Exam R Nov 18 10-11:30 am

1. Exam R Nov 1810-11:30 am Review session W Nov 17 11-noon CF 316

2. Basic Principle: electrons occupy lowest energy levels available Rules for Filling Orbitals Bottom-up (Aufbau’s principle) Fill orbitals singly before doubling up (Hund’s Rule) Paired electrons have opposite spin (Pauli exclusion principle)

3. Identify examples of the following principles: 1) Aufbau 2) Hund’s rule 3) Pauli exclusion

4. Electron spin & magnetism For the ground state oxygen atom: spdf configuration: orbital box notation: Paramagnetic: atoms with unpaired electrons that are weakly attracted to a magnet. Diamagnetic: atoms with paired electrons that are not attracted to a magnet.

5. Apparatus for measuring magnetic properties

6. Electron configuration for As

7. Note: Not written according to Aufbau, but grouping according to n

8. Phosphorus Symbol:P Atomic Number:15 Full Configuration:1s22s22p63s23p3 Valence Configuration:3s23p3 Shorthand Configuration:[Ne]3s23p3          Box Notation          2s 1s 2p 3s 3p

9. Electronic configuration of Br 1s2 2s22p6 3s23p63d10 4s24p5 [Ar]3d104s24p5 [Ar] = “noble gas core” [Ar]3d10 = “pseudo noble gas core” (electrons that tend not to react) Atom’s reactivity is determined by valence electrons valence e’s in Br:4s24p5 highest n electrons

10. Valence e- shells for transition metalsv.main group elements d orbitals not included in valence shell (pseudo noble gas cores) d orbitals sometimes included in valence shell

11. Rule-of-thumb for valence electrons Examples ●Sulfur: 1s22s22p63s23p4 or [Ne]3s23p4 valence electrons:3s23p4 ● Strontium: [Kr]5s2 valence electrons:5s2 ● Gallium: [Ar]4s23d104p1 valence electrons:4s24p1 ● Vanadium: [Ar]4s23d3 valence electrons:4s2or3d34s2 Identify all electrons at the highest principal quantum number (n) Use on exams, but recognize limitations Use Table 8.9 for online HW

12. Selenium’s valence electrons Written for increasing energy: Pseudo noble gas core includes: noble gas electron core d electrons (not very reactive)

13. Core and valence electrons in Germanium Written for increasing energy: Pseudo noble gas core includes: noble gas core d electrons

14. d-block: some exceptions to the Aufbau principle Fig. 8.9: Use this table for online homework

15. Orbital energy ladder f d n = 4 p d s p n = 3 s p n = 2 s n = 1 Energy s

16. Quantum numbers and orbital energiesEach electron in an atom has a unique set of quantum numbers to define it{ n, l, ml, ms } • n = principal quantum number = any non-zero positive integer • electron’s energy depends principally on this • l = azimuthal quantum number = 0, 1, 2, …n-1 • for orbitals of same n, l distinguishes different shapes (angular momentum) • ml = magnetic quantum number = -l …0…+l • for orbitals of same n & l, ml distinguishes different orientations in space • ms = spin quantum number = +/- 1/2 • for orbitals of same n,l & ml, ms identifies the two possible spin orientations

17. Energy levelSublevel# of orbitals/sublevel n = 1 1s (l = 0) 1 (ml has one value) n = 22s (l = 0) 1 (ml has one value) 2p (l = 1) 3 (ml has three values) n = 33s (l = 0) 1 (ml has one value) 3p (l = 1) 3 (ml has three values) 3d (l = 2) 5 (ml has five values) Quantum numbers and orbital energies Each atom’s electron has a unique set of quantum numbers to define it{ n, l, ml, ms } n = principal quantum number (energy) l = azimuthal quantum number (shape) ml = magnetic quantum number (orientation)

18. Concept: Each electron in an atom has a unique set of quantum numbers to define it{ n, l, ml, ms } 18

19. Quantum numbers: unique set for each e- • s orbitals p orbitals d orbitals f orbitals • l = 0 l = 1l = 2l = 3 • ml = 0ml = -1, 0, 1ml = -2, -1, 0, 1, 2 ml=-3,-2,-1,0,1,2,3 • An s subshellA p subshellA d subshell An f subshell • One s orbitalThree p orbitalsFive d orbitals Seven f orbitals • For n=1 l=0 an s subshell (with 1 orbital) • For n=2 l=0,1 an s subshell and a p subshell (with 3 orbitals) • For n=3 l=0,1,2 an s subshell, a p subshell, a d subshell (with 5 orbitals) • For n=4 l=0,1,2,3 an s subshell, a p subshell, a d subshell, an f subshell (with 7 orbitals)

20. Chapter 8. Part B.The periodic table & the quantum model within a period Trends within a group

21. Use the Periodic Table to write the electron configuration and orbital diagram for the following: Na (atomic no. 11) Te (atomic no. 52) Tc (atomic no. 43) [Ne]3s1 [Kr]5s24d105p4 [Kr]5s24d5 3s 5s 5p 4d 5s 4d

22. Coulomb’s law of electrostatics Coulomb’s Law q1 = “-” charge of particle 1 q2 = “+” charge of particle 2 d = distance between particles Key Concept Use this to understand the attraction between valence electrons & nucleus

23. Penetrating and Shielding probability distributions shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel are more shielded from the attractive force of the nucleus the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus the result is that the electrons in the 2s sublevel are lower in energy than the electrons in the 2p

24. Consider these electron (e) probability distributions. Which e’s are most strongly attracted to the nucleus? Which e’s are most shielded?

25. The Periodic Law: Periodicity Higher “effective nuclear” charge Electrons held more tightly “Higher energy” Larger orbitals Electrons held less tightly Many chemical & physical properties vary periodically -density - reactivity - . . . . . 1. Atomic & ionic size 2. Ionization energy 3. Electron affinity 4. Metallic character 5. Electronegativity

26. Metals Properties of Metals • Lustrous • Malleable • Ductile • Good conductors of • electricity & heat Nonmetals Properties of Nonmetals • Dull appearance • Brittle solids • Poor conductors of • electricity & heat

27. Consider the Noble Gases….. They are chemically unreactive, but show systematic property change

28. Tables from N.J. Tro, “Chemistry”, Pearson, 2008. 1st ionization energy electron affinity

29. Trends when atoms form chemical bonds Metals tend to lose e-’s Nonmetals tend to gain e-’s Empirical Observation “when forming ionic compounds, elements tend to lose or gain electrons to be more like the nearest noble gas”

30. What does each category have in common? Noble Gases Alkali Metals Alkaline Earths Halogens

31. Reactivity of the alkali metals Trend? Lithium video 2Li(s) + 2H2O(l)  2LiOH(aq) + H2(g) Sodium video 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) Potassium video 2K(s) + 2H2O(l)  2KOH(aq) + H2(g)

32. More sodium reaction videos 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) http://www.theodoregray.com/PeriodicTable/ Prepping Na 100 g Na in one piece 150 g Na in small pieces

33. What explains the structure of the periodic table and the behavior and properties of the elements?

34. Specify if each pair has chemical properties that are similar () or not similar (X): 1. Cl and Br 2. P and S 3. O and S Concept Check

35. Trends in atomic radii●within a group?●within a period? Period 6 Period 5 Period 4 Period 3 Period 2

36. Atomic radii of main-group elements

37. Understanding “periodicity” Trends within a period what is changing? what is constant? Trends within a group what is changing? what is constant? 11 Na 12 Mg 11 Na 19 K

38. Screening effect of inner electrons Consider Mg [Ne]3s2 Core electrons of [Ne] “shield” valence electrons

39. Periodicity – trends within a given period Trends within a period what is changing? what is constant? 11 Na 12 Mg 13 Al

40. Effective nuclear charge (Z*) Z* = nuclear charge experienced by the outermost electrons. Estimate Z* by  [ Z - (# of shielding electrons)] (Z = total number of electrons) Z* increases across a period