1. Chapter 7 Periodic Properties of the Elements

2. Effective Nuclear Charge Z eff = Z − S Z = atomic number S = number of inner electrons

3. What Is the Size of an Atom?

4. Sizes of Atoms Atomic radius – Decreases from left to right across a row – Increases from top to bottom of a column

5. Sizes of Ions

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

7. Ionization Energy Ionization energy: 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 first electron. – The second ionization energy is that energy required to remove second electron, etc.

8. Ionization Energy

9. Trends in First Ionization Energies First ionization energy decreases as you go down a column – valence electrons are farther from the nucleus

10. Trends in First Ionization Energies First ionization energy increases from left to right – Z eff increases.

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

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

13. Properties of Metal, Nonmetals, and Metalloids

14. Metals versus Nonmetals Metals tend to form cations. Nonmetals tend to form anions.

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

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

17. 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.

18. Nonmetals Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic.

19. Metalloids These have some characteristics of metals and some of nonmetals. For instance, silicon looks shiny, but is brittle and fairly poor conductor.

20. Alkali Metals Alkali metals are soft, metallic solids.

21. 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.

22. Alkali Metals Their reactions with water are famously exothermic.

23. Alkali Metals Alkali metals (except Li) react with oxygen to form peroxides. They produce bright colors when placed in a flame.

24. 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.

25. Alkaline Earth Metals Reactivity tends to increase as you go down the group.

26. Group 6A Oxygen, sulfur, and selenium are nonmetals. Tellurium is a metalloid. The radioactive polonium is a metal.

27. Oxygen There are two allotropes of oxygen: – O 2 – O 3, ozone Two common anions: – O 2−, oxide – O 2 2−, peroxide It tends to take electrons from other elements (oxidation).

28. Sulfur Sulfur is a weaker oxidizer than oxygen. The most stable allotrope is S 8, a ringed molecule.

29. Group VIIA: Halogens The halogens are prototypical nonmetals. The name comes from the Greek words halos and gennao: “salt formers”.

30. Group VIIA: Halogens They have large, negative electron affinities. – Therefore, they tend to oxidize other elements easily. They react directly with metals to form metal halides.

31. Group VIIIA: 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.

32. Group VIIIA: Noble Gases Xe forms three compounds: – XeF 2 – XeF 4 (at right) – XeF 6 Kr forms only one stable compound: – KrF 2 The unstable HArF was synthesized in 2000.