Chapter 7 Periodic Properties of the Elements

1. Chapter 7Periodic Propertiesof the Elements

2. Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

3. Development of the 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.

4. A. Lanthanides because some of the elements are not stable B. Row 6, missing atomic numbers of elements between Ba and Lu C. Row 6, Rn is the start of elements that are radioactive D. Row 7, elements in this row are generally radioactive

5. Atomic weight has uncertainty in measurement, whereas atomic number depends only on number of protons. • Atomic weight depends both on number of protons and neutrons. • Atomic weight is greatly influenced by number of electrons, not protons. • Atomic weight depends both on number of protons and electrons.

6. Co/Ni • Th/Pa • Te/I • All of the above

7. Periodic Trends • In this chapter, we will rationalize observed trends in • Sizes of atoms and ions. • Ionization energy. • Electron affinity.

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. 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.

9. A. 2s B. 2p

10. 2p electron of a Ne atom • 3s electron of a Na atom • Both experience the same effective nuclear charge • Requires a table of shielding constants to make an estimation

11. What Is the Size of an Atom? The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

12. Sizes of Atoms The 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 the increasing value of n).

13. A. Bottom and left B. Bottom and right C. Top and left D. Top and right

14. Sample Exercise 7.1Bond Lengths in a Molecule 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. Use Figure 7.6 to predict the lengths of the C—S, C—H, and S—H bonds in this molecule. FIGURE 7.6 Trends in bonding atomic radii for periods 1 through 5.

15. Trends work with each other but trend in orbital size is more important. • Trends work with each other but trend in effective nuclear charge is more important. • Trends work against each other but trend in orbital size is more important. • Trends work against each other but trend in effective nuclear charge is more important.

16. Sample Exercise 7.2Atomic Radii Referring to a periodic table, arrange (as much as possible) the atoms 15P, 16S, 33As, and 34Se in order of increasing size. (Atomic numbers are given to help you locate the atoms quickly in the table.)

17. Sizes of Ions • Ionic size depends upon • The nuclear charge. • The number of electrons. • The orbitals in which electrons reside. • Cations are smaller than their parent atoms: • The outermost electron is removed and repulsions between electrons are reduced. • Anions are larger than their parent atoms” • Electrons are added and repulsions between electrons are increased.

18. Sizes of Ions • Ions increase in size as you go down a column: • This increase in size is due to the increasing value of n.

19. Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.

20. A. Minimal change B. Larger C. Smaller D. Trend varies by column

21. Sample Exercise 7.3Atomic and Ionic Radii Arrange Mg2+, Ca2+, and Ca in order of decreasing radius..

22. Sample Exercise 7.4Ionic Radii in an Isoelectronic Series Arrange the ions K+, Cl–, Ca2+, and S2– in order of decreasing size.

23. 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 that energy required to remove the first electron. • The second ionization energy is that energy required to remove the second electron, etc.

24. First ionization step, because it takes more energy and hence shorter wavelength. • First ionization step, because it takes less energy and hence shorter wavelength. • Second ionization step, because it takes more energy and hence shorter wavelength. • Second ionization step, because it takes less energy and hence shorter wavelength.

25. 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.

26. A. I1 for the boron atom is equal to I2 for the carbon atom. B. I2 for the carbon atom is greater. C. I1 for the boron atom is greater.

27. Sample Exercise 7.5Trends in Ionization Energy Three elements are indicated in the periodic table in the margin. Which one has the largest second ionization energy?

28. 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.

29. 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. • However, there are two apparent discontinuities in this trend.

30. 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 the nucleus. • There is also a small amount of repulsion by the s electrons.

31. Trends in First Ionization Energies • The second discontinuity occurs between Groups VA and VIA. • The electron removed comes from a doubly occupied orbital. • Repulsion from the other electron in the orbital aids in its removal.

32. A. As; outer electrons experience larger Zeff than those in Ar. B. As; outer electrons experience smaller Zeff than those in Ar. C. Ar; outer electrons experience larger Zeff than those in As. D. Ar; outer electrons experience smaller Zeff than those in As.

33. A. The 2p electron removed from an oxygen atom is in a fully occupied orbital with decreased electron-electron repulsions compared to the one from a nitrogen atom. B. The 2p electron removed from an oxygen atom is in a more than half-filled subshell, which has extra stability compared to the one from a nitrogen atom. C. The 2p electron removed from an oxygen atom is in a fully occupied orbital with increased electron-electron repulsions compared to the one from a nitrogen atom. D. The 2p electron removed from an oxygen atom is in a lower energy state compared to the one from a nitrogen atom.

34. Sample Exercise 7.6Periodic Trends in Ionization Energy Referring to a periodic table, arrange the atoms Ne, Na, P, Ar, K in order of increasing first ionization energy. FIGURE 7.9 Trends in first ionization energies of the elements

35. Same electron configurations: [Ar]3d3 • Different electron configurations

36. Sample Exercise 7.7Electron Configurations of Ions Write the electron configuration for (a) Ca2+, (b) Co3+, and (c) S2–.

37. Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e− Cl−

38. Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. There are again, however, two discontinuities in this trend.

39. 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 the nucleus and feels repulsion from the s electrons.

40. Trends in Electron Affinity • The second discontinuity occurs between Groups IVA and VA. • Group VA has no empty orbitals. • The extra electron must go into an already occupied orbital, creating repulsion.

41. A. Group 4A; outer electrons of atoms in this family experience greater Zeff than the elements in other families of the same period and this results in greater attraction of an electron when electrons are added. B. Group 5A; outer electrons of atoms in this family experience greater Zeff than the elements in other families of the same period and this results greater attraction of an electron when electrons are added. C. Group 6A; outer electrons of atoms in this family experience greater Zeff than the elements in other families of the same period and this results in greater attraction of an electron when electrons are added. D. Group 7A; outer electrons of atoms in this family experience greater Zeff than the elements in other families of the same period and this results in greater attraction of an electron when an electron is added.

42. A. EA for Cl is significantly more endothermic than IE1 for Cl. B. IE1 for Cl– is significantly more exothermic than EA for Cl. C. They are equal in magnitude and opposite in sign. D. They are equal in magnitude and sign.

43. Properties of Metal, Nonmetals,and Metalloids

44. A. Electron affinity is increasing from Ge to Sn and this correlates with increasing metallic character. B. Ionization energy is decreasing from Ge to Sn and this correlates with increasing metallic character. C. Atomic radius is decreasing from Ge to Sn and this correlates with increasing metallic character. D. Number of valence electrons is increasing from Ge to Sn and this correlates with increasing metallic character.

45. Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties.

46. Metals versus Nonmetals • Metals tend to form cations. • Nonmetals tend to form anions.

47. Metals Metals tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

48. A. Cations are to the right of the red line and anions are to the left B. Higher oxidation states are adjacent to the red line and decrease moving to the right or left, excluding the transition elements in terms of trend in oxidation states. C. Higher oxidation states are adjacent to the red line and increase moving to the right or left, excluding the transition elements in terms of trend in oxidation states. D. There is no observable trend.

49. A. There is no correlation between ionization energy and metallic character. B. Increasing ionization energy correlates with decreasing metallic character. C. Increasing ionization energy correlates with increasing metallic character. D. Decreasing ionization energy correlates with decreasing metallic character.

50. Metals • Compounds formed between metals and nonmetals tend to be ionic. • Metal oxides tend to be basic.