1. Chapter 7 Periodic Properties of the Elements
Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 7 Periodic Properties of the Elements
John D. Bookstaver
St. Charles Community College
Cottleville, MO
© 2009, Prentice-Hall, Inc.

2. Overview of Periodic Trends
Why and how show we study trends in the periodic table?
Focus on valence electrons  understand and predict reactivity of elements
© 2009, Prentice-Hall, Inc.

3. Overview of Periodic Trends
Descriptions and explanations of trends in the periodic table:
Key concept: Effective nuclear charge  effects on the valence electron(s)
a) Atomic radius
b) Ionic radius
c) Ionization energy
d) Electron affinity
e) Electronegativity (in bonding chapters)
Be able to predict and explain the trends, using effective nuclear charge and atomic structure.
© 2009, Prentice-Hall, Inc.

4. Overview of Periodic Trends
2. Effects on the properties of metals, nonmetals and metalloids
 group trends in terms of reactions
 Periodic Trends lab and demo
(next week)
© 2009, Prentice-Hall, Inc.

5. Development of Periodic Table
Elements in the same group generally have similar chemical properties.
Physical properties are not necessarily similar, however.
© 2009, Prentice-Hall, Inc.

6. Development of Periodic Table
Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.
© 2009, Prentice-Hall, Inc.

7. Development of Periodic Table
Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.
© 2009, Prentice-Hall, Inc.

8. Effective Nuclear Charge
In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons.
The nuclear charge that an electron experiences depends on both factors.
© 2009, Prentice-Hall, Inc.

9. Effective Nuclear Charge
The effective nuclear charge, Zeff, is found this way:
Zeff = Z − S
where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.
© 2009, Prentice-Hall, Inc.

10. What Is the Size of an Atom?
The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.
© 2009, Prentice-Hall, Inc.

11. Sizes of Atoms Bonding atomic radius tends to…
…decrease from left to right across a row
(due to increasing Zeff).
…increase from top to bottom of a column
(due to increasing value of n).
© 2009, Prentice-Hall, Inc.

12. Sample Exercise 7.1 (p. 262)
Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH3SH, whose structure is shown below. Use Figure 7.7 to predict the lengths of the C-S, C-H, and S-H bonds in this molecule.
© 2009, Prentice-Hall, Inc.

13. Practice Exercise 7.1
Using Figure 7.7, predict which will be greater, the P-Br bond length in PBr3 or the As-Cl bond length in AsCl3.
© 2009, Prentice-Hall, Inc.

14. Sample Exercise 7.2 (p. 263)
Referring to a periodic table, arrange (as much as possible) the following atoms in order of increasing size: 15P, 16S, 33As, 34Se. (Atomic numbers are given for the elements to help you locate them quickly in the periodic table.)
© 2009, Prentice-Hall, Inc.

15. Practice Exercise 7.2
Arrange the following atoms in order of increasing atomic radius: 11Na, 4Be, 12Mg.
© 2009, Prentice-Hall, Inc.

16. Sizes of Ions Ionic size depends upon: The nuclear charge.
The number of electrons.
The orbitals in which electrons reside.
© 2009, Prentice-Hall, Inc.

17. Sizes of Ions Cations are smaller than their parent atoms.
The outermost electron is removed and repulsions between electrons are reduced.
© 2009, Prentice-Hall, Inc.

18. Sizes of Ions Anions are larger than their parent atoms.
Electrons are added and repulsions between electrons are increased.
© 2009, Prentice-Hall, Inc.

19. Sizes of Ions Ions increase in size as you go down a column.
This is due to increasing value of n.
© 2009, Prentice-Hall, Inc.

20. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.
© 2009, Prentice-Hall, Inc.

21. Sample Exercise 7.3 (p. 264)
Arrange these atoms and ions in order of decreasing size: Mg2+, Ca2+, Ca.
© 2009, Prentice-Hall, Inc.

22. Practice Exercise 7.3
Which of the following atoms and ions is largest: S2-, S, O2-?
© 2009, Prentice-Hall, Inc.

23. Sample Exercise 7.4 (p. 264)
Arrange the ions S2-, Cl-, K+ and Ca2+ in order of decreasing size.
© 2009, Prentice-Hall, Inc.

24. Practice Exercise 7.4
Which of the following ions is largest? Rb+, Sr2+, or Y3+?
© 2009, Prentice-Hall, Inc.

25. Ionization Energy
The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
The first ionization energy is the energy required to remove first electron.
The second ionization energy is that energy required to remove second electron, etc.
© 2009, Prentice-Hall, Inc.

26. Ionization Energy
It requires more energy to remove each successive electron.
When all valence electrons have been removed, the ionization energy takes a quantum leap.
© 2009, Prentice-Hall, Inc.

27. Sample Exercise 7.5 (p. 266)
Three elements are indicated in the periodic table below. Based on their locations, predict the one with the largest second ionization energy.
© 2009, Prentice-Hall, Inc.

28. Practice Exercise 7.5
Which will have the greater third ionization energy, Ca or S?
© 2009, Prentice-Hall, Inc.

29. Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.
© 2009, Prentice-Hall, Inc.

30. Trends in First Ionization Energies
Generally, as one goes across a row, it gets harder to remove an electron.
As you go from left to right, Zeff increases.
© 2009, Prentice-Hall, Inc.

31. Trends in First Ionization Energies
However, there are two apparent discontinuities in this trend.
© 2009, Prentice-Hall, Inc.

32. Trends in First Ionization Energies
The first occurs between Groups IIA and IIIA.
In this case the electron is removed from a p-orbital rather than an s-orbital.
The electron removed is farther from nucleus.
There is also a small amount of repulsion by the s electrons.
© 2009, Prentice-Hall, Inc.

33. Trends in First Ionization Energies
The second occurs between Groups VA and VIA.
The electron removed comes from doubly occupied orbital.
Repulsion from the other electron in the orbital aids in its removal.
© 2009, Prentice-Hall, Inc.

34. Sample Exercise 7.6 (p. 268)
Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy:
Ne, Na, P, Ar, K.
© 2009, Prentice-Hall, Inc.

35. Practice Exercise 7.6
Which hs the lowest first ionization energy, B, Al, C, or Si?
Which has the highest first ionization energy?
© 2009, Prentice-Hall, Inc.

36. Sample Exercise 7.7 (p. 269) Write the electron configuration for the
a) Ca2+ ion;
b) Co3+ ion;
c) S2- ion.
© 2009, Prentice-Hall, Inc.

37. Practice Exercise 7.7 Write the electron configuration for the
a) Ga3+ ion;
b) Cr3+ ion;
c) Br- ion.
© 2009, Prentice-Hall, Inc.

38. Electron Affinity
Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:
Cl + e−  Cl−
© 2009, Prentice-Hall, Inc.

39. Ionization Energy vs. Electron Affinity
Ionization energy: how much does it cost to remove an electron?
Always positive (+kJ)
Electron affinity: What is the energy change when an electron is added to an atom?
Positive or negative (+kJ or –kJ)
Together ionization energy and electron affinity
electronegativity
bonding and reactivity
© 2009, Prentice-Hall, Inc.

40. Trends in Electron Affinity
In general, electron affinity becomes more exothermic as you go from left to right across a row.
© 2009, Prentice-Hall, Inc.

41. Trends in Electron Affinity
There are again, however, two discontinuities in this trend.
© 2009, Prentice-Hall, Inc.

42. Trends in Electron Affinity
The first occurs between Groups IA and IIA.
The added electron must go in a p-orbital, not an s-orbital.
The electron is farther from nucleus and feels repulsion from the s-electrons.
© 2009, Prentice-Hall, Inc.

43. Trends in Electron Affinity
The second occurs between Groups IVA and VA.
Group VA has no empty orbitals.
The extra electron must go into an already occupied orbital, creating repulsion.
© 2009, Prentice-Hall, Inc.

44. Properties of Metal, Nonmetals, and Metalloids
© 2009, Prentice-Hall, Inc.

45. Metals
They tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.
© 2009, Prentice-Hall, Inc.

46. Metals Compounds formed between metals and nonmetals tend to be ionic.
Metal oxides tend to be basic when dissolved in water.
© 2009, Prentice-Hall, Inc.

47. Reactions of Metal Oxides
Most metal oxides are basic:
Metal oxide + water  metal hydroxide
(synthesis reaction)
Na2O(s) + H2O(l)  2NaOH(aq)
© 2009, Prentice-Hall, Inc.

48. Reactions of Metal Oxides
Metal oxides are able to react with acids to form salts and water:
Metal oxide + acid  salt + water
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)
© 2009, Prentice-Hall, Inc.

49. Sample Exercise 7.8 (p. 274)
a) Would you expect scandium oxide to be a solid, liquid, or gas at room temperature?
b) Write the balanced chemical equation for the reaction of scandium oxide with nitric acid.
© 2009, Prentice-Hall, Inc.

50. Practice Exercise 7.8
Write the balanced chemical equation for the reaction between copper (II) oxide and sulfuric acid.
© 2009, Prentice-Hall, Inc.

51. Metals versus Nonmetals
Differences between metals and nonmetals tend to revolve around these properties.
© 2009, Prentice-Hall, Inc.

52. Metals versus Nonmetals
Metals tend to form cations (are oxidized), including oxides (think rust).
Nonmetals tend to form anions (are reduced).
© 2009, Prentice-Hall, Inc.

53. Nonmetals
These are dull, brittle substances that are poor conductors of heat and electricity.
They tend to gain electrons in reactions with metals to acquire a noble gas configuration.
© 2009, Prentice-Hall, Inc.

54. Nonmetals
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic when dissolved in water.
© 2009, Prentice-Hall, Inc.

55. Nonmetals
Seven nonmetallic elements exist as diatomic molecules under ordinary conditions:
H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s)
Remember the “7” rule.
© 2009, Prentice-Hall, Inc.

56. Reactions - Nonmetals
When nonmetals react with metals, nonmetals tend to gain electrons:
Metal + nonmetal  salt
2Al(s) + 3Br2(l)  2AlBr3 (s)
© 2009, Prentice-Hall, Inc.

57. Reactions - Nonmetals
Most nonmetal oxides are acidic when dissolved in acid:
Nonmetal oxide + water  acid (synthesis)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
CO2(g) + H2O(l)  H2CO3(aq)
© 2009, Prentice-Hall, Inc.

58. Reactions - Nonmetals
Nonmetal oxides react with bases to form salts and water:
Nonmetal oxide + base  salt + water
CO2(g)+ 2NaOH(aq)  Na2CO3(aq) + H2O(l)
© 2009, Prentice-Hall, Inc.

59. Sample Exercise 7.9 (p. 256)
Write the balanced chemical equations for the reactions of solid selenium dioxide with
a) water
b) aqueous sodium hydroxide
© 2009, Prentice-Hall, Inc.

60. Practice Exercise 7.9
Write the balanced chemical equation for the reaction of solid tetraphosphorus hexoxide with water.
© 2009, Prentice-Hall, Inc.

61. Metalloids
These have some characteristics of metals and some of nonmetals.
For instance, silicon looks shiny, but is brittle and fairly poor conductor.
© 2009, Prentice-Hall, Inc.

62. Group Trends
© 2009, Prentice-Hall, Inc.

63. Alkali Metals Alkali metals are soft, metallic solids.
The name comes from the Arabic word for ashes.
© 2009, Prentice-Hall, Inc.

64. Alkali Metals
They are found only in compounds in nature, not in their elemental forms.
They have low densities and melting points.
They also have low ionization energies.
© 2009, Prentice-Hall, Inc.

65. Alkali Metals Reactions
Alkali metals react with hydrogen to form hydrides.
In hydrides, the hydrogen is present as H–, called the hydride ion.
2M(s) + H2(g)  2MH(s)
© 2009, Prentice-Hall, Inc.

66. Alkali Metals
Alkali metals react with water to form MOH and hydrogen gas:
e.g. Na(s) + H2O(l)  NaOH(aq) + H2(g)
© 2009, Prentice-Hall, Inc.

67. Alkali Metals Reactions
Alkali metals produce different oxides when reacting with O2:
4Li(s) + O2(g)  2Li2O(s) (oxide)
2Na(s) + O2 (g)  Na2O2(s) (peroxide)
K(s) + O2 (g)  KO2(s) (superoxide)
© 2009, Prentice-Hall, Inc.

68. Alkali Metals
Alkali metals (except Li) react with oxygen to form peroxides.
K, Rb, and Cs also form superoxides:
K + O2  KO2
They produce bright colors when placed in a flame.
© 2009, Prentice-Hall, Inc.

69. Sample Exercise 7.10 (p. 279)
Write balanced chemical equations that predict the reactions of cesium metal with
a) Cl2(g)
b) H2O(l)
c) H2(g)
© 2009, Prentice-Hall, Inc.

70. Practice Exercise 7.10
Write a balanced chemical equation that predicts the products of the reaction between potassium metal and elemental sulfur (S8).
© 2009, Prentice-Hall, Inc.

71. Alkaline Earth Metals
Alkaline earth metals have higher densities and melting points than alkali metals.
Their ionization energies are low, but not as low as those of alkali metals.
© 2009, Prentice-Hall, Inc.

72. Alkaline Earth Metals Reactions
Their chemistry is dominated by the loss of two s electrons:
M  M2+ + 2e–
Mg(s) + Cl2(g)  MgCl2(s)
2Mg(s) + O2(g)  2MgO(s)
(remember last year’s Reactions labs – white flame?
© 2009, Prentice-Hall, Inc.

73. Alkaline Earth Metals Reactivity increases down the group.
Beryllium does not react with water and magnesium reacts only with steam, but the others react readily with water.
Thursday’s lab
© 2009, Prentice-Hall, Inc.

74. Hydrogen
Hydrogen is a unique element, most often occurs as a colorless diatomic gas, H2.
Reactions between hydrogen and nonmetals can be very exothermic:
2H2(g) + 2O2(g)  2H2O(l)
∆Ho = –571.7 kJ
© 2009, Prentice-Hall, Inc.

75. Hydrogen
It can either gain another electron to form the hydride ion, H–, or lose its electron to become H+:
2Na(s) + H2(g)  2NaH(s)
2H2(g) + O2(g)  2H2O(l)
© 2009, Prentice-Hall, Inc.

76. Hydrogen
H+ is a proton.
The aqueous chemistry of hydrogen is dominated by H+(aq).
(acids – next semester)
© 2009, Prentice-Hall, Inc.

77. Group 6A Oxygen, sulfur, and selenium are nonmetals.
Tellurium is a metalloid.
The radioactive polonium is a metal.
© 2009, Prentice-Hall, Inc.

78. Oxygen There are two allotropes of oxygen: There can be three anions:
O3, ozone
There can be three anions:
O2−, oxide
O22−, peroxide
O21−, superoxide
It tends to take electrons from other elements (oxidation).
© 2009, Prentice-Hall, Inc.

79. Sulfur Sulfur is a weaker oxidizer than oxygen.
The most stable allotrope is S8, a ringed molecule.
© 2009, Prentice-Hall, Inc.

80. Group 7A: Halogens The halogens are prototypical nonmetals.
The name comes from the Greek words halos and gennao: “salt formers”.
All halogens are diatomic: X2.
© 2009, Prentice-Hall, Inc.

81. Group 7A: Halogens
The chemistry of the halogens is dominated by gaining an electron to form an anion:
X2 + 2e–  2X–
Fluorine is one of the most reactive substances known:
2F2(g) + 2H2O(l)  4HF(aq) + O2(g) DH = –758.9 kJ
© 2009, Prentice-Hall, Inc.

82. Group 7A: Halogens They have large, negative electron affinities.
Therefore, they tend to oxidize other elements easily.
They react directly with metals to form metal halides.
Chlorine is added to water supplies to serve as a disinfectant
© 2009, Prentice-Hall, Inc.

83. Group 7A: Halogens
Chlorine is added to water supplies to serve as a disinfectant
e.g. swimming pools:
Cl2(g)+ H2O(l)  HCl(aq) + HOCl(aq)
© 2009, Prentice-Hall, Inc.

84. Group 8A: Noble Gases
The noble gases have astronomical ionization energies.
Their electron affinities are positive.
Therefore, they are relatively unreactive.
They are found as monatomic gases.
© 2009, Prentice-Hall, Inc.

85. Group 8A: Noble Gases Xe forms three compounds:
XeF4 (at right)
XeF6
Kr forms only one stable compound:
KrF2
The unstable HArF was synthesized in 2000.
© 2009, Prentice-Hall, Inc.

86. Sample Integrative Exercise 7 (p. 285)
The element bismuth (Bi, atomic number 83) is the heaviest member of group 5A. A salt of the element, bismuth subsalicylate, is the active ingredient in Pepto-Bismol, an over-the-counter medication for gastric distress.
The covalent atomic radii of thallium (Tl) and lead (Pb) are 1.48 Å and 1.47 Å, respectively. Using these values and what you know about trends in atomic radii, predict the covalent atomic radius of the element bismuth. Explain your answer.
What accounts for the general increase in atomic radius going down the group 5A elements?
3. Another major use of bismuth has been as an ingredient in low-melting metal alloys, such as those used in fire sprinkler systems and in typesetting. The element itself is a brittle white crystalline solid.
How do these characteristics fit with the fact that bismuth is in the same periodic group with such nonmetallic elements as nitrogen and phosphorus?
© 2009, Prentice-Hall, Inc.

87. Sample Integrative Exercise 7 (p. 285)
Bi2O3 is a basic oxide. Write a balanced equation for its reaction with dilute nitric acid. If 6.77 g of Bi2O3 is dissolved in dilute acidic solution to make up 500 mL of solution, what is the molarity of the Bi3+ ion?
209Bi is the heaviest stable isotope of any element. How many protons and neutrons are present in this nucleus?
6. The density of Bi at 25oC is g/cm3. How many Bi atoms are present in a cube of the element that is 5.00 cm on each edge? How many moles of the element are present?
© 2009, Prentice-Hall, Inc.