3. Atomic Emission Spectra zZumdahl 2 : p. 290-299Atoms Let go A range

4. What is light? zWhite light: reflection of all colors zBlack light: absorption of all colors zColors are each a different wavelength (λ: lamda) of light

5. zColors yDifferent wavelengths of light are seen as different colors. yDifferent colors indicate (show) different energy levels.

6.  c=fλ  c=fλ (velocity of light = frequency x wavelength) ythe greater the frequency the shorter the wavelength zΔE = hf y(energy lost by the electron = h(constant) x frequency yFrequency (and thus, color) of the light depends on the amount of energy lost by the electron.

7. c=fλ

9. zWhen atoms are “exited” (energy is added) they produce light. zNot white or all-colored light, but one color at a time.

10. zStudy the light emitted (produced) by atoms and ions to deduce (find out) the structure of atoms. zWhen an atom is “excited” its electrons gain energy and move to a higher energy level. To return to a lower energy level, electrons must lose energy. They do this by giving off light.

12.  Continuous spectrum:  Continuous spectrum: all wavelengths of visible light contained in white light. line spectrum  Light emitted by an atom can be separated into a line spectrum that shows exactly what frequencies of light are present.

14. Closer to nucleus Further from nucleus Increasing frequency

15. line spectrum  Because the light emitted from atoms is a line spectrum (not a continuous spectrum) we determine that: energy levels  There are “discrete” (separate) energy levels for each atom that can only produce light of certain wavelengths (this is NOT ordinary white light!).

18. Increasing frequency (f) (increasing energy)

21. Hydrogen zOnly certain energy levels can occur (not a continuous spectrum)

22. Energy Level Diagram zThe larger the difference in energy, the greater the frequency (thus, the more purple the light).

23. Increasing potential energy Frequency Visible

24.  convergence:  convergence: the lines in a spectrum converge (get closer together) as frequency increases. yrelated to how much energy is required to remove the electron from the atom (ionize)

25. Closer to nucleus Further from nucleus Increasing frequency

26. zStop

27. Electronic Structure Energy Levels Shells

28. zMost stable = closest to nucleus

29. z1st energy level = 2 z2nd energy level = 8 zElectronic structure: number of electrons in each orbital

31. zH=1 zO=2,6 (two electrons in the first energy level, six in the second) zAl=2,8,3 zCl= zCa= yDifferent isotopes have the same electronic structure and the same chemical properties!

32. Electron Behavior zValence shell: outer shell of an atom ydetermine the physical and chemical properties of an atom Valence Shell

34. zHow many electrons in valence shell? yAl yNe yLi yCa

35. zStop here

36. HL Topic Electronic Structure of Atoms Zumdahl 2 : p. 307-312

37. Electronic Structure I.Energy levels A. Sub-levels 1. Orbital a. Spin

38. Energy Levels zMajor shells (layers) around the nucleus yfilled before higher levels are filled y1st: 2 electrons y2nd: 8 electrons y3rd: 8 electrons

39. Sub-levels zDifferent shapes zs – sphere yone orbital zp – figure eight ythree orbitals zd – yfive orbitals zf – yseven orbitals

41. p Sub-level zp sub-level has three orbitals zp x, p y, p z

42. zd and f sub-levels have very complex shapes

43. Orbitals zEach orbital can hold two electrons. zElectrons spin in opposite directions

44. Energy Level Types of sub-levels Total orbits Electron capacity 1s12 2s, p1+3= 4 8 3s, p, d1+3+5= 9 18 4s, p, d, f1+3+5+7= 16 2

45. Energies of sub-levels

48. Electronic structure of atoms

49. Energy States zDepending on where an electron was around the nucleus, it had different energy states. zGround state: An orbit near the nucleus: not very exited at all. zExcited State: An orbit farther away from the nucleus: much more potential for giving off energy.

50. Heisenberg Uncertainty Principle zIf electrons are both waves and particles, where are they around the atom? zIt is impossible to figure out both the position and velocity of an electron, at the same time. zWe CAN figure out the probability that an electron is in any one spot at any given time.

51. Electron Configuration zThe arrangement of electrons in an atom yEach element’s atoms are different yArrangement with the lowest energy= ground state electron configuration zHow do we figure out what the ground state electron configuration looks like?

52. How do we figure out where the electrons are? z1. Figure out the energy levels of the orbitals z2. Add electrons to the orbitals according to three rules

53. 1: Aufbau Principle zAn electron goes to the lowest-energy orbital that can take it.

54. 2: Pauli Exclusion Principle zNo two electrons can have exactly the same configuration description yCan have the same orbital, but must have opposite spins.

55. 3. Hund’s Rule zOrbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron zAll electrons that are by themselves in an orbit must have the same spin.

56. zElectron configuration: the arrangement of electrons in an atom. zSo…what do we do with this information?!

57. How do we write electron configurations? zElectron Configuration Notation zOrbital Notation zNoble Gas Notation (shorthand) yAll ways to communicate where the electrons are in the ground state of any atom.

58. zs – sphere yone orbital – 2 electrons zp – figure eight ythree orbitals – 6 electrons zd – yfive orbitals – 10 electrons zf – yseven orbitals – 14 electrons

59. Orbital Notation zElectron configuration goes below a line or box zArrows representing electrons go on the lines or in the boxes 1s 2s 2p 3s 3p 4s

60. zFluorine? 9 electrons zMagnesium? 12 electrons 1s 2s 2p 3s 3p 4s

61. Electron Configuration Notation zMain energy level Sub-level Electrons zCarbon (6 electrons) 1s 2 2s 2 2p 2 zAluminum (13 electrons) y1s 2 2s 2 2p 6 3s 2 3p 1 zOxygen zArgon zCopper

62. Orbital Notation zBoron z1s 2 2s 2 2p 1 zAtomic Number? zHow many electrons? zOrbital notation

63. zAufbau Principle: start with 1s and work up in energy level zPauli Exclusion Principle: No two elements can have the same arrangement of electrons zHund’s Rule: Fill in one electron per orbital first, then go back.

64. Orbital Notation zNitrogen z1s 2 2s 2 2p 3 zAtomic Number? zNumber of Electrons? zOrbital Notation?

65. Practice zWrite the orbital notation for: zFluorine zAluminum zCarbon zOxygen

68. Practice zWrite the electron configuration notation of: zBe (Beryllium) zN (Nitrogen) zSi (Silicon) zNa (Sodium)

69. Noble Gas Notation (Shorthand) zNoble gasses have totally filled outer orbitals. zIf Ne (a noble gas) is 1s 2 2s 2 2p 6, we can abbreviate Na as [Ne]3s 1. zSodium has one more electron than Neon, so its Noble Gas Notation is Neon plus one electron in the s sublevel of the third energy level.

70. Practice zFe yElectron Configuration Notation yNoble Gas Notation zKzK yElectron Configuration Notation yNoble Gas Notation zLi zBe