1. Chapter 7; Electronic Structure of Atoms I.Electromagnetic Radiation II.Flame Test/ Emission Spectra III.Quantized Energy Levels IV.Bohr Model/ Rydberg Equation V.Principal Energy Levels, n a)First Ionization Energy b)2 nd, 3 rd, 4 th, etc Ionization Energy

2. Chapter 7; Electronic Structure of Atoms VI.Sublevels (s, p, d, f) a)Photoelectron Spectroscopy VII.Electron Configuration VIII. Valence Electrons/ Core IX.Good/ Bad Point of Atom Model X.Quantum Theory a)Dual Nature of the Electron b)Heisenberg Uncertainty Principle

3. Chapter 7; Electronic Structure of Atoms XI.Quantum Numbers (n, l, m l, m s ) XII.Oribtal Diagrams a)Paramagnetism and Diamagnetism

4. Electronic Structure Model Experimental Evidence 1.Line Spectra 2.Ionization Energies 3.Photoelectron Spectrum 4.Intensity/detail of Line Spectra What it means 1.Electrons in quanitized ‘n’ 2.# electrons in each ‘n’ 3.# electrons in each ‘n’ and each sublevel 4.Indicates ‘n’ have sublevels associated with them

5. Electronic Structure n# of Sublevel # e - in n (2n 2 ) Sublevel Names # e - in each sublevel 112ss-2 228s,ps-2, p-6 3318s,p,ds-2, p-6, d-10 4432s,p,d,fs-2, p-6, d-10, f-14

7. Order of orbitals (filling) in multi-electron atom 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s 7.7

9. What is the electron configuration of Mg? Mg 12 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s 2 2s 2 2p 6 3s 2 2 + 2 + 6 + 2 = 12 electrons 7.7 Abbreviated as [Ne]3s 2 [Ne] 1s 2 2s 2 2p 6 What is the electron configuration of Cl? Cl 17 electrons1s < 2s < 2p < 3s < 3p < 4s 1s 2 2s 2 2p 6 3s 2 3p 5 2 + 2 + 6 + 2 + 5 = 17 electrons

10. Electron Configurations of Cations and Anions Na [Ne]3s 1 Na + [Ne] Ca [Ar]4s 2 Ca 2+ [Ar] Al [Ne]3s 2 3p 1 Al 3+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. H 1s 1 H - 1s 2 or [He] F 1s 2 2s 2 2p 5 F - 1s 2 2s 2 2p 6 or [Ne] O 1s 2 2s 2 2p 4 O 2- 1s 2 2s 2 2p 6 or [Ne] N 1s 2 2s 2 2p 3 N 3- 1s 2 2s 2 2p 6 or [Ne] Atoms gain electrons so that anion has a noble-gas outer electron configuration. Of Representative Elements 8.2

11. Na + : [Ne]Al 3+ : [Ne]F - : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne]N 3- : 1s 2 2s 2 2p 6 or [Ne] Na +, Al 3+, F -, O 2-, and N 3- are all isoelectronic with Ne What neutral atom is isoelectronic with H - ? H - : 1s 2 same electron configuration as He 8.2

12. Electron Configurations of Transition Metals Completely filled or half-completely filled d-orbitals have a special stability –Some “irregularities” are seen in the electron configurations of transition and inner-transition metals.

13. Electron Configurations of Cations of Transition Metals 8.2 When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s 2 3d 6 Fe 2+ : [Ar]4s 0 3d 6 or [Ar]3d 6 Fe 3+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Mn: [Ar]4s 2 3d 5 Mn 2+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Order of filling; 3s<3p<4s<3d But when removing electrons to form + ions for transition metals Order of removing electrons; 4s<3d<3p<3s

14. Electronic Structure Good Points Electrons in Quantized Energy Levels Maximum # electrons in each n is 2n 2 Sublevels (s,p,d,f) and # electrons they hold Bad Points Electrons are placed in orbits about nucleus Only explains emission spectra of H 2 Does not address all interactions Treats electron as particle

15. H +1 Be +4 There are less interactions to take into account in H than other elements Interactions 1.Attraction between + nucleus and negative electrons Interactions 1.Attraction between + nucleus and negative electrons 2.Repulsion between electrons in same energy level. 3.Shielding effect of filled principal energy levels.

16. Quantum Theory – Revised Electronic Structure Model 1.Dual Nature of the Electron 2.Heisenberg Uncertainty Principle

17. Dual Nature of Electron Previous Concept; A Substance is Either Matter or Energy Matter; Definite Mass and Position Made of Particles Energy; Massless and Delocalized Position not Specificed Wave-like

18. Dual Nature of Electron Electron is both “particle-like” and “wave- like” at the same time. Previous model only considered “particle- like” nature of the electron

23. Heisenberg Uncertainty Principle Act of measuring the position and energy of electron changes the position of electron –Better one variable is known (energy); the less well the other variable is known (position)

24. Orbitals Replace Orbits Orbits- Both electron position and energy known with certainty Orbitals – Regions of space where an electrons of a given energy will most likely be found

25. Quantum Theory Orbitals Replace Orbits Orbits Orbitals

26. Schrodinger Wave Equation (  ) Describes size/shape/orientation of orbitals 7.5 Wave Equation is based on… 1.Dual Nature of Electron (Electron both particle and wave-like at the same time.) 2.Heisenberg Uncertainty Principle (Orbitals describe a region in space an electron will most likely be.)

27. Wave Equation (  ) Wave Equation describe the size, shape, and orientation of the orbital the electron (of a given energy) is in. There are four variables in the function -n; Energy and size of orbital –l; Shape of orbital –m l ; Orientation of orbital –m s ; Electron Spin (n, l, m l, m s )

28. 1.Each electron has a unique set of 4 Quantum Numbers 2.Each orbital described by the Quantum Numbers can hold a maximum of 2 electrons.

29. Schrodinger Wave Equation; 1 st Quantum Number  fn(n, l, m l, m s ) principal quantum number n n = 1, 2, 3, 4, …. n=1 n=2 n=3 7.6 distance of e - from the nucleus

30.  = fn(n, l, m l, m s ) angular momentum quantum number l for a given value of n, l = 0, 1, 2, 3, … n-1 n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 Shape of the “volume” of space that the e - occupies l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital Schrodinger Wave Equation 2 nd Quantum Number 7.6

31. Principal Energy Level, n Sublevel, l Quantum #Electron Configuration 10(1,0,, )1s 20(2,0,, )2s 1(2, 1,, )2p 30(3,0,, )3s 1(3, 1,, )3p 2(3,2,, )3d 40(4,0,, )4s 1(4, 1,, )4p 2(4, 2,, )4d 3(4, 3,, )4f

32. l = 0 (s orbitals) l = 1 (p orbitals) 7.6

33. l = 2 (d orbitals)

34. f-orbitals

35. Orbital Shapes Orbital TypeShape Name sSpherical pDumbbell dComplex fMore complex

36.  = fn(n, l, m l, m s ) magnetic quantum number m l for a given value of l m l = -l, …., 0, …. +l orientation of the orbital in space if l = 1 (p orbital), m l = -1, 0, or 1 if l = 2 (d orbital), m l = -2, -1, 0, 1, or 2 Schrodinger Wave Equation 3 rd Quantum Number 7.6

37. Number of Degenerate Orbitals Needed for Each Type of Orbital (Sublevel) Type of OrbitalMaximum # of electrons in Orbital # of Degenerate Orbitals s21 p63 d105 f147

38. m l = -1m l = 0m l = 1 m l = -2m l = -1m l = 0m l = 1m l = 2

39.  = fn(n, l, m l, m s ) spin quantum number m s m s = +½ or -½ Schrodinger Wave Equation 4 th Quantum Number m s = -½m s = +½ 7.6

40. Valid Possibilities for Quantum Numbers Chemistry; The Science in Context; by Thomas R Gilbert, Rein V Kriss, and Geoffrey Davies, Norton Publisher, 2004, p125

42. How many 2p orbitals are there in an atom? 2p n=2 l = 1 If l = 1, then m l = -1, 0, or +1 3 orbitals How many electrons can be placed in the 3d subshell? 3d n=3 l = 2 If l = 2, then m l = -2, -1, 0, +1, or +2 5 orbitals which can hold a total of 10 e - 7.6

44. Three Manners to Convey How Electrons are Arranged 1.Electron Configuration ; List Orbitals and Number of Electrons in Each (1s 2 2s 2 2p 6 3s 2 …) 2.Quantum Numbers (2,0,0,+1/2) 3.Orbital Diagrams; List Orbitals and show location of electrons and their spin

45. Pauli exclusion principle - no two electrons in an atom can have the same four quantum numbers. The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule) or maximum # of unpaired electrons. Orbital Diagrams

46. 7.7 Orbital Diagrams Carbon; 6 electrons Electron Configuration; 1s 2 2s 2 2p 2 Orbital Diagram

47. Orbital Diagrams Oxygen; 8 electrons Electron Configuration; 1s 2 2s 2 2p 4 Orbital Diagram

48. Paramagnetic unpaired electrons 2p Diamagnetic all electrons paired 2p 7.8