Chapter 7 Periodic Properties of the Elements

1. Chapter 7Periodic Propertiesof the Elements

2. Ch 7.1 Development of Periodic Table • The development of the periodic table has been ongoing since ancient times. (See left) • Since the 19th century discovery of new elements has been rapid. • In 1869, Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped. • Their arrangement of elements according to their atomic weights and the periodicity of their chemical and physical properties led to the modern periodic table. https://www.youtube.com/watch?v=0l_Z2zdT9gU

3. Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon, i.e.,“under silicon”) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon. • Elements within groups on the periodic table generally have the same characteristics due to having the same valence electron configurations, though at different principle energy levels. • However, properties are not always similar. For example, O is above S in group 16. They share some similarities due to their valance shell arrangements (ns2 np4), but are very different in other characteristics. For example, oxygen is a colorless gas but sulfur is a yellow solid.

4. 7.2 Effective Nuclear Charge • Properties of many-electron atoms depend not only on their electron configurations but also on how strongly their outer electrons are attracted to the nucleus. • The strength of the interaction depends on the magnitude of the nuclear charge acting on an electron and the average distance between the nucleus and the electron. • The force of attraction increases as the nuclear charge (atomic number) increases and decreases as the electron moves farther from the nucleus. • So, in a many-electron atom, each electron is simultaneously attracted to the nucleus and repelled by other electrons. • Each electron can be treated as though it were moving in a net electric field created by the nucleus and the electron density of the other electrons.

5. The net positive charge experienced by a valence electron is called the effective nuclear charge, Zeff. This charge is not the full nuclear charge due to the shielding of the nucleus by other electrons in the atom. • Zeff, is found this way: Zeff = Z−S where Z is the atomic number and S is a screening constant. • S represents the portion of the nuclear charge that is screened from the valence electron by other electrons. • Because inner (core) electrons are most effective at screening a valance electron, the value of S is usually close to the number of core electrons in an atom.

6. Example: Zeff for sodium. • The expected magnitude of Zeff for the Na valence (3s) electron is 1+, as shown in fig (a) at the left. • Now, if the 3s electron were completely outside the electron density for the core electrons this would be true. However, as as you can see from fig (b), a 3s electron has a small probability of being found inside the core electrons and much closer to the nucleus. • For this reason, the actual Zeff for the Na 3s electron is 2.5+. (The calculations for this go beyond our discussion at this level.)

7. Trends for Zeff: • Zeff increases moving across the PT. • Why? Well, the overall nuclear charge (Z) increases. • The number of core electrons, however, stay the same. The added valence electrons ineffectively shield each other. • Zeff increases slightly going down a group/family • The change is less than going across a row. • The Zeff increases because larger electron cores are less able to screen the outer electrons from the nuclear charge

8. 7.3 Sizes of Atoms and Ions

9. 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

10. Sizes of Ions • Ionic size depends upon: • Nuclear charge. • Number of electrons. • Orbitals in which electrons reside.

11. Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions are reduced.

12. Anions are larger than their parent atoms. • Electrons are added and repulsions are increased.

13. Ions increase in size as you go down a column. • Due to increasing value of n.

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

15. 7.4 Ionization Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc.

16. It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

17. 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.

18. Generally, as one goes across a row, it gets harder to remove an electron. • As you go from left to right, Zeff increases.

19. However, there are two apparent discontinuities in this trend.

20. The first occurs between Groups IIA and IIIA. • Electron removed from p-orbital rather than s-orbital • Electron farther from nucleus • Small amount of repulsion by s electrons.

21. The second occurs between Groups VA and VIA. • Electron removed comes from doubly occupied orbital. • Repulsion from other electron in orbital helps in its removal.

22. 7.5 Electron Affinity Energy change accompanying addition of electron to gaseous atom: Cl + e− Cl−

23. Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row.

24. There are again, however, two discontinuities in this trend.

25. The first occurs between Groups IA and IIA. • Added electron must go in p-orbital, not s-orbital. • Electron is farther from nucleus and feels repulsion from s-electrons.

26. The second occurs between Groups IVA and VA. • Group VA has no empty orbitals. • Extra electron must go into occupied orbital, creating repulsion.

27. 7.6 Properties of Metal, Nonmetals,and Metalloids

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

29. Metals tend to form cations. • Nonmetals tend to form anions.

30. Metals Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

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

32. Nonmetals • Dull, brittle substances that are poor conductors of heat and electricity. • Tend to gain electrons in reactions with metals to acquire noble gas configuration.

33. Substances containing only nonmetals are molecular compounds. • Most nonmetal oxides are acidic.

34. Metalloids • Have some characteristics of metals, some of nonmetals. • For instance, silicon looks shiny, but is brittle and fairly poor conductor.

35. 7.7 Group Trends for the Active Metals

36. Alkali Metals • Soft, metallic solids. • Name comes from Arabic word for ashes.

37. Found only as compounds in nature. • Have low densities and melting points. • Also have low ionization energies.

38. Their reactions with water are famously exothermic.

39. Alkali metals (except Li) react with oxygen to form peroxides. • K, Rb, and Cs also form superoxides: K + O2 KO2 • Produce bright colors when placed in flame.

40. Alkaline Earth Metals • Have higher densities and melting points than alkali metals. • Have low ionization energies, but not as low as alkali metals.

41. Bedoes not react with water, Mg reacts only with steam, but others react readily with water. • Reactivity tends to increase as go down group.

42. Group 6A (16) • Oxygen, sulfur, and selenium are nonmetals. • Tellurium is a metalloid. • The radioactive polonium is a metal.

43. Oxygen • Two allotropes: • O2 • O3, ozone • Three anions: • O2−, oxide • O22−, peroxide • O21−, superoxide • Tends to take electrons from other elements (oxidation)

44. Sulfur • Weaker oxidizing agent than oxygen. • Most stable allotrope is S8, a ringed molecule.

45. Group VIIA: Halogens • Prototypical nonmetals • Name comes from the Greek halos and gennao: “salt formers”

46. Large, negative electron affinities • Therefore, tend to oxidize other elements easily • React directly with metals to form metal halides • Chlorine added to water supplies to serve as disinfectant

47. Group VIIIA: Noble Gases • Astronomical ionization energies • Positive electron affinities • Therefore, relatively unreactive • Monatomic gases

48. Xe forms three compounds: • XeF2 • XeF4 (at right) • XeF6 • Kr forms only one stable compound: • KrF2 • The unstable HArF was synthesized in 2000.