AP Chapter 7

1. AP Chapter 7 Periodic Properties of the Elements HW: 25 31 33 35 36 37 39 41 47 53 63 67 69 73

2. 7.1 – Development of the PT -Newlands = Law of Octaves – Noticed that when organized by atomic mass, similar properties appear every 8th element -Mendeleev – Father of the PT. Made the first accepted PT. Based on periodicity and atomic mass. Could predict missing elements. -Meyer – Did same work as Mendeleev -Rutherford – Gold Foil Experiment – Estimated # of positive charges in the nucleus -Moseley – Determined atomic #

3. Development of 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.

4. Element Discoveries: Page 256

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

6. Periodic Classifications -Representative Elements – s and p block. AKA “main group” elements -IUPAC numbers groups as 1 – 18 -American style uses IA-IIA and IIIA – VIIIA -Noble Gases – Group 18 -Transition Metals – d block (Groups 3 – 12) - # System - We use IUPAC, not American

7. 7.2 - Effective Nuclear ChargeIn 10th Grade we called this “Strength of the Nucleus” • 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. • Nuclear Charge = The amount of pull the electron is experiencing.

8. 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. (In same period, Z increases across the PT and # shielding remains the same, so Zeff INCREASES across the PT) (Down a group, the valence e-s are farther away, so Zeff decreases down a group) Mg = 12-10 = 2

9. 7.3 - Atomic Radius The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

10. Atomic Radius 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

11. 7.3 - Ionic Radius • Ionic size depends upon: • Nuclear charge. • Number of electrons. • Orbitals in which electrons reside.

12. Ionic Radius • Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions are reduced.

13. Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions are increased.

14. Sizes of Ions • Ions increase in size as you go down a column. • Due to increasing value of n.

15. Ionic Radius • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing atomic #. =[He] =[Ne]

16. 7.4 - Ionization Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. Always requires energy (+). • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc.

17. Ionization Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

18. Trends in First Ionization Energies • Down a group, less energy is required to remove the first electron (lower IE). • For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

19. Trends in First Ionization Energies • Generally, across a period, it gets harder to remove an electron (higher IE). • As you go from left to right, Zeff increases.

20. Trends in First Ionization Energies However, there are two apparent exceptions to this trend.

21. Trends in First Ionization Energies • The first occurs between Groups 2 and 13. • Group 13 – The electron removed from p-orbital. Group 2 it is removed from an s-orbital • Electron farther from nucleus, so easier to remove • s electrons are stable as a pair

22. Trends in First Ionization Energies • The second occurs between Groups 15 and 16. • Group 15 has one e- in each orbital of the p subshell. • Group 16 has a pair and two singles in the p-subshell. Repulsion between the electrons in the same orbital makes it easier to remove.

23. Electron Affinity Energy change accompanying addition of electron to gaseous atom: Cl + e− Cl− -If EA is + = the process is endothermic and energy must be added to make it occur -If EA is - = the process is exothermic and energy is released when the process occurs (this is different than what your book says!)

24. Trends in Electron Affinity -In general, electron affinity becomes more exothermic (easier to add an e-) as you go from left to right across a period. -Becomes more endothermic as you go down a group.

25. Trends in Electron Affinity There are two discontinuities in this trend.

26. Trends in Electron Affinity • The first occurs between Groups 1 and 2. • For Group 2, the added electron must go into a p-orbital, not s-orbital. • Electron is farther from nucleus and feels repulsion from s-electrons, so requires more energy than adding to Group 1 atom.

27. Trends in Electron Affinity • The second occurs between Groups 14 and 15. • Group 15 has all half-filled p orbitals in the p-subshell. • Extra electron must go in to form a pair, which has repulsion, which takes more energy, so overall the process is less exothermic than it was for Group 14.

28. Valence Electrons and Configurations Valence Electrons – Outer shell electrons. Involved in bonding. Configurations of ions – Gain (anions) or Lose (cations) to obtain the configuration of a noble gas. Ex: Na F (Isoelectronic = same number and location of the electrons)

29. 7.6 - METALLIC PROPERTIES

30. Metals versus Nonmetals

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

32. Metals Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

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

34. 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.

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

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

37. 7.7 – Group Trends and the Variations in Chemical Properties • (sorry about the grammar…) The jump from Period 2 to 3 is “more different” than the rest of the group trends because Period 1 and 2 elements do NOT have the option for d orbitals • Ex: Li more different than Na, than Na is to K • Diagonal relationships: There is a similarity between elements in different group and period (ex: Li and Mg are similar; Be and Al; B and Si, etc.) • Groups 13 – 16 have more variation in properties because of the “steps” separating metals and nonmetals

38. Group 1 Hydrogen: Can be +1 or -1; Forms both molecular (water) and ionic substances (NaH) Alkali Metals: • Soft, metallic solids. • Name comes from Arabic word for ashes.

39. Alkali Metals • Found only as compounds in nature. • Have low densities and melting points. • Also have low ionization energies.

40. Alkali Metals Their reactions with water are famously exothermic. -Form the hydroxide and hydrogen gas

41. Alkali Metals • Alkali metals (except Li) react with oxygen to form peroxides. • K, Rb, and Cs also form superoxides: K + O2 KO2 • Produce bright colors when placed in flame.

42. Group 2: Alkaline Earth Metals • Less reactive than Group 1 • Have higher densities and melting points than alkali metals. • Have low ionization energies, but not as low as alkali metals.

43. Alkaline Earth Metals • Bedoes not react with water, Mg reacts only with steam, but others react readily with water. • Reactivity tends to increase as go down group.

44. Group 13 • B is a metalloid. Does not form binary ionic compounds. • Al, Ga, In, Tl are metals • Al forms +3 ions • Ga, In, Tl form +1 and +3 ions (lose p or lose s and p valence e-)

45. Group 14 C = Nonmetal Si, Ge = Metalloid Sn, Pb = Metal -All can be +2 or +4 oxidation states -Sn and Pb more stable at +2; C and Si more stable at +4 -C can be -4 oxidation state

46. Group 15 • N, P = nonmetals ex: N2,NO, N2O, NO2 P4, P4O6, P4O10 • As, Sb = metalloids • Bi = metal

47. Group 16 • Oxygen, sulfur, and selenium are nonmetals. • Tellurium is a metalloid. • The radioactive polonium is a metal.

48. Oxygen • Two allotropes: • O2 • O3, ozone • Three anions: • O2−, oxide • O22−, peroxide • O21−, superoxide • Tends to take electrons from other elements (oxidation)

49. Sulfur • Weaker oxidizing agent than oxygen. • Most stable allotrope is S8, a ringed molecule.

50. Group 17: Halogens • Prototypical nonmetals • Name comes from the Greek halos and gennao: “salt formers”