22 November 2011

1. 22 November 2011 • Take out your Problem Set (if you haven’t already handed it in) • Objective: You will be able to: • describe, calculate and compare the effective nuclear charge of elements

2. Agenda • Questions about electron configuration? • Effective nuclear charge notes and problems Homework: Quiz Monday 6-8 multiple choice, 2 multi-part free response

3. Periodic Properties of the Elements

4. Today • attraction between electrons and the nucleus • repulsion between electrons • and some properties and their trends on the periodic table that this attraction/repulsion causes

5. Effective Nuclear Charge • First, some definitions: • force of attraction (between two charged particles) where Q1 and Q2 are the charges of the particles and d is the distance between them • in general, electrons close to the nucleus will be held with greater force than those that are more distant from the nucleus! • higher positive nuclear charges will draw electrons closer to the nucleus and hold them tighter.

6. Valence vs. Core Electrons • Valence electrons: electrons in the outermost orbitals of atoms, farthest from the nucleus • Core electrons: inner electrons, include electrons in completely full “d” orbitals

7. Effective Nuclear Charge • http://www.youtube.com/watch?v=MtP5mWLB-ys • Zeff, the net positive charge experienced by an electron in an atom. • Not the full nuclear charge because the core electrons “shield” (cancel) part of the positive nuclear charge. • Valence electrons experience less-than-full pull from the nuclear charge.

8. Shielding Effect • the reduction of the full nuclear charge experienced by an outer electron as a result of screening (cancelling) by inner core electrons

9. Trends in Effective Nuclear Charge • Zeffincreases from left to right across any period • core electrons maintain constant across any row, but the nuclear charge increases, so Zeff increases • Zeff = Z – S • S: screening constant, about equal to the number of core electrons in an atom

10. Examples

11. Zeffis constant going down a group because valence electrons going down a group are constant, but there is an increase in number of protons to balance this

13. Problems • What is the approximate Zeff of scandium? • Are the valence electrons of Sc held more or less tightly than those of K? Use scandium’s electron configuration to explain your answer.

14. Zeffproduces trends… • Sizes of atoms and ions • Ionization energy • Electron affinity (electronegativity)

15. 29 November 2011 • Objective: You will be able to: describe trends on the periodic table caused by effective nuclear charge. • Do now: (On page 5, 3rd slide): • How many energy levels do the following atoms have? • sodium b. potassium c. rubidium • Calculate Zeff for Na, Mg and Al.

16. Sizes of Atoms and Ions • atomic radius: an estimate of the size of an atom • atoms don’t have sharply defined boundaries because orbitals are areas of probability, so definite sizes can’t be determined

17. Atomic Radius (size of the atom) Atomic radius increases Atomic radius increases

19. Atomic Radius • Increases top to bottom: outer electrons are on higher energy levels, which are further from the nucleus • Decreases left to right: shielding remains constant as nuclear charge increases • no more core electrons are added, but more protons are, which pull the valence electrons closer to the nucleus

20. Ionic Radius • http://www.youtube.com/watch?v=hkyxQjKwBU4 • Cations (+) are smaller than their parent atoms because the electron is lost from the valence shell, and e--e- repulsionsare decreased • Anions (-) are larger than their parent atoms because additional electrons cause increased e--e- repulsions, causing the electrons to spread out more in space

22. Ionization Energy • first ionization energy (I1): the energy required to remove the outermost electron from the ground state of a gaseous atom. • Ex: 495 kJ + Na(g) → Na+(g) + e− • second ionization energy (I2): the energy required to remove the second electron • etc. • I1<I2<I3 because with each successive removal, an electron is pulled away from an increasingly positive ion.

23. http://www.youtube.com/watch?v=6e4uoWQeM4s&feature=related

24. Ionization Energy Increases Ionization Energy Increases

27. Exceptions • I1 decreases from Be to B and Mg to Al • electrons in filled s or d orbitals provide limited screening for electrons in p subshells • I1 decreases from N to O, P to S and As to Se • due to repulsion of paired electrons in the p4 configuration of group 16 atoms

30. Noble Gases • have the highest ionization energies of their periods because their valence electrons are poorly screened. • very high Zeff • They are also the smallest in their periods

31. Problem • Arrange the period 3 elements in order of increasing first ionization energy, lowest to highest. Note any anomalies.

32. Electron Affinity • a.k.a. electronegativity • ∆Hea, the energy change when an electron is added to a gaseous atom • F(g) + e- → F-(g) ∆Hea = -328 kJ/mol • Energy is released when an atom attracts an electron. • http://www.youtube.com/watch?v=scvNYZD3jrI

33. Electronegativity Increases Electronegativity Increases

34. Trend in Electron Affinity • Increases from left to right along a period • Increases from bottom to top within a group • smaller atoms are less shielded and attract electrons more easily • Exception: F has less electron affinity than chlorine because of the small size of F causes greater e--e- repulsion

35. Summary of Exceptions

36. 30 November 2011 • Objective: You will be able to describe and write chemical questions for patterns in reactivity on the periodic table. • Homework quiz (week of Nov. 28) • Compare the radius and ionization energy of oxygen and sulfur. Explain your answer.

37. Agenda • Homework Quiz • Homework answers • More trends on the periodic table: reactivity and compounds formed Homework: Read lab – be familiar with the procedure for tomorrow! p. 359 #61, 72, 73, 74, 82, 86, 95: Mon.

39. Metals, Non-Metals, Metaloids • Metals: low ionization energy, lose electrons readily • have luster, conduct heat and electricity, malleable, ductile • Metallic character increases right to left along a period and top to bottom within a group • Metal hydrides, oxides and nitrides are basic • Li2O(s) + H2O(l) → 2Li+(aq) + 2OH-(aq)

40. Non-metals have high electron affinity and gain electrons readily • form negative ions • do not have luster, and are poor conductors of heat and electricity • form molecular compounds • non-metal oxides are acidic • SO2(g) + H2O(l) → H2SO3(aq)

41. metalloids have properties intermediate between those of metals and non-metals

42. Trends for Group 1 and 2 Metals • Group 1 • alkali metals (group 1): soft, metallic solids • s1 valence electron configurations • lose one electron to form 1+ cations • become more reactive moving down the group • http://www.youtube.com/watch?v=uixxJtJPVXk

43. Alkali metal + Water • all alkali metals react with water to produce hydrogen gas 2Na(s) + 2H2O(l) → 2Na+(aq) + 2OH-(aq) + H2(g)

44. Alkali metal + H2 gas • all alkali metals react with hydrogen gas to form hydrides • 2Li(s) + H2(g) → 2LiH(s)

45. Alkali metals + Non-metals • all alkali metals react with most non-metals • 2K(s) + S(s) → K2S(s) • 6Li(s) + N2(g) → 2Li3N(s)

46. Forming Peroxides • Na, K, Rb and Cs form peroxides • 2Na(s) + O2(g) → Na2O2(s) oxide ion: O2- peroxide ion: O22- superoxide ion: O2-

47. Forming Superoxides • K, Rb and Cs form superoxides • K(s) + O2(g) → KO2(s)

48. When burned • Li: crimson-red • Na: yellow • K: violet

49. Practice • Write and balance a chemical equation to describe what happens when solid potassium is added to water. • Classify as acid-base, redox, or precipitation reaction. • Describe what you would observe when the reaction takes place.

50. 6 December 2011 • Objective: You will be able to: • Describe trends in alkaline earth metals, some non-metals, allotropes, halogens and noble gases • Do now: Find the final mass of your copper + filter paper.