Atomic structure part 2

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Shielding effects of electrons and consequences for atomic properties across the periodic table (atomic radii, electronegativities, ionization energy)

Transcript of Atomic structure part 2

  • 1. ATOMIC STRUCTURE (PART 2) Lesson by Dr.Chris UP, May 2014

2. WHAT WE WILL LEARN PART 1: 3. REVIEW LESSON 1 Atomic nucleus atomic mass unit amu Isotopes / MS Bohr model of the Hydrogen atom absorption and emission spectra energy levels in the atom = orbits Electron as standing wave deBroglie: = h/p = h/(mv) Schroedinger eq. for 3D el.waves => 3 quantum numbers n, l and m Aufbau principle Shielding of electrons 4. ATOMIC SPECTRA Show us the energy levels in an atom/molecle (n) and also splitting (l) 5. MAIN ENERGY LEVELS (N) Each line with a wavelength corresponds to an electron transition with an energy E = E2 E1 = h * c / 1/ = wavenumber in cm-1 Planck constant h = 6.626 x 10-34 J s C speed of light = 300000 km/s 6. Energy unit electronVolt eV = energy of 1 electron in a field of 1 V 1 eV => = 1240 nm = 1240 nm / E [eV] 7. CHECK Which light will an electron emit, when it falls from energy level 4 to 2 ? (is it shorter or longer than from 4 to 3 ?) Which energy in eV will an electron bring from its ground level to the first excited state ? 8. WHY ONLY CERTAIN ENERGY LEVELS ? The Bohr model cannot explain why electrons can only live in certain orbits ! When we look at electrons as WAVES, we can understand that each orbit must be a mulitple of /2 9. CLASSICAL VS. QUANTUM MECHANICS Quantum mechanics describes the function which represents these waves 10. FROM WAVE TO ENERGY Euler representation 11. DIFFERENTIATE with respect to x => impuls p = m*v - - t => energy E = mv2 = p2/2m 12. FROM LINE SPECTRA TO WAVEFUNCTIONS (ORBITALS) Model the electron as a standing wave in 3D, we can describe the most likely places of an electron and its energy from the Schroedinger Equation If you want to know this in detail: and: This equation leads to 3 quantum numbers which describe the energy and the distribution of the electron in an atom 13. QUICK OVERVIEW ABOUT SCHROEDINGER EQ. From: 14. SCHROEDINGER EQUATION It can describe the behavior of any small particle in the micro cosmos in theory 15. Main spectral lines = n Fine structure = l with magnetic field: Zeeman effect magnetic quantum no. m => 3 quantum numbers 16. ELECTRONIC SHIELDING AUFBAU PRINCIPLE (PERIODIC TABLE) 17. SHIELDING EFFECT Simplified demonstration 18. SLATERS RULES ZEFF = Z S ESTIMATION OF S: 19. EXAMPLE: K WHERE IS THE 19TH ELECTRON ? 20. HOMEWORK (PRESENT NEXT LESSON) Calculate the shielding for the valence electron(s) of: Ca compare 4s2 3 d2 Sc compare 3d1 4 p1 Cu (1) compare 4s1 4 p1 Cu (2) compare 3d10 4s1 3d9 4s2 Mn compare 3d5 4s2 3d7 Co (1) compare 3d7 4s2 3 d9 Co (2) compare 3d7 4s2 3d8 4s1 Cr (1) compare 3d5 4s1 3d4 4s2 Cr (2) compare 3d5 4s1 3 d6 Questions: explain 1. How shielding determines the AUFBAU principle 2. trend of atomic radius in PT (left to right) 3. -- ionization energies -- -- 4. - - electronegativities -- 21. ATOMIC RADII 22. IONIZATION ENERGIES 23. ELECTRO NEGATIVITY 24. PERIODIC TABLE AND TRENDS Watch clip on: 25. ***** BREAK ***** 26. 3 RULES FOR CONFIGURATIONS Aufbau Principle: Electrons are filled according to their lowest energy possible Pauli exclusion principle: Electrons must differ in one of 4 quantum numbers => max 2 electrons in one orbital Hunds Rules: Electrons want to have maximum SPIN 27. ELECTRONIC SHIELDING AUFBAU PRINCIPLE (PERIODIC TABLE) 28. ARE THESE CONFIGURATIONS FOR GROUND STATE POSSIBLE ? 1s 1s 2s 1s 2s 2p 1s 2s 29. SHAPES OF ORBITALS CHARACTERISTIC IS THE NO. OF NODE PLANES 2p orbital: 1 node 3d orbital: 2 nodes 30. FINE STRUCTURE OF HYDROGEN SPECTRA When an electron is in an orbital, it can cause different energies because of 2 forces: SPIN ANGULAR MOMENTUM 31. 3 MOMENTUMS: SPIN, ANGULAR AND SUM Watch clip until t=6:20 mins 32. ELECTRON HAS 3 MOMENTUMS An electron rotates around itself (like the earth) producing a SPIN S This produces a magnetic field around the electron 33. ANGULAR MOMENTUM L Because the electron circles around the nucleus, it creates an angular momentum ( L ) which creates also a magnetic field Since only certain radius are possible, L can have only discrete values 34. SPIN-ORBIT COUPLING J Both momentums combine to the total angular momentum J Lower energy ! 35. TOTAL ANGULAR MOMENTUM J 2D5/2 and 2D3/2 state 36. CONSEQUENCE An electron in an s-orbital has angular momentum of zero on average (L=0) An electron in a p orbital can have 2 different energies: depending if the spin momentum points in the same direction of the angular momentum or opposite. A smaller J (L-S) means lower energy than higher J (L+S) 37. TERM SYMBOLS (RUSELL-SAUNDER) To describe the electron configuration in an atom: 2S+1LJ L: orbital of the electrons (S =0, P =1, D =2, F =3) S: total spin of all these electrons J = L+S, L+S-1, L+S-2, ....|L-S| Orbital angular momentum multiplicity Total angular momentum 38. EXAMPLES Hydrogen ground state 1s1 2S1/2 Helium: 1s2 1S0 He(1s12s1) Excited State Configuration Terms: 1S0 , 3S1 Boron B(1s22s22p1) Terms: 2P1/2, 2P3/2 39. ROUSELL-SAUNDERS COUPLING (LS) Watch the clip 40. APPLICATION ON THE NA SPECTRUM 41. HYDROGEN EXCITED STATES 42. WATCH A DEMO VIDEO TO LEARN ABOUT S AND L 43. POSSIBLE CONFIGURATIONS FOR H 44. POSSIBLE TRANSITIONS 45. ELECTRON TRANSITIONS IN NA ATOM 46. 3 NEW QUANTUM NUMBERS (1) Angular Momentum L = the unfilled highest shell of the electron(s) Add l of each electron in this shell: l1 + l2 , . , |l1 - l2| (min.0) Example: 2 electrons in p shell: L = (1+1 ) 2 -> 1 -> 0 47. (2) MULTIPLICITY S For 2 valence electrons S can be 1 or 3 48. (3) TOTAL ANGULAR MOMENTUM J Example: Carbon Atom 1s2 2s2 2p2 (1) unfilled shell: p (l1, l2=1) => L= 2,1,0 (D,P,S) (2) possible Spin: S = +1/2 +1/2 = 1 or S = +1/2 -1/2 = 0 => multiplicity 3 (triplet) or 1 (singlet) (3) total momentum J: S=0: L+S: 2,1,0 S =1: L+S: 3,2,1,0 States: 1S0 3S1 1P1 and 3P0, 3P1, 3P2 and 1D2 and 3D1 3D2 3D3 49. CONFIGURATIONS FOR CARBON 50. ENERGY ORDER Apply Hunds rules to find the lowest energy ground state: 1. Term with highest S 2. Within the same S, the term with highest L 3. Within same S and L: a) shell half-filled or less: lowest J b) more than half-filled: highest J (for excited states, we cannot get the lowest energy from these rules) 51. EXAMPLE C ATOM GROUND STATE 2 p electrons with max. L in a configuration with highest spin S=1 => L=1 (P) => J = 2,1,0 Less than half-filled => lowest J: => ground state: S=1, L=1, J=0: 3P0 52. EXAMPLE C ATOM EXCITED STATE Configuration : 2 p1 and 3 s1 Spin: +1/2 +1/2 or +1/2 -1/2 => S= 0,1 1 electron in s orbital => l1 = 0 and 1 el. in p orbital => l2 = 1 => L= 1+0, 1-0 = 1 (P) => S=0: J=1 S=1: J = 2,1,0 Possible states: 1P1 and 3P0, 3P1, 3P2 53. MICROSTATES FOR CARBON C(1S22S22P2) 54. EXAMPLE FE ATOM Configuration 4s2 3d6 (1) Spin: 2 (2) L: -2, -1, 0, 1, 2 the max.L is in this state: L=2 (-2 -1 +0 +1 +2 +2) (3) J: 4,3,2,1,0 Hunds rule: max. spin, L=2 (D) highest J (more than half-filled): 4 => Ground state: 5D4 55. HOMEWORK (1) 1. Ground state conf. for He 2. Excited state (a) 1s1 2s1 3. Excited state (b) 1s1 2p1 4. Ground state conf. for Be 2s2 5. Excited state 2s1 2p1 56. SUMMARY 57. ***** BREAK ***** Go to Molecular Structure and Bonding