Chapter 7 Periodic Properties of the Elements

1. CHEMISTRYThe Central Science 9th Edition Chapter 7 Periodic Properties of the Elements Chapter 7

2. 7.1: Development of the Periodic Table • 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 Chapter 7

3. 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 Chapter 7

4. 7.2: Effective Nuclear Charge • 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 (Zeff) decreases • OR…as the distance from the nucleus increases, Sincreases andZeffdecreases Chapter 7

5. 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) Text, P. 240 Zeff = Z – S or Protons – Electrons The radial electron density is the probability of finding an electron at a given distance Chapter 7

6. Zeff increases across a period • Added electrons don’t increase shielding • Zeff 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 Chapter 7

7. 7.3: Sizes of Atoms and Ions • 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 Chapter 7

8. 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 Chapter 7

9. Text, P. 242

10. 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 Chapter 7

11. Text P. 244

12. 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 (O2-, F-, Na+, Mg2+, Al3+) • As nuclear charge increases in an isoelectronic series, the ions become smaller: • O2- > F- > Na+ > Mg2+ > Al3+ • 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 Chapter 7

13. 7.4: Ionization Energy • The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom: • Na(g)  Na+(g) + e- • The second ionization energy, I2, is the energy required to remove an electron from a gaseous ion: • Na+(g)  Na2+(g) + e- • The larger ionization energy, the more difficult it is to remove the electron Chapter 7

14. Chapter 7

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

16. 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 I1 Chapter 7

17. Chapter 7 Text, P. 248

18. 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 I1 (half-filled sublevel) • Transition metals and f-block elements show slow variations in I1 due to shielding Chapter 7

19. Text, P. 248

20. Text, P. 247

21. Electron Configuration of Ions • Cations: electrons removed from orbital with highest principle quantum number, n, first: • Li (1s2 2s1)  Li+ (1s2) • Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5) • Anions: electrons added to the orbital with highest n: • F (1s2 2s2 2p5)  F-(1s2 2s2 2p6) Chapter 7

22. 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) Chapter 7

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. 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 Chapter 7

26. 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 Chapter 7

27. 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 Chapter 7

28. 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) + 2H2O(l)  2MOH(aq) + H2(g) Chapter 7

29. Alkali metals produce different oxides when reacting with O2: • 4Li(s) + O2(g)  2Li2O(s) (oxide) • 2Na(s) + O2(g)  Na2O2(s) (peroxide) • K(s) + O2(g)  KO2(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 Chapter 7

30. Li line: 2p 2s transition Na line (589 nm): 3p 3s transition K line: 4p 4s transition Chapter 7

31. Text, P. 257 Chapter 7

32. 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  M2+ + 2e- • Mg(s) + Cl2(g)  MgCl2(s) • Increasing activity down the group: • Be does not react with water • Mg will only react with steam • Ca, Sr and Ba: • M(s) + 2H2O(l)  M(OH)2(aq) + H2(g) Chapter 7

33. Group 2A: The Alkaline Earth Metals Text, P. 260 Chapter 7

34. 7.8: Group Trends for Selected Nonmetals • Hydrogen • Most often occurs as a colorless diatomic gas, H2 • 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) + H2(g)  2NaH(s) • 2H2(g) + O2(g)  2H2O(g) • H+ is a proton • The aqueous chemistry of hydrogen is dominated by H+(aq) Chapter 7

35. Group 16: The Oxygen Group • As we move down the group the metallic character increases (O2 is a gas, Te is a metalloid, Po is a metal) • There are two important forms of oxygen: O2 and ozone, O3 • Ozone can be prepared from oxygen: • 3O2(g)  2O3(g) H = +284.6 kJ. • Ozone is pungent and toxic Chapter 7

36. Group 16: The Oxygen Group Text, P. 261 Chapter 7

37. Oxygen (O2) is a potent oxidizing agent since the O2- 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 S8 • Sulfur tends to form S2- in compounds (sulfides) • Allotropes: different forms of the same element in the same state Chapter 7

38. Group 17: The Halogens • The chemistry of the halogens is dominated by gaining an electron to form an anion: • X2 + 2e- 2X- • Fluorine is one of the most reactive substances known • All halogens consists of diatomic molecules, X2 Chapter 7

39. Group 17: The Halogens Text, P. 262 Chapter 7

40. Chlorine is the most industrially useful halogen • Produced by the electrolysis of brine (NaCl): • 2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g) • Hydrogen compounds of the halogens are all strong acids with the exception of HF Chapter 7

41. 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: XeF2, XeF4, and XeF6 • Xe has a low enough IE to react with substances that readily remove electrons • To date the only other noble gas compounds known are KrF2 and HArF • “Argon hydrofluoride” Chapter 7

42. Group 18: The Noble Gases Text, P. 263 Chapter 7