7.f^Chapter 7 Ending^289^295^,,^17226^18143%
Page 289
Chapter Summary
Section 7.1
  • The modern periodic table was devised independently by Dmitri Mendeleev and Lothar Meyer in the nineteenth century. The elements that were known at the time were grouped based on their physical and chemical properties. Using his arrangements of the elements, Mendeleev successfully predicted the existence of elements that had not yet been discovered.

  • Early in the twentieth century, Henry Moseley refined the periodic table with the concept of the atomic number, thus resolving a few inconsistencies in the tables proposed by Mendeleev and Meyer.

  • Elements in the same group of the periodic table tend to have similar physical and chemical properties.

Section 7.2
  • The periodic table can be divided into the main group elements (also known as the representative elements) and the transition metals. It is further divided into smaller groups or columns of elements that all have the same configuration of valence electrons.

  • The 18 columns of the periodic table are labeled 1A through 8A (s and p-block elements) and 1B through 8B (d-block elements), or by the numbers 1 through 18.

Section 7.3
  • Effective nuclear charge (Zeff) is the nuclear charge that is “felt” by the valence electrons. It is usually lower than the nuclear charge due to shielding by the core electrons.

  • According to Coulomb's law, the attractive force (F) between two oppositely charged particles (Q1 and Q2) is directly proportional to the product of the charges and inversely proportional to the distance (d) between the objects squared: (F α Q1 · Q2/d2).

Section 7.4
  • Atomic radius is the distance between an atom's nucleus and its valence shell. The atomic radius of a metal atom is defined as the metallic radius, which is one-half the distance between adjacent, identical nuclei in a metal solid. The atomic radius of a nonmetal is defined as the covalent radius, which is one-half the distance between adjacent, identical nuclei in a molecule. In general, atomic radii decrease from left to right across a period of the periodic table and increase from top to bottom down a group.

  • Ionization energy (IE) is the energy required to remove an electron from a gaseous atom. The first ionization energy (IE1) is smaller than subsequent ionization energies [e.g., second (IE2), third (IE3), and so on]. The first ionization of any atom removes a valence electron. Ionization energies increase dramatically when core electrons are being removed.

  • First ionization energies (IE1 values) tend to increase across a period and decrease down a group. Exceptions to this trend can be explained based upon the electron configuration of the element.

  • Electron affinity (EA) is the energy released when an atom in the gas phase accepts an electron. EA is equal to −ΔH for the process .

  • Electron affinities tend to increase across a period. As with first ionization energies, exceptions to the trend can be explained based on the electron configuration of the element.

  • Metals tend to be shiny, lustrous, malleable, ductile, and conducting (for both heat and electricity). Metals typically lose electrons to form cations, and they tend to form ionic compounds (including basic oxides).

  • Nonmetals tend to be brittle and not good conductors (for either heat or electricity). They can gain electrons to form anions but they commonly form molecular compounds (including acidic oxides).

  • In general, metallic character decreases across a period and increases down a group of the periodic table. Metalloids are elements with properties intermediate between metals and nonmetals.

Section 7.5
  • Ions of main group elements are isoelectronic with noble gases. When a d-block element loses one or more electrons, it loses them first from the shell with the highest principal quantum number (e.g., electrons in the 4s subshell are lost before electrons in the 3d subshell).

Section 7.6
  • Ionic radius is the distance between the nucleus and valence shell of a cation or an anion. A cation is smaller than its parent atom. An anion is larger than its parent atom.

  • An isoelectronic series consists of one or more ions and a noble gas, all of which have identical electron configurations. Within an isoelectronic series, the greater the nuclear charge, the smaller the radius.

Section 7.7
  • A diagonal relationship describes similarities in the chemical properties of elements that are in different groups, but that are positioned diagonally from each other in the periodic table.

  • Although members of a group in the periodic table exhibit similar chemical and physical properties, the first member of each group tends to be significantly different from the other members. Hydrogen is essentially a group unto itself.

  • The alkali metals (Group 1A) tend to be highly reactive toward oxygen, water, and acid. Group 2A metals are less reactive than Group 1A metals, but the heavier members all react with water to produce metal hydroxides and hydrogen gas. Groups that contain both metals and nonmetals (e.g., Groups 4A, 5A, and 6A) tend to show greater variability in their physical and chemical properties.

  • Amphoteric oxides, such as Al2O3, are those that exhibit both acidic and basic behavior.

Key Words
Page 290
Key Equations
Questions and Problems
Section 7.1: Development of the Periodic Table
Review Questions
  • 7.1

    Briefly describe the significance of Mendeleev's periodic table.

  • 7.2

    What is Moseley's contribution to the modern periodic table?

  • 7.3

    Describe the general layout of a modern periodic table.

  • 7.4

    What is the most important relationship among elements in the same group in the periodic table?

Section 7.2: The Modern Periodic Table
Review Questions
  • 7.5

    Classify each of the following elements as a metal, a nonmetal, or a metalloid: As, Xe, Fe, Li, B, Cl, Ba, P, I, Si.

  • 7.6

    Compare the physical and chemical properties of metals and nonmetals.

  • 7.7

    Draw a rough sketch of a periodic table (no details are required). Indicate regions where metals, nonmetals, and metalloids are located.

  • 7.8

    What is a main group element? Give names and symbols of four main group elements.

  • 7.9

    Without referring to a periodic table, write the name and give the symbol for one element in each of the following groups: 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A, transition metals.

  • 7.10

    Indicate whether the following elements exist as atomic species, molecular species, or extensive three-dimensional structures in their most stable states at room temperature, and write the molecular or empirical formula for each one: phosphorus, iodine, magnesium, neon, carbon, sulfur, cesium, and oxygen.

  • 7.11

    You are given a sample of a dark, shiny solid and asked to determine whether it is the nonmetal iodine or a metallic element. What test could you do that would enable you to answer the question without destroying the sample?

  • 7.12

    What are valence electrons? For main group elements, the number of valence electrons of an element is equal to its group number. Show that this is true for the following elements: Al, Sr, K, Br, P, S, C.

  • 7.13

    Write the outer electron configurations for the (a) alkali metals, (b) alkaline earth metals, (c) halogens, (d) noble gases.

  • 7.14

    Use the first-row transition metals (Sc to Cu) as an example to illustrate the characteristics of the electron configurations of transition metals.

  • 7.15

    Arsenic is not an essential element for the human body. Based on its position in the periodic table, suggest a reason for its toxicity.

  • 7.16

    In the periodic table, the element hydrogen is sometimes grouped with the alkali metals and sometimes with the halogens. Explain why hydrogen can resemble the Group 1A and the Group 7A elements.

  • 7.17

    A neutral atom of a certain element has 16 electrons. Consulting only the periodic table, identify the element and write its ground-state electron configuration.


  • 7.18

    Group the following electron configurations in pairs that would represent elements with similar chemical properties:

    1. 1s22s22p63s2

    2. 1s22s22p3

    3. 1s22s22p63s23p64s23d104p6

    4. 1s22s2

    5. 1s22s22p6

    6. 1s22s22p63s23p3

  • 7.19

    Group the following electron configurations in pairs that would represent elements with similar chemical properties:

    1. 1s22s22p5

    2. 1s22s1

    3. 1s22s22p6

    4. 1s22s22p63s23p5

    5. 1s22s22p63s23p64s1

    6. 1s22s22p63s23p64s23d104p6


  • 7.20

    Without referring to a periodic table, write the electron configuration of elements with the following atomic numbers: (a) 9, (b) 20, (c) 26, (d) 33.

  • 7.21

    Specify the group of the periodic table in which each of the following elements is found: (a) [Ne]3s1, (b) [Ne]3s23p3, (c) [Ne]3s23p6, (d) [Ar]4s23d8.


Section 7.3: Effective Nuclear Charge
Review Questions
  • 7.22

    Explain the term effective nuclear charge.

  • 7.23

    Explain why the atomic radius of Be is smaller than that of Li.

  • 7.24

    The electron configuration of B is 1s22s22p1. (a) If each core electron (that is, the 1s electrons) were totally effective in shielding the valence electrons (that is, the 2s and 2p electrons) from the nucleus and the valence electrons did not shield one another, what would be the shielding constant (σ) and the effective nuclear charge (Zeff) for the 2s and 2p electrons? (b) In reality, the shielding constants for the 2s and 2p electrons in B are slightly different. They are 2.42 and 2.58, respectively. Calculate Zeff for these electrons, and explain the differences from the values you determined in part (a).

  • Page 291
  • 7.25

    The electron configuration of C is 1s22s22p2. (a) If each core electron (that is, the 1s electrons) were totally effective in shielding the valence electrons (that is, the 2s and 2p electrons) from the nucleus and the valence electrons did not shield one another, what would be the shielding constant (σ) and the effective nuclear charge (Zeff) for the 2s and 2p electrons? (b) In reality, the shielding constants for the 2s and 2p electrons in C are slightly different. They are 2.78 and 2.86, respectively. Calculate Zeff for these electrons, and explain the differences from the values you determined in part (a).


Section 7.4: Periodic Trends in Properties of Elements
Review Questions
  • 7.26

    Define atomic radius. Does the size of an atom have a precise meaning?

  • 7.27

    How does atomic radius change (a) from left to right across a period and (b) from top to bottom in a group?

  • 7.28

    Define ionization energy. Explain why ionization energy measurements are usually made when atoms are in the gaseous state. Why is the second ionization energy always greater than the first ionization energy for any element?

  • 7.29

    Sketch the outline of the periodic table, and show group and period trends in the first ionization energy of the elements. What types of elements have the highest ionization energies and what types have the lowest ionization energies?

  • 7.30

    (a) Define electron affinity. (b) Explain why electron affinity measurements are made with gaseous atoms. (c) Ionization energy is always a positive quantity, whereas electron affinity may be either positive or negative. Explain.

  • 7.31

    Explain the trends in electron affinity from aluminum to chlorine (see Figure 7.10).

  • 7.32

    On the basis of their positions in the periodic table, select the atom with the larger atomic radius in each of the following pairs:

    1. Na, Si;

    2. Ba, Be;

    3. N, F;

    4. Br, Cl;

    5. Ne, Kr.

  • 7.33

    Arrange the following atoms in order of increasing atomic radius: Na, Al, P, Cl, Mg.


  • 7.34

    Which is the largest atom in the third period of the periodic table?

  • 7.35

    Which is the smallest atom in Group 7A?


  • 7.36

    Based on size, identify the spheres shown as Na, Mg, O, and S.

  • 7.37

    Based on size, identify the spheres shown as K, Ca, S, and Se.


  • 7.38

    Why is the radius of the lithium atom considerably larger than the radius of the hydrogen atom?

  • 7.39

    Use the second period of the periodic table as an example to show that the size of atoms decreases as we move from left to right. Explain the trend.


  • 7.40

    Arrange the following in order of increasing first ionization energy: Na, Cl, Al, S, and Cs.

  • 7.41

    Arrange the following in order of increasing first ionization energy: F, K, P, Ca, and Ne.


  • 7.42

    Use the third period of the periodic table as an example to illustrate the change in first ionization energies of the elements as we move from left to right. Explain the trend.

  • 7.43

    In general, the first ionization energy increases from left to right across a given period. Aluminum, however, has a lower first ionization energy than magnesium. Explain.


  • 7.44

    The first and second ionization energies of K are 419 and 3052 kJ/mol, and those of Ca are 590 and 1145 kJ/mol, respectively. Compare their values and comment on the differences.

  • 7.45

    Two atoms have the electron configurations 1s22s22p6 and 1s22s22p63s1. The first ionization energy of one is 2080 kJ/mol, and that of the other is 496 kJ/mol. Match each ionization energy with one of the given electron configurations. Justify your choice.


  • 7.46

    A hydrogen-like ion is an ion containing only one electron. The energies of the electron in a hydrogen-like ion are given by

    where n is the principal quantum number and Z is the atomic number of the element. Calculate the ionization energy (in kJ/mol) of the He+ ion.

  • 7.47

    Plasma is a state of matter consisting of positive gaseous ions and electrons. In the plasma state, a mercury atom could be stripped of its 80 electrons and therefore would exist as Hg80+. Use the equation in Problem 7.46 to calculate the energy required for the last ionization step, that is,


  • 7.48

    Arrange the elements in each of the following groups in order of increasing electron affinity:

    1. Li, Na, K;

    2. F, Cl, Br, I.

  • 7.49

    Specify which of the following elements you would expect to have the greatest electron affinity: He, K, Co, S, Cl.


  • 7.50

    Considering their electron affinities, do you think it is possible for the alkali metals to form an anion like M, where M represents an alkali metal?

  • 7.51

    Explain why alkali metals have a greater affinity for electrons than alkaline earth metals.


Section 7.5: Electron Configuration of Ions
Review Questions
  • 7.52

    How does the electron configuration of ions derived from main group elements give them stability?

  • 7.53

    What do we mean when we say that two ions or an atom and an ion are isoelectronic?

  • 7.54

    Is it possible for the atoms of one element to be isoelectronic with the atoms of another element? Explain.

  • 7.55

    Give three examples of first-row transition metal (Se to Cu) ions that are isoelectronic with argon.

  • 7.56

    A M2+ ion derived from a metal in the first transition metal series has four electrons in the 3d subshell. What element might M be?

  • 7.57

    A metal ion with a net +3 charge has five electrons in the 3d subshell. Identify the metal.


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  • 7.58

    Write the ground-state electron configurations of the following ions:

    1. Li+,

    2. H,

    3. N3−,

    4. F,

    5. S2−,

    6. Al3+,

    7. Se2−,

    8. Br,

    9. Rb+,

    10. Sr2+,

    11. Sn2+,

    12. Te2−,

    13. Ba2+,

    14. Pb2+,

    15. In3+,

    16. Tl+,

    17. Tl3+.

  • 7.59

    Write the ground-state electron configurations of the following ions, which play important roles in biochemical processes in our bodies:

    1. Na+,


    2. Mg2+,

    3. Cl,

    4. K+,

    5. Ca2+,

    6. Fe2+,

    7. Cu2+,

    8. Zn2+.

  • 7.60

    Write the ground-state electron configurations of the following transition metal ions:

    1. Sc3+,

    2. Ti4+,

    3. V5+,

    4. Cr3+,

    5. Mn2+,

    6. Fe2+,

    7. Fe3+,

    8. Co2+,

    9. Ni2+,

    10. Cu+,

    11. Cu2+,

    12. Ag+,

    13. Au+,

    14. Au3+,

    15. Pt2+.

  • 7.61

    Name the ions with three charges that have the following electron configurations:

    1. [Ar]3d3,

    2. [Ar],

    3. [Kr]4d6,

    4. [Xe]4f145d6.


  • 7.62

    Which of the following species are isoelectronic with each other: C, Cl, Mn2+, B, Ar, Zn, Fe3+, Ge2+?

  • 7.63

    Group the species that are isoelectronic: Be2+, F, Fe2+, N3−, He, S2−, Co3+, Ar.


  • 7.64

    Thallium (Tl) is a neurotoxin and exists mostly in the Tl(I) oxidation state in its compounds. Aluminum (Al), which causes anemia and dementia, is only stable in the Al(III) form. The first, second, and third ionization energies of Tl are 589, 1971, and 2878 kJ/mol, respectively. The first, second, and third ionization energies of Al are 577.5, 1817, and 2745 kJ/mol, respectively. Plot the ionization energies of Al and Tl versus atomic number and explain the trends.

Section 7.6: Ionic Radius
Review Questions
  • 7.65

    Define ionic radius. How does the size of an atom change when it is converted to (a) an anion and (b) a cation?

  • 7.66

    Explain why, for isoelectronic ions, the anions are larger than the cations.

  • 7.67

    Indicate which one of the two species in each of the following pairs is smaller:

    1. Cl or Cl,


    2. Na or Na+,

    3. O2− or S2−,

    4. Mg2+ or Al3+,

    5. Au+ or Au3+.

  • 7.68

    List the following ions in order of increasing ionic radius: N3−, Na+, F, Mg2+, O2−.

  • 7.69

    Explain which of the following cations is larger, and why: Cu+ or Cu2+.


  • 7.70

    Explain which of the following anions is larger, and why: Se2− or Te2−.

  • 7.71

    Both Mg2+ and Ca2+ are important biological ions. One of their functions is to bind to the phosphate group of ATP molecules or amino acids of proteins. For Group 2A metals in general, the tendency for binding to the anions increases in the order Ba2+ < Sr2+ < Ca2+ < Mg2+. Explain this trend.


Section 7.7: Periodic Trends in Chemical Properties of the Main Group Elements
Review Questions
  • 7.72

    Why do members of a group exhibit similar chemical properties?

  • 7.73

    Which elements are more likely to form acidic oxides? Basic oxides? Amphoteric oxides?

  • 7.74

    Give the physical states (gas, liquid, or solid) of the main group elements in the fourth period (K, Ca, Ga, Ge, As, Se, Br) at room temperature.

  • 7.75

    The boiling points of neon and krypton are − 246.1°C and −153.2°C, respectively. Using these data, estimate the boiling point of argon.


  • 7.76

    Use the alkali metals and alkaline earth metals as examples to show how we can predict the chemical properties of elements simply from their electron configurations.

  • 7.77

    Based on your knowledge of the chemistry of the alkali metals, predict some of the chemical properties of francium, the last member of the group.


  • 7.78

    As a group, the noble gases are very stable chemically (only Kr and Xe are known to form compounds). Why?

  • 7.79

    Why are Group 1B elements more stable than Group 1A elements even though they seem to have the same outer electron configuration, ns1, where n is the principal quantum number of the outermost shell?


  • 7.80

    How do the chemical properties of oxides change from left to right across a period? How do they change from top to bottom within a particular group?

  • 7.81

    Write balanced equations for the reactions between each of the following oxides and water:

    1. Li2O,


    2. CaO,

    3. SO3.

  • 7.82

    Write formulas for and name the binary hydrogen compounds of the second-period elements (Li to F). Describe how the physical and chemical properties of these compounds change from left to right across the period.

  • 7.83

    Which oxide is more basic, MgO or BaO? Why?


Additional Problems
  • 7.84

    State whether each of the following properties of the main group elements generally increases or decreases (a) from left to right across a period and (b) from top to bottom within a group: metallic character, atomic size, ionization energy, acidity of oxides.

  • 7.85

    Referring to the periodic table, name (a) the halogen in the fourth period, (b) an element similar to phosphorus in chemical properties, (c) the most reactive metal in the fifth period, (d) an element that has an atomic number smaller than 20 and is similar to strontium.


  • 7.86

    Write equations representing the following processes:

    1. The electron affinity of S

    2. The third ionization energy of titanium

    3. The electron affinity of Mg2+

    4. The ionization energy of O2−

  • 7.87

    Arrange the following isoelectronic species in order of increasing ionization energy: O2−, F, Na+, Mg2+.


  • 7.88

    Write the empirical (or molecular) formulas of compounds that the elements in the third period (sodium to chlorine) should form with (a) molecular oxygen and (b) molecular chlorine. In each case indicate whether you would expect the compound to be ionic or molecular in character.

  • 7.89

    Element M is a shiny and highly reactive metal (melting point 63°C), and element X is a highly reactive nonmetal (melting point 27.2°C). They react to form a compound with the empirical formula MX, a colorless, brittle white solid that melts at 734°C. When dissolved in water or when in the molten state, the substance conducts electricity. When chlorine gas is bubbled through an aqueous solution containing MX, a reddish-brown liquid appears and Cl ions are formed. From these observations, identify M and X. (You may need to consult a handbook of chemistry for the melting-point values.)


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  • 7.90

    Match each of the elements on the right with its description on the left:

    (a) A dark-red liquid Calcium (Ca)
    (b) A colorless gas that burns in oxygen gas Gold (Au)
    (c) A metal that reacts violently with water Hydrogen (H2)
    (d) A shiny metal that is used in jewelry Argon (Ar)
    (e) An inert gas Bromine (Br2)
  • 7.91

    Arrange the following species in isoelectronic pairs: O+, Ar, S2−, Ne, Zn, Cs+, N3−, As3+, N, Xe.


  • 7.92

    In which of the following are the species written in decreasing order by size of radius:

    1. Be, Mg, Ba,

    2. N3−, O2−, F,

    3. Tl3+, Tl2+, Tl+?

  • 7.93

    Which of the following properties show a clear periodic variation:

    1. first ionization energy,

    2. molar mass of the elements,

    3. number of isotopes of an element,

    4. atomic radius?


  • 7.94

    When carbon dioxide is bubbled through a clear calcium hydroxide solution, the solution appears milky. Write an equation for the reaction, and explain how this reaction illustrates that CO2 is an acidic oxide.

  • 7.95

    You are given four substances: a fuming red liquid, a dark metallic-looking solid, a pale-yellow gas, and a yellow-green gas that attacks glass. You are told that these substances are the first four members of Group 7A, the halogens. Name each one.


  • 7.96

    For each pair of elements listed, give three properties that show their chemical similarity:

    1. sodium and potassium and

    2. chlorine and bromine.

  • 7.97

    Name the element that forms compounds, under appropriate conditions, with every other element in the periodic table except He, Ne, and Ar.


  • 7.98

    Explain why the first electron affinity of sulfur is 200 kJ/mol but the second electron affinity is −649 kJ/mol.

  • 7.99

    The H ion and the He atom have two 1s electrons each. Which of the two species is larger? Explain.


  • 7.100

    Predict the products of the following oxides with water: Na2O, BaO, CO2, N2O5, P4O10, SO3. Write an equation for each of the reactions. Specify whether the oxides are acidic, basic, or amphoteric.

  • 7.101

    Write the formulas and names of the oxides of the second-period elements (Li to N). Identify the oxides as acidic, basic, or amphoteric. Use the highest oxidation state of each element.


  • 7.102

    State whether each of the following elements is a gas, liquid, or solid under atmospheric conditions. Also state whether it exists in the elemental form as atoms, molecules, or a three-dimensional network: Mg, Cl, Si, Kr, O, I, Hg, Br.

  • 7.103

    What factors account for the unique nature of hydrogen?


  • 7.104

    The air in a manned spacecraft or submarine needs to be purified of exhaled carbon dioxide. Write equations for the reactions between carbon dioxide and (a) lithium oxide (Li2O), (b) sodium peroxide (Na2O2), and (c) potassium superoxide (KO2).

  • 7.105

    The formula for calculating the energies of an electron in a hydrogen-like ion is given in Problem 7.46. This equation can be applied only to one-electron atoms or ions. One way to modify it for more complex species is to replace Z with Z − σ or Zeff. Calculate the value of σ if the first ionization energy of helium is 3.94 × 10−18 J per atom. (Disregard the minus sign in the given equation in your calculation.)


  • 7.106

    Why do noble gases have negative electron affinity values?

  • 7.107

    The atomic radius of K is 227 pm and that of K1 is 138 pm. Calculate the percent decrease in volume that occurs when K(g) is converted to K+(g). (The volume of a sphere is , where r is the radius of the sphere.)


  • 7.108

    The atomic radius of F is 72 pm and that of F is 133 pm. Calculate the percent increase in volume that occurs when F(g) is converted to F(g). (See Problem 7.107 for the volume of a sphere.)

  • 7.109

    A technique called photoelectron spectroscopy is used to measure the ionization energy of atoms. A gaseous sample is irradiated with UV light, and electrons are ejected from the valence shell. The kinetic energies of the ejected electrons are measured. Because the energy of the UV photon and the kinetic energy of the ejected electron are known, we can write

    where ν is the frequency of the UV light, and m and u are the mass and velocity of the electron, respectively. In one experiment the kinetic energy of the ejected electron from potassium is found to be 5.34 × 10−19J using a UV source of wavelength 162 nm. Calculate the ionization energy of potassium. How can you be sure that this ionization energy corresponds to the electron in the valence shell (that is, the most loosely held electron)?


  • 7.110

    The energy needed for the following process is 1.96 × 104kJ/mol:

    If the first ionization energy of lithium is 520 kJ/mol, calculate the second ionization energy of lithium, that is, the energy required for the process

    (Hint: You need the equation in Problem 7.46.)

  • 7.111

    A student is given samples of three elements, X, Y, and Z, which could be an alkali metal, a member of Group 4A, or a member of Group 5A. She makes the following observations: Element X has a metallic luster and conducts electricity. It reacts slowly with hydrochloric acid to produce hydrogen gas. Element Y is a light yellow solid that does not conduct electricity. Element Z has a metallic luster and conducts electricity. When exposed to air, it slowly forms a white powder. A solution of the white powder in water is basic. What can you conclude about the elements from these observations?


  • 7.112

    What is the electron affinity of the Na+ ion?

  • 7.113

    The ionization energies of sodium (in kJ/mol), starting with the first and ending with the eleventh, are 496, 4562, 6910, 9543, 13,354, 16,613, 20,117, 25,496, 28,932, 141,362, 159,075. Plot the log of ionization energy (y axis) versus the number of ionization (x axis); for example, log 496 is plotted versus 1 (labeled IE1, the first ionization energy), log 4562 is plotted versus 2 (labeled IE2, the second ionization energy), and so on. (a) Label IE1 through IE11 with the electrons in orbitals such as 1s, 2s, 2p, and 3s. (b) What can you deduce about electron shells from the breaks in the curve?


  • Page 294
  • 7.114

    Experimentally, the electron affinity of an element can be determined by using a laser light to ionize the anion of the element in the gas phase:

    Referring to Figure 7.10, calculate the photon wavelength (in nm) corresponding to the electron affinity for chlorine. In what region of the electromagnetic spectrum does this wavelength fall?

  • 7.115

    Explain, in terms of their electron configurations, why Fe2+ is more easily oxidized to Fe3+ than Mn2+ is to Mn3+.


  • 7.116

    Write the formulas and names of the hydrides of the following second-period elements: Li, C, N, O, F. Predict their reactions with water.

  • 7.117

    Based on knowledge of the electronic configuration of titanium, state which of the following compounds of titanium is unlikely to exist: K3TiF6, K2Ti2O5, TiCl3, K2TiO4, K2TiF6.


  • 7.118

    In halogen displacement reactions a halogen element can be generated by oxidizing its anions with a halogen element that lies above it in the periodic table. This means that there is no way to prepare elemental fluorine, because it is the first member of Group 7A. Indeed, for years the only way to prepare elemental fluorine was to oxidize F ions by electrolytic means. Then, in 1986, a chemist reported that by combining potassium hexafluoromanganate(IV) (K2MnF6) with antimony pentafluoride (SbF5) at 150°C, he had generated elemental fluorine. Balance the following equation representing the reaction:

  • 7.119

    Write a balanced equation for the preparation of (a) molecular oxygen, (b) ammonia, (c) carbon dioxide, (d) molecular hydrogen, (e) calcium oxide. Indicate the physical state of the reactants and products in each equation.


  • 7.120

    Write chemical formulas for oxides of nitrogen with the following oxidation numbers: +1, +2, +3, +4, +5. (Hint: There are two oxides of nitrogen with a +4 oxidation number.)

  • 7.121

    Most transition metal ions are colored. For example, a solution of CuSO4 is blue. How would you show that the blue color is due to the hydrated Cu2+ ions and not the ions?


  • 7.122

    In general, atomic radius and ionization energy have opposite periodic trends. Why?

  • 7.123

    Explain why the electron affinity of nitrogen is approximately zero, while the elements on either side, carbon and oxygen, have substantial positive electron affinities.


  • 7.124

    Consider the halogens chlorine, bromine, and iodine. The melting point and boiling point of chlorine are −101.5°C and −34.0°C and those of iodine are 113.7°C and 184.3°C, respectively. Thus chlorine is a gas and iodine is a solid under room conditions. Estimate the melting point and boiling point of bromine. Compare your values with those from the webelements.com website.

  • 7.125

    Although it is possible to determine the second, third, and higher ionization energies of an element, the same cannot usually be done with the electron affinities of an element. Explain.


  • 7.126

    Little is known of the chemistry of astatine, the last member of Group 7A. Describe the physical characteristics that you would expect this halogen to have. Predict the products of the reaction between sodium astatide (NaAt) and sulfuric acid. (Hint: Sulfuric acid is an oxidizing agent.)

  • 7.127

    As discussed in the chapter, the atomic mass of argon is greater than that of potassium. This observation created a problem in the early development of the periodic table because it meant that argon should be placed after potassium. (a) How was this difficulty resolved? (b) From the following data, calculate the average atomic masses of argon and potassium: Ar-36 (35.9675 amu, 0.337 percent), Ar-38 (37.9627 amu, 0.063 percent), Ar-40 (39.9624 amu, 99.60 percent), K-39 (38.9637 amu, 93.258 percent), K-40 (39.9640 amu, 0.0117 percent), K-41 (40.9618 amu, 6.730 percent).


  • 7.128

    Calculate the maximum wavelength of light (in nm) required to ionize a single sodium atom.

  • 7.129

    Predict the atomic number and ground-state electron configuration of the next member of the alkali metals after francium.


  • 7.130

    Why do elements that have high ionization energies also have more positive electron affinities? Which group of elements would be an exception to this generalization?

  • 7.131

    The first four ionization energies of an element are approximately 738, 1450, 7.7 × 103, and 1.1 × 104 kJ/mol. To which periodic group does this element belong? Explain your answer.


  • 7.132

    Some chemists think that helium should properly be called “helon.” Why? What does the ending in helium (-ium) suggest?

  • 7.133

    (a) The formula of the simplest hydrocarbon is CH4 (methane). Predict the formulas of the simplest compounds formed between hydrogen and the following elements: silicon, germanium, tin, and lead. (b) Sodium hydride (NaH) is an ionic compound. Would you expect rubidium hydride (RbH) to be more or less ionic than NaH? (c) Predict the reaction between radium (Ra) and water. (d) When exposed to air, aluminum forms a tenacious oxide (Al2O3) coating that protects the metal from corrosion. Which metal in Group 2A would you expect to exhibit similar properties? (See Page 278.)


  • 7.134

    Match each of the elements on the right with its description on the left:

    (a) A pale yellow gas that reacts with water Nitrogen (N2)
    (b) A soft metal that reacts with water to produce hydrogen Boron (B)
    (c) A metalloid that is hard and has a high melting point. Fluorine (F2)
    (d) A colorless, odorless gas. Aluminum (Al)
    (d) A metal that is more reactive than iron, but does not corrode in air. Sodium (Na)
  • 7.135

    Write at least two paragraphs describing the importance of the periodic table. Pay particular attention to the significance of the position of an element in the table and how the position relates to the chemical and physical properties of the element.


  • 7.136

    On one graph, plot the effective nuclear charge (shown in parentheses) and atomic radius (see Figure 7.6) versus atomic number for the second-period elements: Li(1.28), Be(1.91), B(2.42), C(3.14), N(3.83), O(4.45), F(5.10), Ne(5.76). Comment on the trends.

  • 7.137

    One allotropic form of an element X is a colorless crystalline solid. The reaction of X with an excess amount of oxygen produces a colorless gas. This gas dissolves in water to yield an acidic solution. Choose one of the following elements that matches X:

    1. sulfur,

    2. phosphorus,

    3. carbon,

    4. boron,

    5. silicon.


  • Page 295
  • 7.138

    The ionization energy of a certain element is 412 kJ/mol. When the atoms of this element are in the first excited state, however, the ionization energy is only 126 kJ/mol. Based on this information, calculate the wavelength of light emitted in a transition from the first excited state to the ground state.

  • 7.139

    One way to estimate the effective charge (Zeff) of a many-electron atom is to use the equation IE1 = (1312 kJ/mol)(Z2eff/n2), where IE1 is the first ionization energy and n is the principal quantum number of the shell in which the electron resides. Use this equation to calculate the effective charges of Li, Na, and K. Also calculate Zeff/n for each metal. Comment on your results.


  • 7.140

    Use your knowledge of thermochemistry to calculate the ΔH for the following processes: and

  • 7.141

    To prevent the formation of oxides, peroxides, and superoxides, alkali metals are sometimes stored in an inert atmosphere. Which of the following gases should not be used for lithium: Ne, Ar, N2, Kr? Explain. (Hint: As mentioned in the chapter, Li and Mg exhibit a diagonal relationship. Compare the common compounds of these two elements.)


Pre-Professional Practice Exam Problems: Physical And Biological Sciences

These questions are not based on a descriptive passage.

  1. A halogen has valence electrons in which orbitals?

    1. s

    2. s and p

    3. p

    4. s, p, and d

  2. How many subshells does a shell with principal quantum number n contain?

    1. n

    2. n2

    3. n −1

    4. 2n −1

  3. In a shell that contains an f subshell, what is the ratio of f orbitals to s orbitals?

    1. 14:1

    2. 7:1

    3. 7:3

    4. 7:5

  4. What is the maximum number of electrons that can be in the n 5 = shell?

    1. 2

    2. 6

    3. 8

    4. 18

Answers to In-Chapter Materials
Answers to Practice Problems
  • 7.1A


  • 7.1B

    Bi < As < P.

  • 7.2A
    1. 1s2 2s2 2p6 3s2 3p3, p-block,

    2. 1s2 2s2 2p6 3s2 3p6 4s2, s-block,

    3. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5, p-block.

  • 7.2B
    1. Al,

    2. Zn,

    3. Sr.

  • 7.3A

    F < Se < Ge.

  • 7.3B

    P and Se.

  • 7.4A

    Mg, Mg.

  • 7.4B

    Rb has a smaller Zeff, IE2 for Rb corresponds to the removal of a core electron.

  • 7.5A


  • 7.5B

    Adding an electron to As involves pairing.

  • 7.6A

    The attractive force is slightly larger between +3.26 and −1.15 separated by 1.5 pm.

  • 7.6B

    1.51 pm.

  • 7.7A
    1. [Ne],

    2. [Ar],

    3. [Kr].

  • 7.7B

    N3−, O2−, F, Ne, Na+, Mg2+, Al3+.

  • 7.8
    1. [Ar] 3d6,

    2. [Ar] 3d9,

    3. [Kr] 4d10.

  • 7.8B


  • 7.9A

    Rb+ < Kr < Br < Se2−.

  • 7.9B

    F, O2−, N3−, Na+, Mg2+, Al3+.

Answers to Checkpoints
  • 7.1.1


  • 7.1.2


  • 7.1.3


  • 7.2.1


  • 7.2.2

    a, d, e.

  • 7.4.1


  • 7.4.2


  • 7.4.3


  • 7.4.4


  • 7.5.1

    b, c, e.

  • 7.5.2

    b, d.

  • 7.5.3


  • 7.5.4


  • 7.5.5

    d, f.

  • 7.6.1

    d, e.

  • 7.6.2

    a, c.

  • 7.6.3


  • 7.6.4


Answers to Applying What You've Learned
  1. 1s2 2s1.

  2. In order of increasing atomic radius: Li < Na < K < Rb < Cs.

  3. In order of increasing ionization energy (IE1): Cs < Rb < K < Na < Li.

  4. Li+: 1s2 or [He]; Na1: 1s22s22p6 or [Ne]; K+: 1s22s22p63s23p6 or [Ar]; Rb+: 1s22s22p63s23p64s23d104p6 or [Kr]; Cs+: 1s22s22p63s23p64s23d104p65s24d105p6 or [Xe].

  5. Isoelectronic with Li+: He. Isoelectronic with Na+: Al3+, Mg2+, Ne, F, O2−, and N3−. Isoelectronic with K1: Ca2+, Ar, Cl, S2−, and P3−. Isoelectronic with Rb+: Sr2+, Kr, Br, and Se2−. Isoelectronic with Cs+: Ba2+, Xe, I, and Te2−.