Unit 8: Periodic Properties of the Elements

1. CHM 1045: General Chemistry and Qualitative Analysis Unit 8:Periodic Propertiesof the Elements Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL • Textbook Reference: • Chapter #6 • Module #10

2. Development of Periodic Table Dmitry Ivanovich Mendeleyev (1834-1907) noticed a repeating pattern of properties and reactivities, when the elements are ordered by increasing atomic numbers.

3. Development of Periodic Table

4. Development of Periodic Table Mendeleyev predicted the discovery and properties of Sc, Ga, Ge

5. Development of Periodic Table *

6. {Intro.Per.Prop*} Periodic Trends • Size of atoms and ions: Determined by electrons which are both attracted to the nucleus and repelled by other electrons. • Ionization energy:energy required to remove an electron from and atom or ion. (Measure of how strongly an atom holds its own electrons) • Electron affinity:energy released when an atoms accept electrons. (Measure of how strongly an atom wants other electrons) • Metal, Non-metals, Metalloids • Group Trends: reactive characteristicsdetermined by valence electrons.

7. Effective Nuclear Charge (Zeff ) Felt by Shielding electrons (S) Effective Nuclear Charge (Zeff) Atomic Number (Z = #p+) Zeff = Z - S Inner Core of Shielding e- = screening constant Attractive Charge felt by valence electrons Atomic Number (#p+) {Eff.Nuclear.Charge*}

8. Effective Nuclear Charge (Zeff ) depends on both the fact that valence electrons are both: • attracted to the nucleus • repelled by the other (shielding) electrons. Effective Nuclear Charge (Zeff.) Increasing +1 Shielding electrons +8 (0e-) +2 +3 +4 +5 +6 +7 +2 (2e-) (10e-) (18e-) (19e-) (20e-) (21e-) (22e-) (23e-) (24e-) (25e-) (26e-) (27e-) ( 28e-)

9. (1) Sizes of Atoms(Atomic Radii) increasing value of n 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 increasing Zeff. {AtomicRadii*}

10. Sizes of Atoms (Atomic Radii)

15. Sizes of Atoms and Ions • Experimental Techniques for determining radii: • X-ray crystallography • Neutron scattering

16. Sizes of Atoms The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

17. {SizeOfIons*} Sizes of Ions • Anions:larger than parent atoms. • Electrons are added and repulsions are increased. {Grey = atom,Blue = anion} {Pink = cation, Grey = atom} • Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions are reduced.

18. Sizes of Ions {Pink = cation, Grey = atom , Blue = anion} • Ions increase in size as you go down a column. • Due to increasing value of n. Anions larger Cations larger

19. Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge. {Pink = cation, Grey = atom , Blue = anion}

20. (2) Ionization Energy • the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • A good measure of how strongly metal ions hold-on to their electrons as they become cations: M  M+ + e- Electron gun

21. {IonizationEnergy} Ionization Energy • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc. • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap (large jump).

22. Trends in First Ionization Energies • As one goes up in a column, more energy is required to remove the first electron. • Zeff is essentially the same, but the valence electrons are closer to the nucleus. +1 +8 Zeff +2 +5 +3 +4 +6 +7

23. Trends in First Ionization Energies • As one goes across a row, the Zeff increases, so more energy is required to remove the first electron. • Zeff increases, but the valence electrons essentially in the same distance (energy level) from the nucleus. +1 +8 Zeff +2 +5 +7 +4 +6 +3

24. Trends in First Ionization Energies • Generally, across a row, the Zeff increases. {IonizationEnergy}

25. Trends in First Ionization Energies there are two apparent discontinuities in this trend.

26. Trends in First Ionization Energies • 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.

27. Trends in First Ionization Energies • 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.

28. (3) Electron Affinity Energy change accompanying addition of electron to gaseous atom: Cl + e− Cl− A good measure of how strongly non-metal atom acquire electrons to become anions: N + e- N- {ElectronAffinity}

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

30. Trends in Electron Affinity There are again, however, two discontinuities in this trend.

31. Trends in Electron Affinity • 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.

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

33. (4) Properties of Metal, Nonmetals,and Metalloids

34. Metals versus Nonmetals • Metals tend to form cations. • Nonmetals tend to form anions.

35. Metals vs. Nonmetals Metals tend to be: • lustrous • Malleable and ductile • Good conductors of heat and electricity. • Loose electrons (cations) in reactions with nonmetals in order to acquire noble gas configuration • Unreactive with each, instead they mix with each other to form alloys (solutions) • Nonmetals tend to be: • Dull • Brittle • Poor conductors of heat and electricity. • Gain electrons (anions) in reactions with metals in order to acquire noble gas configuration. • Reactive with each otherto form molecular (covalent) compounds.

36. Metals do not react with each otherto form compounds, instead they mix with each other to form alloys (solutions) Gold Gold Jewelry 24k =100% Au; 12k =50% Au, usually mixed with Ag, Pd or Ni (white) or Cu (rose); colored alloys coated with Rh or Pt to prevent rusting. 24k =100% Au 12k = 50%    mixed with Ag, Pd mixed with Cu mixed with Ni {coated with Rh or Pt to prevent rusting}

37. Metal oxides vs. Nonmetal oxides H2SO3(aq) weak acid H2SO4(aq) strong acid Note that the more oxidized the oxide the stronger its acid base strength! Nonmetal oxides tend to be acidic. SO2 + H2O  SO3 + H2O  CO2(s) CO2 SO3 MgO Na2O Amphoteric: capable of behaving as an acid or a base. Acid H2O + bromthymol blue Indicator Basic H2O + bromthymol blue Indicator Metal oxides tend to be basic. CaO + HOH  Ca2+ + 2OH- CO2(s) + H2O(l)  H2CO3(aq) weak acid H2CO3(aq)  CO2(g) + H2O(l) {Acid/Base behavior of Oxides}

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

39. Not a Metalloid: Silicone • An organic compound made primarily from silicon and oxygen Silicone overdose

41. Alkali Metals Nobel Gases Alkaline Earth Metals Halogens (5) Group Trends

42. Alkali Metals • Soft, metallic solids. • Name comes from Arabic word for ashes. • So reactive, that they are found only as compounds in nature. • Have low densities and melting points. • Also have low ionization energies.

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

44. Alkali Metals • Alkali metals react with oxygen to form oxides (O2-) • 4K + O2 2K2O • Alkali metals (except Li) react with oxygen to form peroxides (O22-) 2K + O2 K2O2 • K, Rb, and Cs also form superoxides (O2-): K + O2 KO2 • Produce bright colors when placed in flame. . 2 {ox # = -1} {ox # = -½}

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

46. Group 4A: The Carbon Family • Allotropes of Carbon: • a) Diamond • b) Graphite • c) Lonsdaleite • d) C60 (Buckminsterfullerene) • e) C540 (see Fullerene) • f) C70 (see Fullerene) • g) Amorphous carbon • h) single-walled carbon nanotube

47. Group 6A: The Oxygen Family • Oxygen, sulfur, and selenium are nonmetals. • Tellurium is a metalloid. • The radioactive polonium is a metal.

48. Oxygen O O • Two allotropes: • O2 • O3, ozone • Oxidizing agent: Tends to take electrons from other elements • Three anions: • O2−, oxide • O22−, peroxide • O21−, superoxide Sulfur • Weaker oxidizing agent than oxygen. • Most stable allotrope is S8, a ringed molecule.

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

50. Group VIIA: Halogens • 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