Chapter 7; Electronic Structure of Atoms

1. Chapter 7; Electronic Structure of Atoms • Electromagnetic Radiation • Flame Test/ Emission Spectra • Quantized Energy Levels • Bohr Model/ Rydberg Equation • Principal Energy Levels, n • First Ionization Energy • 2nd , 3rd, 4th, etc Ionization Energy

2. Chapter 7; Electronic Structureof Atoms • Sublevels (s, p, d, f) • Photoelectron Spectroscopy • Electron Configuration • Valence Electrons/ Core • Good/ Bad Point of Atom Model • Quantum Theory • Dual Nature of the Electron • Heisenberg Uncertainty Principle

3. Chapter 7; ElectronicStructure of Atoms • Quantum Numbers (n, l, ml, ms) • Oribtal Diagrams • Paramagnetism and Diamagnetism

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

5. Electronic Structure

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? What is the electron configuration of Cl? Mg 12 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s22s22p63s2 2 + 2 + 6 + 2 = 12 electrons Abbreviated as [Ne]3s2 [Ne] 1s22s22p6 Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s22s22p63s23p5 2 + 2 + 6 + 2 + 5 = 17 electrons 7.7

10. Electron Configurations of Cations and Anions Of Representative Elements Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne] 8.2

11. What neutral atom is isoelectronic with H- ? Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne H-: 1s2 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. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5 Electron Configurations of Cations of Transition Metals 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. 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 8.2

14. Good Points Electrons in Quantized Energy Levels Maximum # electrons in each n is 2n2 Sublevels (s,p,d,f) and # electrons they hold Bad Points Electrons are placed in orbits about nucleus Only explains emission spectra of H2 Does not address all interactions Treats electron as particle Electronic Structure

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

16. Quantum Theory – Revised Electronic Structure Model • Dual Nature of the Electron • 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 TheoryOrbitals Replace Orbits Orbitals Orbits

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

27. Wave Equation (Y) • 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 • ml; Orientation of orbital • ms; Electron Spin (n, l, ml, ms)

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

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

30. Schrodinger Wave Equation 2nd Quantum Number Y = fn(n, l, ml, ms) angular momentum quantum number l for a given value of n, l= 0, 1, 2, 3, … n-1 l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital 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 7.6

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

33. l = 2 (d orbitals)

34. f-orbitals

35. Orbital Shapes

36. Schrodinger Wave Equation 3rd Quantum Number Y = fn(n, l, ml, ms) magnetic quantum number ml for a given value of l ml = -l, …., 0, …. +l if l = 1 (p orbital), ml = -1, 0, or1 if l = 2 (d orbital), ml = -2, -1, 0, 1, or2 orientation of the orbital in space 7.6

37. Number of Degenerate Orbitals Needed for Each Type of Orbital (Sublevel)

38. ml = -2 ml = -1 ml = 0 ml = 1 ml = 2 ml = -1 ml = 0 ml = 1

39. Schrodinger Wave Equation 4th Quantum Number Y = fn(n, l, ml, ms) spin quantum number ms ms = +½or -½ ms = +½ ms = -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 electrons can be placed in the 3d subshell? n=2 n=3 l = 1 l = 2 How many 2p orbitals are there in an atom? If l = 1, then ml = -1, 0, or +1 2p 3 orbitals If l = 2, then ml = -2, -1, 0, +1, or +2 3d 5 orbitals which can hold a total of 10 e- 7.6

44. Three Manners to Convey How Electrons are Arranged • Electron Configuration ; List Orbitals and Number of Electrons in Each (1s22s22p63s2…) • Quantum Numbers (2,0,0,+1/2) • Orbital Diagrams; List Orbitals and show location of electrons and their spin

45. 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 Pauli exclusion principle - no two electrons in an atom can have the same four quantum numbers.

46. Orbital Diagrams Carbon; 6 electrons Electron Configuration; 1s22s22p2 Orbital Diagram 7.7

47. Orbital Diagrams Oxygen; 8 electrons Electron Configuration; 1s22s22p4 Orbital Diagram

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