Quantum numbers and orbital energies Each atom’s electron has a unique set of quantum numbers to define it { n , l ,

1. 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)

2. Concept: Each electron in an atom has a unique set of quantum numbers to define it{ n, l, ml, ms } 2

3. Fig. 7.14

4. What is the reason that the periodic table organizes elements according to similarities in chemical properties?

5. There is a relationship between the quantum number (n) and its the number of subshells. Arrangement of electrons in atoms Principal quantum number (n) =number of subshells

6. multi-electron atoms H atom All other atoms As n increases, the difference in energy level _______. 6

7. Energy levels for multi-electron atoms 2n2 electrons

8. Orbital energy ladder n = 3  18 e’s n = 2  8 e’s n = 1  2 e’s

9. Basic Principle: electrons occupy lowest energy levels available

12. Aufbau Principle -- “Bottom Up Rule”

13. Stern-Gerlach Experiment   How could an orbital hold two electrons without electrostatic repulsion? Electron spin

14. 2 ways to write electron configurations spdf Notation spdf NOTATION for H, atomic number = 1 1 no. of electrons s 1 sublevel value of energy level Orbital Box Notation ORBITAL BOX NOTATION for He, atomic number = 2 Arrows show electron spin (+½ or -½) 2  1s 1s

15. Example: Determine the electron configuration and orbital notation for the ground state neon atom. Pauli exclusion principle An orbital can contain a maximum of 2 electrons, and they must have the opposite “spin.”

16. Write the ground state configuration and the orbital diagram for oxygen in its ground state Hund’s Rule -

17. 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)

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

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

20. 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.

21. Apparatus for measuring magnetic properties

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

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

24. Electron distribution for the argon atom Never zero electron distribution

25. Electron configuration for As

26. Silicon’s valence electrons 1s22s22p63s23p2 [Ne]3s23p2  [Ne] Shorthand notation for silicon

27. Shorthand notation practice Examples ●Aluminum: 1s22s22p63s23p1[Ne]3s23p1 ● Calcium: 1s22s22p63s23p64s2 [Ar]4s2 ● Nickel: 1s22s22p63s23p64s23d8 [Ar]4s23d8 {or [Ar]3d84s2} ● Iodine: [Kr]5s24d105p5 {or [Kr]4d105s25p5} ● Astatine (At): [Xe]6s24f145d106p5 {or [Xe]4f145d106s26p5} [Noble Gas Core] + higher energy electrons

28. Outer electron configuration for the elements

29. 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

30. Valence e’s for “main group” elements

31. 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

32. 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)

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

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

35. d-block: some exceptions to the Aufbau principle

36. Mendeleev’s periodic table generally organized elements by increasing atomic mass and with similar properties in columns. In some places, there were missing elements whose properties he predicted. When gallium, scandium, and germanium were isolated and characterized, their properties were almost identical to those predicted by Mendeleev for eka-aluminum, eka-boron, and eka-silicon, respectively.

37. Figure 8.14: Mendeleev’s periodic table.

38. Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. We will look in more detail at three periodic properties: atomic radius, ionization energy, and electron affinity.

39. Effective Nuclear Charge Effective nuclear charge is the positive charge that an electron experiences from the nucleus. It is equal to the nuclear charge, but is reduced by shielding or screening from any intervening electron distribution (inner shell electrons).

40. Effective nuclear charge increases across a period. Because the shell number (n) is the same across a period, each successive atom experiences a stronger nuclear charge. As a result, the atomic size decreases across a period.

41. Atomic Radius While an atom does not have a definite size, we can define it in terms of covalent radii (the radius in covalent compounds).

43. Figure 8.17: Representation of atomic radii (covalent radii) of the main-group elements.

44. Atomic radius is plotted against atomic number in the graph below. Note the regular (periodic) variation.

45. A representation of atomic radii is shown below.

46. Refer to a periodic table and arrange the following elements in order of increasing atomic radius: Br, Se, Te. 34 Se 35 Br 52 Te Te is larger than Se. Se is larger than Br. Br < Se < Te

47. First Ionization Energy (first ionization potential) The minimum energy needed to remove the highest-energy (outermost) electron from a neutral atom in the gaseous state, thereby forming a positive ion

48. Periodicity of First Ionization Energy (IE1) Like Figure 8-18

49. Fig. 8.15

50. Left of the line, valence shell electrons are being removed. Right of the line, noble-gas core electrons are being removed.