outside the nucleus in the sodium atom. Do your results support Rutherford's model of an atom? [The volume of a sphere of radius r is .]

  • 2.121

    Draw all possible structural formulas of the following hydrocarbons: CH4, C2H6, C3H8, C4H10, C5H12.

    Answer

  • 2.122

    Draw two different structural formulas based on the molecular formula C2H6O. Is the fact that you can have more than one compound with the same molecular formula consistent with Dalton's atomic theory?

  • 2.123

    Ethane and acetylene are two gaseous hydrocarbons. Chemical analyses show that in one sample of ethane, 2.65 g of carbon are combined with 0.665 g of hydrogen, and in one sample of acetylene, 4.56 g of carbon are combined with 0.383 g of hydrogen. (a) Are these results consistent with the law of multiple proportions? (b) Write reasonable molecular formulas for these compounds.

    Answer

  • 2.124

    A cube made of platinum (Pt) has an edge length of 1.0 cm. (a) Calculate the number of Pt atoms in the cube. (b) Atoms are spherical in shape. Therefore, the Pt atoms in the cube cannot fill all the available space. If only 74 percent of the space inside the cube is taken up by Pt atoms, calculate the radius in picometers of a Pt atom. The density Pt is 21.45 g/cm3, and the mass of a single Pt atom is 3.240 × 10−22 g. [The volume of a sphere of radius r is .]

  • 2.125

    A monatomic ion has a charge of +2. The nucleus of the parent atom has a mass number of 55. If the number of neutrons in the nucleus is 1.2 times that of the number of protons, what is the name and symbol of the element?

    Answer

  • 2.126

    In the following 2 × 2 crossword, each letter must be correct in four ways: horizontally, vertically, diagonally, and by itself. When the puzzle is complete, the four spaces will contain the overlapping symbols of 10 elements. Use capital letters for each square. There is only one correct solution.

    Horizontal

    1–2: Two-letter symbol for a metal used in ancient times
    3–4: Two-letter symbol for a metal that burns in air and is found in Group 5A

    Vertical

    1–3: Two-letter symbol for a metalloid
    2–4: Two-letter symbol for a metal used in U.S. coins

    Single Square

    1: A colorful nonmetal
    2: A colorless gaseous nonmetal
    3: An element that makes fireworks green
    4: An element that has medicinal uses

    Diagonal

    1–4: Two-letter symbol for an element used in electronics
    2–3: Two-letter symbol for a metal used with Zr to make wires for superconducting magnets
  • 2.127

    Name the given acids.

    Answer

  • 2.128

    (a) Assuming an atomic nucleus is spherical in shape, show that its radius r is proportional to the cube root of the mass number (A). (b) In general, the radius of a nucleus is given by r = r0A1/3, where r0 is a proportionality constant given by 1.2 × 10−15 m. Calculate the volume of the Li nucleus. (c) Given that the radius of a Li atom is 152 pm, calculate what fraction of the atom's volume is occupied by its nucleus. Does your result support Rutherford's model of the atom?

Page 75
Pre-Professional Practice Exam Problems: Physical and Biological Sciences

Carbon-14, a radioactive isotope of carbon, is used to determine the ages of fossils in a technique called carbon dating. Carbon-14 is produced in the upper atmosphere when nitrogen-14 atoms are bombarded by neutrons from cosmic rays. 14C undergoes a process called β emission in which a neutron in the nucleus decays to form a proton and an electron. The electron, or β particle, is ejected from the nucleus. Because the production and decay of 14C occur simultaneously, the total amount of 14C in the atmosphere is constant. Plants absorb 14C in the form of CO2 and animals consume plants and other animals. Thus, all living things contain a constant ratio of 12C to 14C. When a living thing dies, the 14C it contains continues to decay but because replenishment ceases, the ratio of 12C to 14C changes over time. Scientists use the 12C to 14C ratio to determine the age of material that was once living.

  1. If atmospheric conditions were to change such that 14C were produced at twice the current rate,

    1. the world's supply of 14N would be consumed completely.

    2. the 12C to 14C ratio in living things would increase.

    3. the 12C to 14C ratio in living things would decrease.

    4. the 12C to 14C ratio in living things would not change.

  2. When a 14N nucleus is bombarded by a neutron to produce a 14C nucleus, what else is produced?

    1. Nothing

    2. Another neutron

    3. An electron

    4. A proton

  3. Based on the description of β emission in the passage, what nucleus results from the decay of a 14C nucleus by β emission?

    1. 14N

    2. 13N

    3. 12C

    4. 13C

  4. The accuracy of carbon dating depends on the assumption that

    1. 14C is the only radioactive species in the material being tested.

    2. the rate of decay of 14C is constant.

    3. 12C and 14C undergo radioactive decay at the same rate.

    4. each 14C nucleus decays to give a 12C nucleus.

Answers to In-Chapter Materials
Practice Problems
  • 2.1A
    1. 3:2,

    2. 2:1.

  • 2.1B
    1. 0.2518 g,

    2. n = 6 (XeF6).

  • 2.2A
    1. p = 5, n = 5, e = 5.

    2. p = 18, n = 18, e = 18.

    3. p = 38, n = 47, e = 38.

    4. p = 6, n = 5, e = 6.

  • 2.2B
  • 2.3A

    63.55 amu.

  • 2.3B

    99.64% 14N, 0.36% 15N.

  • 2.4A

    CHCl3.

  • 2.4B

    C3H6O.

  • 2.5A
    1. dichlorine monoxide,

    2. silicon tetrachloride.

  • 2.5B
    1. chlorine dioxide,

    2. carbon tetrabromide.

  • 2.6A
    1. CS2,

    2. N2O3.

  • 2.6B
    1. SF6,

    2. S2F10.

  • 2.7A
    1. C4H5N2O,

    2. C2H5,

    3. C2H5NO2.

  • 2.7B
  • 2.8A
    1. hypobromous acid,

    2. hydrogen sulfate ion,

    3. oxalic acid.

  • 2.8B
    1. iodic acid,

    2. hydrogen chromate ion,

    3. hydrogen oxalate ion.

  • 2.9A

    HBrO4.

  • 2.9B

    H2CrO4.

  • 2.10A
    1. sodium sulfate,

    2. copper(II) nitrate,

    3. iron(III) carbonate.

  • 2.10B
    1. potassium dichromate,

    2. lithium oxalate,

    3. copper(I) nitrate.

  • 2.11A
    1. PbCl2,

    2. MgCO3,

    3. (NH4)3PO4.

  • 2.11B
    1. Fe2S3,

    2. Hg(NO3)2,

    3. K2SO3.

Checkpoints
  • 2.1.1

    e.

  • 2.1.2

    b.

  • 2.3.1

    d.

  • 2.3.2

    c.

  • 2.4.1

    c.

  • 2.4.2

    b.

  • 2.5.1

    b.

  • 2.5.2

    c.

  • 2.6.1

    b.

  • 2.6.2

    c.

  • 2.6.3

    c.

  • 2.6.4

    e.

  • 2.7.1

    c.

  • 2.7.2

    d.

  • 2.7.3

    d.

  • 2.7.4

    c.

  • 2.7.5

    c.

  • 2.7.6

    a.

Applying What You've Learned
  1. There are 54 – 26 = 28 neutrons in the 54Fe nucleus, 30 neutrons in the 56Fe nucleus, 31 neutrons in the 57Fe nucleus, and 32 neutrons in the 58Fe nucleus.

  2. The average atomic mass of iron is 55.845 amu.

  3. The molecular formula for ascorbic acid is C6H8O6.

  4. The empirical formula for ascorbic acid is C3H4O3.

  5. The formula for ferrous sulfate is FeSO4.