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Chapter 6

Chapter 6. Periodic Properties of the Elements. Development of the Periodic Table. Elements in the same group generally have similar chemical properties. Properties are not identical, however. O 2 gas In a gas burette. S 8 rings stack to make Orthorhombic Crystals.

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Chapter 6

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  1. Chapter 6 Periodic Propertiesof the Elements

  2. Development of the Periodic Table • Elements in the same group generally have similar chemical properties. • Properties are not identical, however.

  3. O2 gas In a gas burette S8 rings stack to make Orthorhombic Crystals

  4. Development of the Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped. Figure 6.2

  5. 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. Table 6.1

  6. 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.

  7. 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. Figure 6.3

  8. Sizes of Atoms and Ions The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. Figure 6.4

  9. Sizes of Atoms Bonding atomic radius tends to: • decrease from left to right across a row due to increasing Zeff • increase from the top to the bottom of a column due to increasing value of n Figure 6.5

  10. Sizes of Ions • Ionic size depends upon: • Nuclear charge • Number of electrons • Orbitals in which electrons reside Figure 6.6

  11. Sizes of Ions • Cations are smaller than their parent atoms: • The outermost electron is removed and repulsions are reduced. Figure 6.6

  12. Sizes of Ions • Anions are larger than their parent atoms: • Electrons are added and repulsions are increased. Figure 6.6

  13. Sizes of Ions • Ions increase in size as you go down a column • Due to increasing value of n. • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.

  14. Ionisation Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion: • First ionisation energy is that energy required to remove the first electron. • Second ionisation energy is that energy required to remove the second electron, etc.

  15. Ionisation Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionisation energy takes a quantum leap. Table 6.2

  16. Trends in First Ionisation 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. Figure 6.8

  17. Trends in First Ionisation Energies • Generally, as one goes across a row, it gets harder to remove an electron: • As you go from left to right, Zeff increases. Figure 6.7

  18. Electron Affinities Energy change accompanying the addition of an electron to a gaseous atom: Cl + e− Cl− E = -349 kJ/mol

  19. Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. Figure 6.9

  20. Metal, Nonmetals and Metalloids Figure 6.10

  21. Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties. Table 6.3

  22. Metals versus Nonmetals • Metals tend to form cations. • Nonmetals tend to form anions. Figure 6.12

  23. Metals Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity. Figure 6.11

  24. Metals • Compounds formed between metals and nonmetals tend to be ionic. • Metal oxides tend to be basic. Figure 6.15

  25. Reactivity of Metals • Metal reactivity varies with the type of metal. • Alkali metals are very reactive because it is very easy to remove the 1s1 electron. All alkali metals react vigorously with water to produce metal hydroxide and hydrogen gas: 2M(s) + 2H2O(l) -> 2MOH(aq) + H2(g) Figure 6.13

  26. Reactivity of Metals • Alkaline earth metals are less reactive than alkali metals. • Mg reacts only with steam, Bedoes not react with water, but others react readily with water. • Reactivity tends to increase as you move down groups. Figure 6.14

  27. 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. • Substances containing only nonmetals are molecular compounds. • Most nonmetal oxides are acidic. Figure 6.16

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

  29. Mn Oxides

  30. Hg and Br

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