Arrangement of Electrons in Atoms

1. Arrangement of Electrons in Atoms Chapter 4

2. The New Atomic Model • Investigations  relationship between light and atom’s electrons • How are electrons arranged? Why don’t they fall into the nucleus?

3. Light a wave or particle? • Wave Description: • Electromagnetic Radiation: energy that acts like a wave in space • All forms create Electromagnetic Spectrum

4. Electromagnetic Spectrum

5. Electromagnetic Spectrum • All forms move at speed of light, c, 3.00x108 m/s • Forms identified by: • wavelength, , the distance b/ corresponding points on adjacent waves. Units: nm, cm, or m • Frequency, , # of waves that pass a given point in a specific time, 1 sec. Unit: 1/s = Hertz, Hz

6. Wavelength and Frequency

7. Wavelength and Frequency c =  • Inverse proportion equation!! Frequency, 1/s speed of light, m/s wavelength, m

8. Calculation • Calculate the wavelength of a radio wave with a frequency of 102.7 x 106s-1 • Determine the frequency of light whose wavelength is 5.267 nm.

9. Particle Nature of Light • Photoelectric Effect: emission of electrons from a metal when light shines on the metal • Video -13

10. Photoelectric Effect • Light had to be certain frequency to knock e- loose • Wave theory  any frequency should work (just might take a while) • Light must also be a particle! • Max Planck(1900) explanation: objects emit energy in small packets called quanta • Video - 16

11. Max Planck • Quantum of energy is the smallest amount of energy that can be lost or gained by an atom E = h Frequency, s-1 Energy of quantum, in joules, J Planck’s constant, 6.626x10-34 Js

12. Energy Calculation • What is the energy of green light, with a wavelength of 500. nm?

13. Albert Einstein • Light is both wave and particle! • Particle of light = photon, having zero mass and a quantum of energy • Photons hit metal and knock e- out, but photon has to have enough energy

14. H-atom Emission Spectrum • Pass a current through gas at low pressure it excites the atoms • Ground state: lowest energy state of an atom • Excited state: atom has higher potential energy than it has in ground state

15. H – Atom Spectrum • When atom jumps from excited state to ground state it gives off energy  LIGHT! E2 Ephoton = E2 – E1 = hv E1

16. H-atom Line Emission Spectrum

17. Element Emission Spectras Helium – 23 lines Neon – 75 lines Argon - 159 lines Xenon – 139 lines Mercury – 40 lines

18. H-atom Line Emission Spectrum • More lines in UV (Lyman series) and IR(Paschen series) • Why did H-atom only emit certain colors of light?

19. Bohr Model of H-atom • 1913 – Niels Bohr • e- circles nucleus in certain paths, orbits or atomic energy levels • e- is higher in energy the farther away from nucleus • e- cannot be between orbits • Video - 23

20. Bohr Model of H-atom

21. Bohr Model of H-atom • From wavelengths of emission spectrum Bohr calculated energy levels of H-atom • Model worked ONLY for H-atom • End Part 1

22. Quantum Model of Atom – Part 2 • Can e- behave as a wave? • Yes! • To find e- use a photon, but photon will knock the e- off course • Heisenberg Uncertainty Principle: impossible to determine position and velocity of a particle at the same time.

23. Schrödinger Wave Equation • 1926 – developed equation and only e- waves of certain frequencies were solutions • Quantization of e- probability of finding e- in atom • No neat orbits  probability clouds or orbitals

24. Electron Configurations

25. Atomic Orbitals • Def: 3-D region around nucleus that indicates the probable location of an electron • Energy levels or shells: • Numbered 1-7 • Smaller number = closer to nucleus, lower energy

26. Sublevels • Each shell has sublevels • s • 1 – s orbital • p • 3 – p orbitals • d • 5 – d orbitals • f • 7 – f orbitals

27. Shells and Sublevels Shells and sublevels together: 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f, etc. s is the lowest energy and f is the highest

28. Orbitals • Each orbital in a sublevel can hold a maximum of 2 e- • 1 – s 2 e- max. • 3 – p orbitals 6 e- max. • 5 – d orbitals 10 e- max. • 7 – f orbitals 14 e- max.

29. Electron Configurations • Arrangement of e- in atom • Orbital Notation: • H has 1e- • Rules: • Aufbau Principle: electron occupies lowest energy level that can receive it

30. Electron Configurations 2. Pauli Exclusion Principle: no two e- in an sublevel orbital can have the same spin He – 2e- 3. Hund’sRule: orbitals of equal energy are occupied by one e- before pairing up e-. All single occupied orbitals must have same spin.

31. Energy of sublevels

32. Electron Configurations • N • S • Ti • I

33. Electron Configuration Notation • B • Ni • Hg

34. Noble Gas Notation • Use noble gas from previous row • Al • Pb

35. Special Cases • d sublevel more stable with half-filled or completely filled sublevel • Cr • Cu

36. Quantum Numbers Def: specify the properties of atomic orbitals and the properties of electrons in orbitals. 4 quantum #s - first 3 from Schrödinger

37. Principal Quantum Number n, indicates the main energy level only + integers Total # of orbitals in an energy level is n2

38. Angular Momentum Quantum # l, indicates the shape of orbital (sublevel) # of shapes possible = n (# of energy level) l = 0 and all + integers less than or equal to n-1 1st E level, l = 0 (s) 2nd E level, l = 0 (s), 1(p) 3rd E level, l = 0(s), 1(p), 2(d)

39. Magnetic Quantum Number m, indicates orientation(direction) of orbital m = -l to +l, including 0 l = 0, m = 0(s) l = 1, m = -1(px), 0(py), 1(pz) l = 2, m = -2, -1, 0, 1, 2 (5 d orbitals)

40. Spin Quantum Number Only two values: - ½ and + ½ Indicates the spin, up or down

41. Pauli Exclusion Principle No two e- can have the same four quantum numbers 2s2 n = 2, l = 0, m = 0, + ½ (1st e-) n = 2, l = 0, m = 0, - ½ (1st e-)

42. Examples 3d3 n = 3, l = 2, m = -2, + ½ (1st e-) n = 3, l = 2, m = -1, + ½ (2nd e-) n = 3, l = 2, m = 0, + ½ (3rd e-)

43. Examples What is the configuration of the electron with the following quantum numbers? n = 4, l = 1, m = -1, + ½