1. I. Waves & Particles Ch. 4 - Electrons in Atoms

2. A. Waves zWavelength ( ) - length of one complete wave zFrequency ( ) - # of waves that pass a point during a certain time period yhertz (Hz) = 1/s zAmplitude (A) - distance from the origin to the trough or crest

3. A. Waves A greater amplitude (intensity) greater frequency (color) crest origin trough A

4. B. EM Spectrum LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY

5. LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY ROYG.BIV redorangeyellowgreenblueindigoviolet

6. B. EM Spectrum zFrequency & wavelength are inversely proportional c = c:speed of light (3.00  10 8 m/s) :wavelength (m, nm, etc.) :frequency (Hz)

7. B. EM Spectrum GIVEN: = ? = 434 nm = 4.34  10 -7 m c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 4.34  10 -7 m = 6.91  10 14 Hz zEX: Find the frequency of a photon with a wavelength of 434 nm.

8. C. Quantum Theory zPlanck (1900) yObserved - emission of light from hot objects yConcluded - energy is emitted in small, specific amounts (quanta) yQuantum - minimum amount of energy change

9. C. Quantum Theory zPlanck (1900) vs. Classical TheoryQuantum Theory

10. C. Quantum Theory zEinstein (1905) yObserved - photoelectric effect

11. C. Quantum Theory zEinstein (1905) yConcluded - light has properties of both waves and particles “wave-particle duality” yPhoton - particle of light that carries a quantum of energy

12. C. Quantum Theory E:energy (J, joules) h:Planck’s constant (6.6262  10 -34 J·s) :frequency (Hz) E = h zThe energy of a photon is proportional to its frequency.

13. C. Quantum Theory GIVEN: E = ? = 4.57  10 14 Hz h = 6.6262  10 -34 J·s WORK : E = h E = ( 6.6262  10 -34 J·s ) ( 4.57  10 14 Hz ) E = 3.03  10 -19 J zEX: Find the energy of a red photon with a frequency of 4.57  10 14 Hz.

14. II. Bohr Model of the Atom Ch. 4 - Electrons in Atoms

15. A. Line-Emission Spectrum ground state excited state ENERGY IN PHOTON OUT

16. B. Bohr Model ze - exist only in orbits with specific amounts of energy called energy levels zTherefore… ye - can only gain or lose certain amounts of energy yonly certain photons are produced

17. B. Bohr Model 1 2 3 4 5 6 zEnergy of photon depends on the difference in energy levels zBohr’s calculated energies matched the IR, visible, and UV lines for the H atom

18. C. Other Elements zEach element has a unique bright-line emission spectrum. y“Atomic Fingerprint” Helium zBohr’s calculations only worked for hydrogen! 

19. III. Quantum Model of the Atom Ch. 4 - Electrons in Atoms

20. A. Electrons as Waves zLouis de Broglie (1924) yApplied wave-particle theory to e - ye - exhibit wave properties QUANTIZED WAVELENGTHS

21. A. Electrons as Waves QUANTIZED WAVELENGTHS

22. A. Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT

23. B. Quantum Mechanics zHeisenberg Uncertainty Principle yImpossible to know both the velocity and position of an electron at the same time

24. B. Quantum Mechanics zSchrödinger Wave Equation (1926) yfinite # of solutions  quantized energy levels ydefines probability of finding an e -

25. B. Quantum Mechanics Radial Distribution Curve Orbital zOrbital (“electron cloud”) yRegion in space where there is 90% probability of finding an e -

26. C. Quantum Numbers UPPER LEVEL zFour Quantum Numbers: ySpecify the “address” of each electron in an atom

27. C. Quantum Numbers 1. Principal Quantum Number ( n ) yEnergy level ySize of the orbital yn 2 = # of orbitals in the energy level

28. C. Quantum Numbers s p d f 2. Angular Momentum Quantum # ( l ) yEnergy sublevel yShape of the orbital

29. C. Quantum Numbers zn=# of sublevels per level zn 2 =# of orbitals per level zSublevel sets: 1 s, 3 p, 5 d, 7 f

30. C. Quantum Numbers 3. Magnetic Quantum Number ( m l ) yOrientation of orbital  Specifies the exact orbital within each sublevel

31. C. Quantum Numbers pxpx pypy pzpz

32. zOrbitals combine to form a spherical shape. 2s 2p z 2p y 2p x

33. C. Quantum Numbers 4. Spin Quantum Number ( m s ) yElectron spin  +½ or -½ yAn orbital can hold 2 electrons that spin in opposite directions.

34. C. Quantum Numbers 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orbital electron zPauli Exclusion Principle yNo two electrons in an atom can have the same 4 quantum numbers. yEach e - has a unique “address”:

35. Feeling overwhelmed? Read Section 4-2!

36. IV. Electron Configuration (p. 105 - 116, 128 - 139) Ch. 4 - Electrons in Atoms

37. A. General Rules zPauli Exclusion Principle yEach orbital can hold TWO electrons with opposite spins.

38. A. General Rules zAufbau Principle yElectrons fill the lowest energy orbitals first. y“Lazy Tenant Rule”

39. RIGHT WRONG A. General Rules zHund’s Rule yWithin a sublevel, place one e - per orbital before pairing them. y“Empty Bus Seat Rule”

40. O 8e - zOrbital Diagram zElectron Configuration 1s 2 2s 2 2p 4 B. Notation 1s 2s 2p

41. zShorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 B. Notation zLonghand Configuration

42. © 1998 by Harcourt Brace & Company s p d (n-1) f (n-2) 12345671234567 6767 C. Periodic Patterns

43. zPeriod # yenergy level (subtract for d & f) zA/B Group # ytotal # of valence e - zColumn within sublevel block y# of e - in sublevel

44. s-block1st Period 1s 1 1st column of s-block C. Periodic Patterns zExample - Hydrogen

45. C. Periodic Patterns zShorthand Configuration yCore e - : Go up one row and over to the Noble Gas. yValence e - : On the next row, fill in the # of e - in each sublevel.

46. [Ar]4s 2 3d 10 4p 2 C. Periodic Patterns zExample - Germanium

47. zFull energy level zFull sublevel (s, p, d, f) zHalf-full sublevel D. Stability

48. zElectron Configuration Exceptions yCopper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10 yCopper gains stability with a full d-sublevel. D. Stability

49. zElectron Configuration Exceptions yChromium EXPECT :[Ar] 4s 2 3d 4 ACTUALLY :[Ar] 4s 1 3d 5 yChromium gains stability with a half-full d-sublevel. D. Stability

50. zIon Formation yAtoms gain or lose electrons to become more stable. yIsoelectronic with the Noble Gases.

51. O 2- 10e - [He] 2s 2 2p 6 D. Stability zIon Electron Configuration yWrite the e - config for the closest Noble Gas yEX: Oxygen ion  O 2-  Ne

52. Ch. 5 - The Periodic Table I. History

53. A. Mendeleev zDmitri Mendeleev (1869, Russian) yOrganized elements by increasing atomic mass. yElements with similar properties were grouped together. yThere were some discrepancies.

54. A. Mendeleev zDmitri Mendeleev (1869, Russian) yPredicted properties of undiscovered elements.

55. B. Moseley zHenry Mosely (1913, British) yOrganized elements by increasing atomic number. yResolved discrepancies in Mendeleev’s arrangement.

56. II. Organization of the Elements Ch. 5 - The Periodic Table

57. zMetals zNonmetals zMetalloids A. Metallic Character

58. zMain Group Elements zTransition Metals zInner Transition Metals B. Blocks

59. III. Periodic Trends Ch. 5 - The Periodic Table

60. A. Periodic Law zWhen elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

61. B. Chemical Reactivity zFamilies ySimilar valence e - within a group result in similar chemical properties

62. B. Chemical Reactivity zAlkali Metals zAlkaline Earth Metals zTransition Metals zHalogens zNoble Gases

63. zAtomic Radius ysize of atom © 1998 LOGAL zFirst Ionization Energy yEnergy required to remove one e - from a neutral atom. © 1998 LOGAL zMelting/Boiling Point C. Other Properties

64. zAtomic Radius D. Atomic Radius Li Ar Ne K Na

65. zAtomic Radius yIncreases to the LEFT and DOWN D. Atomic Radius

66. zWhy larger going down? yHigher energy levels have larger orbitals yShielding - core e - block the attraction between the nucleus and the valence e - zWhy smaller to the right? yIncreased nuclear charge without additional shielding pulls e - in tighter D. Atomic Radius

67. zFirst Ionization Energy E. Ionization Energy K Na Li Ar Ne He

68. zFirst Ionization Energy yIncreases UP and to the RIGHT E. Ionization Energy

69. zWhy opposite of atomic radius? yIn small atoms, e - are close to the nucleus where the attraction is stronger zWhy small jumps within each group? yStable e - configurations don’t want to lose e - E. Ionization Energy

70. zSuccessive Ionization Energies yMg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ yLarge jump in I.E. occurs when a CORE e - is removed. E. Ionization Energy

71. yAl1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ zSuccessive Ionization Energies yLarge jump in I.E. occurs when a CORE e - is removed. E. Ionization Energy

72. zMelting/Boiling Point yHighest in the middle of a period. F. Melting/Boiling Point

73. zIonic Radius yCations (+) xlose e - xsmaller © 2002 Prentice-Hall, Inc. yAnions (–) xgain e - xlarger G. Ionic Radius

74. zWhich atom has the larger radius? yBeorBa yCaorBr Ba Ca Examples

75. zWhich atom has the higher 1st I.E.? yNorBi yBaorNe N Ne Examples

76. zWhich atom has the higher melting/boiling point? yLiorC yCrorKr C Cr Examples

77. zWhich particle has the larger radius? ySorS 2- yAlorAl 3+ S 2- Al Examples

78. Periodic Trends Summary