1. Prentice Hall © 2003Chapter 7 Chapter 7 Periodic Properties of the Elements CHEMISTRY The Central Science 9th Edition

2. Prentice Hall © 2003Chapter 7 Dimitri Mendeleev and Lothar Meyer arranged the elements in order of increasing atomic weight Certain elements were missing from this scheme: Mendeleev noted: As properly belonged underneath P and not Si A missing element underneath Si He predicted a number of properties for this element In 1886 Ge was discovered Properties of Ge match Mendeleev’s predictions Modern periodic table: arrange elements in order of increasing atomic number 7.1: Development of the Periodic Table

3. Prentice Hall © 2003Chapter 7 Effective Nuclear Charge Effective nuclear charge is the charge experienced by an electron on a many-electron atom The effective nuclear charge is not the same as the charge on the nucleus Due to the effect of the inner electrons

4. Prentice Hall © 2003Chapter 7 Electrons are attracted to the nucleus, but repelled by the electrons that screen them from the nuclear charge The nuclear charge experienced by an electron depends on its distance from the nucleus the number of core electrons As the average number of screening electrons (S) increases, the effective nuclear charge (Z eff ) decreases OR…as the distance from the nucleus increases, S increases and Z eff decreases 7.2: Effective Nuclear Charge

5. Prentice Hall © 2003Chapter 7 The ns orbitals all have the same shape, but have different sizes and different numbers of nodes (places where the probability function = 0…see text, P. 215) The radial electron density is the probability of finding an electron at a given distance Text, P. 240 Z eff = Z – S or Protons – Electrons

6. Prentice Hall © 2003Chapter 7 Z eff increases across a period –Added electrons don’t increase shielding Z eff increases down a group –Added electrons are in increasingly higher energy levels but they are less able to shield the outer electrons from the nucleus The period trend is of greater significance than the group trend

7. Prentice Hall © 2003Chapter 7 Consider a simple diatomic molecule: The distance between the two nuclei is called the bond distance Half of the bond distance is called the covalent radius of the atom 7.3: Sizes of Atoms and Ions

8. Prentice Hall © 2003Chapter 7 Periodic Trends in Atomic Radii The properties of elements vary periodically Atomic size varies consistently through the periodic table Atomic radius increases down a group (more noticeable) Atomic radius decreases across a period There are two factors at work: Principal quantum number Electrons fill higher energy levels Effective nuclear charge Electrons fill the same energy level, no increase in shielding

9. Text, P. 242

10. Prentice Hall © 2003Chapter 7 Trends in the Sizes of Ions Ion size is the distance between ions in an ionic compound Cations vacate the most spatially extended orbital and are smaller than the parent atom Anions add electrons to the most spatially extended orbital and are larger than the parent atom Remember: electrons are gained or lost to achieve a noble gas configuration

11. Text P. 244

12. Prentice Hall © 2003Chapter 7 For ions of the same charge, ion size increases down a group Electrons in higher energy levels All the members of an isoelectronic series have the same number of electrons ( O 2-, F -, Na +, Mg 2+, Al 3+ ) As nuclear charge increases in an isoelectronic series, the ions become smaller: O 2- > F - > Na + > Mg 2+ > Al 3+ Added electrons are in the same energy level, so no increase in shielding Removed electrons result in an increasingly positive nucleus which pulls electrons toward it

13. Prentice Hall © 2003Chapter 7 7.4: Ionization Energy The first ionization energy, I 1, is the amount of energy required to remove an electron from a gaseous atom: Na(g)  Na + (g) + e - The second ionization energy, I 2, is the energy required to remove an electron from a gaseous ion: Na + (g)  Na 2+ (g) + e - The larger ionization energy, the more difficult it is to remove the electron

14. Prentice Hall © 2003Chapter 7

15. Prentice Hall © 2003Chapter 7 Variations in Successive Ionization Energies There is a sharp increase in ionization energy when a core electron is removed: Text, 246

16. Prentice Hall © 2003Chapter 7 Periodic Trends in Ionization Energies Think: does it want to lose an electron??? Ionization energy decreases down a group The outermost electron is more readily removed as electrons fill higher energy levels The s electrons are more effective at shielding than p electrons Irregularity in group 2: sharp increase in I 1

17. Prentice Hall © 2003Chapter 7 Text, P. 248

18. Prentice Hall © 2003Chapter 7 Ionization energy generally increases across a period Across a period, Zeff increases It becomes more difficult to remove an electron Irregularity in group 15: sharp increase in I 1 (half-filled sublevel) Transition metals and f-block elements show slow variations in I 1 due to shielding

19. Text, P. 248

20. Text, P. 247

21. Prentice Hall © 2003Chapter 7 Electron Configuration of Ions Cations: electrons removed from orbital with highest principle quantum number, n, first: Li (1s 2 2s 1 )  Li + (1s 2 ) Fe ([Ar]3d 6 4s 2 )  Fe 3+ ([Ar]3d 5 ) Anions: electrons added to the orbital with highest n: F (1s 2 2s 2 2p 5 )  F  (1s 2 2s 2 2p 6 )

22. Prentice Hall © 2003Chapter 7 7.5: Electron Affinities Think: does it want to gain an electron??? Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion: Cl(g) + e -  Cl - (g) Electron affinity can either be exothermic (above) or endothermic: Ar(g) + e -  Ar - (g)

23. EA becomes more negative across a period A greater attraction between the atom and the added electron Groups 2, 15 and 18 EA has no real group trend Text, P. 251

24. 7.6: Metals, Nonmetals, and Metalloids Text, P. 252

25. Prentice Hall © 2003Chapter 7 Metals Metallic character refers to the properties of metals shiny or lustrous malleable and ductile oxides form basic ionic solids form cations in aqueous solution Metallic character increases down a group Metallic character decreases across a period Metals have low ionization energies Most neutral metals are oxidized rather than reduced form cations

26. Prentice Hall © 2003Chapter 7 Text, P. 253

27. Prentice Hall © 2003Chapter 7 Most metal oxides are basic: Metal oxide + water  metal hydroxide Nonmetals Nonmetals are more diverse in their behavior than metals Properties are opposite those of metals Compounds composed only of nonmetals are molecular When nonmetals react with metals, nonmetals tend to gain electrons: metal + nonmetal  salt

28. Prentice Hall © 2003Chapter 7 Most nonmetal oxides are acidic: nonmetal oxide + water  acid Nonmetal oxide + base  salt + water The greater the nonmetallic character of the central atom, the stronger the acidic character of the oxide Metalloids Metalloids have properties that are intermediate between metals and nonmetals Si has a metallic luster but it is brittle Semiconductors

29. Prentice Hall © 2003Chapter 7 7.7: Group Trends for the Active Metals Group 1A: The Alkali Metals Alkali metals are all soft Chemistry dominated by the loss of their single s electron: M  M + + e - Reactivity increases down the group Alkali metals react with water to form MOH and hydrogen gas: 2M(s) + 2H 2 O(l)  2MOH(aq) + H 2 (g)

30. Prentice Hall © 2003Chapter 7 Alkali metals produce different oxides when reacting with O 2 : 4Li(s) + O 2 (g)  2Li 2 O(s)(oxide) 2Na(s) + O 2 (g)  Na 2 O 2 (s)(peroxide) K(s) + O 2 (g)  KO 2 (s)(superoxide) (Rb and Cs, too) Alkali metals: Flame tests The s electron is excited by the flame and emits energy when it returns to the ground state

31. Prentice Hall © 2003Chapter 7 Na line (589 nm): 3p  3s transition Li line: 2p  2s transition K line: 4p  4s transition

32. Prentice Hall © 2003Chapter 7 Text, P. 257

33. Prentice Hall © 2003Chapter 7 Group 2A: The Alkaline Earth Metals Alkaline earth metals are harder and more dense than the alkali metals The chemistry is dominated by the loss of two s electrons: M  M 2+ + 2e - Mg(s) + Cl 2 (g)  MgCl 2 (s) Increasing activity down the group: Be does not react with water Mg will only react with steam Ca, Sr and Ba: M(s) + 2H 2 O(l)  M(OH) 2 (aq) + H 2 (g)

34. Prentice Hall © 2003Chapter 7 Group 2A: The Alkaline Earth Metals Text, P. 260

35. Prentice Hall © 2003Chapter 7 7.8: Group Trends for Selected Nonmetals Hydrogen Most often occurs as a colorless diatomic gas, H 2 High IE because there are no inner electrons It can either gain another electron to form the hydride ion, H , or lose its electron to become H + : 2Na(s) + H 2 (g)  2NaH(s) 2H 2 (g) + O 2 (g)  2H 2 O(g) H + is a proton The aqueous chemistry of hydrogen is dominated by H + (aq)

36. Prentice Hall © 2003Chapter 7 Group 16: The Oxygen Group As we move down the group the metallic character increases (O 2 is a gas, Te is a metalloid, Po is a metal) There are two important forms of oxygen: O 2 and ozone, O 3 Ozone can be prepared from oxygen: 3O 2 (g)  2O 3 (g)  H = +284.6 kJ. Ozone is pungent and toxic

37. Prentice Hall © 2003Chapter 7 Group 16: The Oxygen Group Text, P. 261

38. Prentice Hall © 2003Chapter 7 Oxygen (O 2 ) is a potent oxidizing agent since the O 2- ion has a noble gas configuration There are two oxidation states for oxygen: 2- and 1- Sulfur is another important member of this group Most common form of sulfur is yellow S 8 Sulfur tends to form S 2- in compounds (sulfides) Allotropes: different forms of the same element in the same state

39. Prentice Hall © 2003Chapter 7 Group 17: The Halogens The chemistry of the halogens is dominated by gaining an electron to form an anion: X 2 + 2e -  2X - Fluorine is one of the most reactive substances known All halogens consists of diatomic molecules, X 2

40. Prentice Hall © 2003Chapter 7 Group 17: The Halogens Text, P. 262

41. Prentice Hall © 2003Chapter 7 Chlorine is the most industrially useful halogen Produced by the electrolysis of brine (NaCl): 2NaCl(aq) + 2H 2 O(l)  2NaOH(aq) + H 2 (g) + Cl 2 (g) Hydrogen compounds of the halogens are all strong acids with the exception of HF

42. Prentice Hall © 2003Chapter 7 Group 18: The Noble Gases These are all nonmetals and monatomic They are notoriously unreactive because they have completely filled s and p sub-shells In 1962 the first compound of the noble gases was prepared: XeF 2, XeF 4, and XeF 6 Xe has a low enough IE to react with substances that readily remove electrons To date the only other noble gas compounds known are KrF 2 and HArF “Argon hydrofluoride”

43. Prentice Hall © 2003Chapter 7 Group 18: The Noble Gases Text, P. 263

44. Prentice Hall © 2003Chapter 7 End of Chapter 7: Periodic Properties of the Elements