1. John C. Kotz State University of New York, College at Oneonta John C. Kotz Paul M. Treichel John Townsend http://academic.cengage.com/kotz Chapter 6 Atomic Structure

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3. 3 © 2009 Brooks/Cole - Cengage ATOMIC STRUCTURE PLAY MOVIE

4. 4 © 2009 Brooks/Cole - Cengage ELECTROMAGNETIC RADIATION PLAY MOVIE

5. 5 © 2009 Brooks/Cole - Cengage Electromagnetic Radiation Most subatomic particles behave as PARTICLES and obey the physics of waves.Most subatomic particles behave as PARTICLES and obey the physics of waves. PLAY MOVIE

6. 6 © 2009 Brooks/Cole - Cengage wavelength Visible light wavelength Ultraviolet radiation Amplitude Node Electromagnetic Radiation

7. 7 © 2009 Brooks/Cole - Cengage Electromagnetic Radiation See Figure 6.1

8. 8 © 2009 Brooks/Cole - Cengage Waves have a frequencyWaves have a frequency Use the Greek letter “nu”,, for frequency, and units are “cycles per sec”Use the Greek letter “nu”,, for frequency, and units are “cycles per sec” All radiation: = cAll radiation: = c c = velocity of light = 3.00 x 10 8 m/sec c = velocity of light = 3.00 x 10 8 m/sec Long wavelength f small frequencyLong wavelength f small frequency Short wavelength f high frequencyShort wavelength f high frequency Electromagnetic Radiation

9. 9 © 2009 Brooks/Cole - Cengage Long wavelength f small frequency Short wavelength f high frequency increasing frequency increasing wavelength Electromagnetic Radiation PLAY MOVIE

10. 10 © 2009 Brooks/Cole - Cengage Red light has = 700 nm. Calculate the frequency. Electromagnetic Radiation

11. 11 © 2009 Brooks/Cole - Cengage Long wavelength f small frequency small frequency low energy Short wavelength f high frequency high frequency high energy Electromagnetic Radiation

12. 12 © 2009 Brooks/Cole - Cengage Electromagnetic Spectrum See Active Figure 6.2

13. 13 © 2009 Brooks/Cole - Cengage Quantization of Energy Max Planck (1858-1947) Solved the “ultraviolet catastrophe” See Chem & Chem Reactivity, Figure 6.3 PLAY MOVIE

14. 14 © 2009 Brooks/Cole - Cengage E = h · E = h · Quantization of Energy Energy of radiation is proportional to frequency h = Planck’s constant = 6.6262 x 10 -34 J·s An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA.

15. 15 © 2009 Brooks/Cole - Cengage Light with a short (large ) has a large E. Light with large (small ) has a small E. E = h · E = h · Quantization of Energy

16. 16 © 2009 Brooks/Cole - Cengage Photoelectric Effect Experiment demonstrates the particle nature of light.

17. 17 © 2009 Brooks/Cole - Cengage Photoelectric Effect Classical theory said that E of ejected electron should increase with increase in light intensity—not observed! No e - observed until light of a certain minimum E is used.No e - observed until light of a certain minimum E is used. Number of e - ejected depends on light intensity.Number of e - ejected depends on light intensity. A. Einstein (1879-1955)

18. 18 © 2009 Brooks/Cole - Cengage PROBLEM: Calculate the energy of 1.00 mol of photons of red light. = 700. nm = 700. nm = 4.29 x 10 14 sec -1 = 4.29 x 10 14 sec -1 PROBLEM: Calculate the energy of 1.00 mol of photons of red light. = 700. nm = 700. nm = 4.29 x 10 14 sec -1 = 4.29 x 10 14 sec -1 Photoelectric Effect Understand experimental observations if light consists of particles called PHOTONS of discrete energy.

19. 19 © 2009 Brooks/Cole - Cengage Energy of Radiation Energy of 1.00 mol of photons of red light. E = h· E = h· = (6.63 x 10 -34 J·s)(4.29 x 10 14 s -1 ) = (6.63 x 10 -34 J·s)(4.29 x 10 14 s -1 ) = 2.85 x 10 -19 J per photon = 2.85 x 10 -19 J per photon E per mol = (2.85 x 10 -19 J/ph)(6.02 x 10 23 ph/mol) (2.85 x 10 -19 J/ph)(6.02 x 10 23 ph/mol) = 172 kJ/mol = 172 kJ/mol This is in the range of energies that can break bonds.

20. 20 © 2009 Brooks/Cole - Cengage Excited Atoms & Atomic Structure

21. 21 © 2009 Brooks/Cole - Cengage Atomic Line Emission Spectra and Niels Bohr Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the SHARP LINE EMISSION SPECTRA of excited atoms. Niels Bohr (1885-1962)

22. 22 © 2009 Brooks/Cole - Cengage Spectrum of White Light

23. 23 © 2009 Brooks/Cole - Cengage Line Emission Spectra of Excited Atoms Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the element. PLAY MOVIE

24. 24 © 2009 Brooks/Cole - Cengage Spectrum of Excited Hydrogen Gas See Active Figure 6.6

25. 25 © 2009 Brooks/Cole - Cengage Visible lines in H atom spectrum are called the BALMER series. High E Short Short High High Low E Long Long Low Low Line Emission Spectra of Excited Atoms

26. 26 © 2009 Brooks/Cole - Cengage Line Spectra of Other Elements See Figure 6.7

27. 27 © 2009 Brooks/Cole - Cengage The Electric Pickle Excited atoms can emit light. Here the solution in a pickle is excited electrically. The Na + ions in the pickle juice give off light characteristic of that element. PLAY MOVIE

28. 28 © 2009 Brooks/Cole - Cengage Atomic Spectra and Bohr 1.Any orbit should be possible and so is any energy. 2.But a charged particle moving in an electric field should emit energy. End result should be destruction! One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.

29. 29 © 2009 Brooks/Cole - Cengage Atomic Spectra and Bohr Bohr said classical view is wrong. Need a new theory — now called QUANTUM or WAVE MECHANICS. e- can only exist in certain discrete orbits — called stationary states. e- is restricted to QUANTIZED energy states. Energy of state = - C/n 2 where n = quantum no. = 1, 2, 3, 4,....

30. 30 © 2009 Brooks/Cole - Cengage Atomic Spectra and Bohr Only orbits where n = integral no. are permitted.Only orbits where n = integral no. are permitted. Radius of allowed orbitals = n 2 · (0.0529 nm)Radius of allowed orbitals = n 2 · (0.0529 nm) But note — same eqns. come from modern wave mechanics approach.But note — same eqns. come from modern wave mechanics approach. Results can be used to explain atomic spectra.Results can be used to explain atomic spectra. Energy of quantized state = - C/n 2

31. 31 © 2009 Brooks/Cole - Cengage Atomic Spectra and Bohr If e-’s are in quantized energy states, then ∆E of states can have only certain values. This explain sharp line spectra. PLAY MOVIE

32. 32 © 2009 Brooks/Cole - Cengage Energy Adsorption/Emission See Active Figure 6.9

33. 33 © 2009 Brooks/Cole - Cengage Atomic Spectra and Bohr Calculate ∆E for e- “falling” from high energy level (n = 2) to low energy level (n = 1). ∆E = E final - E initial = -C[(1/1 2 ) - (1/2) 2 ] ∆E = -(3/4)C Note that the process is EXOTHERMIC

34. 34 © 2009 Brooks/Cole - Cengage Atomic Spectra and Bohr ∆E = -(3/4)C C has been found from experiment (and is now called R, the Rydberg constant) R (= C) = 1312 kJ/mol or 3.29 x 10 15 cycles/sec so, E of emitted light = (3/4)R = 2.47 x 10 15 sec -1 = (3/4)R = 2.47 x 10 15 sec -1 and = c/ = 121.6 nm This is exactly in agreement with experiment!

35. 35 © 2009 Brooks/Cole - Cengage Origin of Line Spectra Balmer series See Active Figure 6.10

36. 36 © 2009 Brooks/Cole - Cengage Atomic Line Spectra and Niels Bohr Bohr’s theory was a great accomplishment. Rec’d Nobel Prize, 1922 Problems with theory — theory only successful for H.theory only successful for H. introduced quantum idea artificially.introduced quantum idea artificially. So, we go on to QUANTUM or WAVE MECHANICSSo, we go on to QUANTUM or WAVE MECHANICS Niels Bohr (1885-1962)

37. 37 © 2009 Brooks/Cole - Cengage Quantum or Wave Mechanics de Broglie (1924) proposed that all moving objects have wave properties. For light: E = mc 2 E = h = hc / E = h = hc / Therefore, mc = h / Therefore, mc = h / and for particles (mass)(velocity) = h / (mass)(velocity) = h / de Broglie (1924) proposed that all moving objects have wave properties. For light: E = mc 2 E = h = hc / E = h = hc / Therefore, mc = h / Therefore, mc = h / and for particles (mass)(velocity) = h / (mass)(velocity) = h / L. de Broglie (1892-1987)

38. 38 © 2009 Brooks/Cole - Cengage Baseball (115 g) at 100 mph = 1.3 x 10 -32 cm = 1.3 x 10 -32 cm e- with velocity = 1.9 x 10 8 cm/sec = 0.388 nm = 0.388 nm Experimental proof of wave properties of electrons Quantum or Wave Mechanics PLAY MOVIE

39. 39 © 2009 Brooks/Cole - Cengage Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. He developed the WAVE EQUATION Solution gives set of math expressions called WAVE FUNCTIONS,  Each describes an allowed energy state of an e- Quantization introduced naturally. E. Schrodinger 1887-1961 Quantum or Wave Mechanics

40. 40 © 2009 Brooks/Cole - Cengage Wave motion: wave length and nodes “Quantization” in a standing wave

41. 41 © 2009 Brooks/Cole - Cengage WAVE FUNCTIONS,   is a function of distance and two angles.  is a function of distance and two angles. Each  corresponds to an ORBITAL — the region of space within which an electron is found. Each  corresponds to an ORBITAL — the region of space within which an electron is found.  does NOT describe the exact location of the electron.  does NOT describe the exact location of the electron.  2 is proportional to the probability of finding an e- at a given point.  2 is proportional to the probability of finding an e- at a given point.

42. 42 © 2009 Brooks/Cole - Cengage Uncertainty Principle Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= mv) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= mv) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. W. Heisenberg 1901-1976

43. 43 © 2009 Brooks/Cole - Cengage Types of Orbitals s orbital p orbital d orbital PLAY MOVIE

44. 44 © 2009 Brooks/Cole - Cengage Orbitals No more than 2 e- assigned to an orbitalNo more than 2 e- assigned to an orbital Orbitals grouped in s, p, d (and f) subshellsOrbitals grouped in s, p, d (and f) subshells s orbitals d orbitals p orbitals

45. 45 © 2009 Brooks/Cole - Cengage s orbitals d orbitals p orbitals s orbitals p orbitals d orbitals No.orbs. No. e- 135 2610

46. 46 © 2009 Brooks/Cole - Cengage Subshells & Shells Subshells grouped in shells.Subshells grouped in shells. Each shell has a number called the PRINCIPAL QUANTUM NUMBER, nEach shell has a number called the PRINCIPAL QUANTUM NUMBER, n The principal quantum number of the shell is the number of the period or row of the periodic table where that shell begins.The principal quantum number of the shell is the number of the period or row of the periodic table where that shell begins.

47. 47 © 2009 Brooks/Cole - Cengage Subshells & Shells n = 1 n = 2 n = 3 n = 4

48. 48 © 2009 Brooks/Cole - Cengage QUANTUM NUMBERS The shape, size, and energy of each orbital is a function of 3 quantum numbers: n (major) f shell s (angular) f subshell m s (magnetic) f designates an orbital within a subshell

49. 49 © 2009 Brooks/Cole - Cengage SymbolValuesDescription n (major)1, 2, 3,..Orbital size and energy where E = -R(1/n 2 ) s (angular)0, 1, 2,.. n-1Orbital shape or type (subshell) m s (magnetic)-s..0..+sOrbital orientation # of orbitals in subshell = 2s + 1 # of orbitals in subshell = 2s + 1 QUANTUM NUMBERS

50. 50 © 2009 Brooks/Cole - Cengage Types of Atomic Orbitals See Active Figure 6.14

51. 51 © 2009 Brooks/Cole - Cengage Shells and Subshells When n = 1, then s = 0 and m s = 0 Therefore, in n = 1, there is 1 type of subshell and that subshell has a single orbital (m s has a single value f 1 orbital) This subshell is labeled s (“ess”) Each shell has 1 orbital labeled s, and it is SPHERICAL in shape.

52. 52 © 2009 Brooks/Cole - Cengage s Orbitals— Always Spherical Dot picture of electron cloud in 1s orbital. Surface density 4πr 2  versus distance Surface of 90% probability sphere See Active Figure 6.13

53. 53 © 2009 Brooks/Cole - Cengage 1s Orbital PLAY MOVIE

54. 54 © 2009 Brooks/Cole - Cengage 2s Orbital PLAY MOVIE

55. 55 © 2009 Brooks/Cole - Cengage 3s Orbital PLAY MOVIE

56. 56 © 2009 Brooks/Cole - Cengage p Orbitals When n = 2, then s = 0 and 1 Therefore, in n = 2 shell there are 2 types of orbitals — 2 subshells For s = 0m s = 0 this is a s subshell this is a s subshell For s = 1 m s = -1, 0, +1 this is a p subshell with 3 orbitals this is a p subshell with 3 orbitals When n = 2, then s = 0 and 1 Therefore, in n = 2 shell there are 2 types of orbitals — 2 subshells For s = 0m s = 0 this is a s subshell this is a s subshell For s = 1 m s = -1, 0, +1 this is a p subshell with 3 orbitals this is a p subshell with 3 orbitals See Screen 6.15 When s = 1, there is a PLANAR NODE thru the nucleus.

57. 57 © 2009 Brooks/Cole - Cengage p Orbitals The three p orbitals lie 90 o apart in space

58. 58 © 2009 Brooks/Cole - Cengage 2p x Orbital 3p x Orbital PLAY MOVIE

59. 59 © 2009 Brooks/Cole - Cengage d Orbitals When n = 3, what are the values of s? s = 0, 1, 2 and so there are 3 subshells in the shell. For s = 0, m s = 0 f s subshell with single orbital f s subshell with single orbital For s = 1, m s = -1, 0, +1 f p subshell with 3 orbitals f p subshell with 3 orbitals For s = 2, m s = -2, -1, 0, +1, +2 f d subshell with 5 orbitals f d subshell with 5 orbitals

60. 60 © 2009 Brooks/Cole - Cengage d Orbitals s orbitals have no planar node (s = 0) and so are spherical. p orbitals have s = 1, and have 1 planar node, and so are “dumbbell” shaped. This means d orbitals (with s = 2) have 2 planar nodes See Figure 6.15

61. 61 © 2009 Brooks/Cole - Cengage 3d xy Orbital PLAY MOVIE

62. 62 © 2009 Brooks/Cole - Cengage 3d xz Orbital PLAY MOVIE

63. 63 © 2009 Brooks/Cole - Cengage 3d yz Orbital PLAY MOVIE

64. 64 © 2009 Brooks/Cole - Cengage 3d x 2 - y 2 Orbital PLAY MOVIE

65. 65 © 2009 Brooks/Cole - Cengage 3d z 2 Orbital PLAY MOVIE

66. 66 © 2009 Brooks/Cole - Cengage f Orbitals When n = 4, s = 0, 1, 2, 3 so there are 4 subshells in the shell. For s = 0, m s = 0 f s subshell with single orbital f s subshell with single orbital For s = 1, m s = -1, 0, +1 f p subshell with 3 orbitals f p subshell with 3 orbitals For s = 2, m s = -2, -1, 0, +1, +2 f d subshell with 5 orbitals f d subshell with 5 orbitals For s = 3, m s = -3, -2, -1, 0, +1, +2, +3 f f subshell with 7 orbitals f f subshell with 7 orbitals

67. 67 © 2009 Brooks/Cole - Cengage f Orbitals One of 7 possible f orbitals. All have 3 planar surfaces. Can you find the 3 surfaces here?

68. 68 © 2009 Brooks/Cole - Cengage Spherical Nodes Orbitals also have spherical nodesOrbitals also have spherical nodes Number of spherical nodes = n - s - 1Number of spherical nodes = n - s - 1 For a 2s orbital: No. of nodes = 2 - 0 - 1 = 1For a 2s orbital: No. of nodes = 2 - 0 - 1 = 1 2 s orbital

69. 69 © 2009 Brooks/Cole - Cengage Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (s) ORBITALS (m s )

70. 70 © 2009 Brooks/Cole - Cengage Each orbital can be assigned no more than 2 electrons! This is tied to the existence of a 4th quantum number, the electron spin quantum number, m s. Arrangement of Electrons in Atoms

71. 71 © 2009 Brooks/Cole - Cengage Electron Spin Quantum Number, m s Can be proved experimentally that electron has an intrinsic property referred to as “spin.” Two spin directions are given by m s where m s = +1/2 and -1/2. Can be proved experimentally that electron has an intrinsic property referred to as “spin.” Two spin directions are given by m s where m s = +1/2 and -1/2. PLAY MOVIE

72. 72 © 2009 Brooks/Cole - Cengage Electron Spin and Magnetism Diamagnetic : NOT attracted to a magnetic fieldDiamagnetic : NOT attracted to a magnetic field Paramagnetic : substance is attracted to a magnetic field.Paramagnetic : substance is attracted to a magnetic field. Substances with unpaired electrons are paramagnetic.Substances with unpaired electrons are paramagnetic. PLAY MOVIE

73. 73 © 2009 Brooks/Cole - Cengage Measuring Paramagnetism Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons. Diamagnetic : NOT attracted to a magnetic field See Active Figure 6.18

74. 74 © 2009 Brooks/Cole - Cengage n f shell1, 2, 3, 4,... s f subshell0, 1, 2,... n - 1 m s f orbital - s... 0... + s m s f electron spin+1/2 and -1/2 n f shell1, 2, 3, 4,... s f subshell0, 1, 2,... n - 1 m s f orbital - s... 0... + s m s f electron spin+1/2 and -1/2 QUANTUM NUMBERS Now there are four!