1. Chapter 7 Atomic Structure and Periodicity

2. Chapter 7 Table of Contents Copyright © Cengage Learning. All rights reserved 2 7.1 Electromagnetic Radiation 7.2 The Nature of Matter 7.3 The Atomic Spectrum of Hydrogen 7.4 The Bohr Model 7.5 The Quantum Mechanical Model of the Atom 7.6Quantum Numbers 7.7Orbital Shapes and Energies 7.8Electron Spin and the Pauli Principle 7.9Polyelectronic Atoms 7.10The History of the Periodic Table 7.11The Aufbau Principle and the Periodic Table 7.12Periodic Trends in Atomic Properties 7.13The Properties of a Group: The Alkali Metals

3. Section 7.1 Electromagnetic Radiation Return to TOC Copyright © Cengage Learning. All rights reserved 3 Different Colored Fireworks

4. Section 7.1 Electromagnetic Radiation Return to TOC Copyright © Cengage Learning. All rights reserved 4 Questions to Consider Why do we get colors? Why do different chemicals give us different colors?

5. Section 7.1 Electromagnetic Radiation Return to TOC Copyright © Cengage Learning. All rights reserved 5 Electromagnetic Radiation One of the ways that energy travels through space. Three characteristics:  Wavelength  Frequency  Speed

6. Section 7.1 Electromagnetic Radiation Return to TOC Copyright © Cengage Learning. All rights reserved 6 Characteristics Wavelength ( ) – distance between two peaks or troughs in a wave. Frequency ( ) – number of waves (cycles) per second that pass a given point in space Speed (c) – speed of light (2.9979×10 8 m/s)

7. Section 7.1 Electromagnetic Radiation Return to TOC Copyright © Cengage Learning. All rights reserved 7 The Nature of Waves

8. Section 7.1 Electromagnetic Radiation Return to TOC Copyright © Cengage Learning. All rights reserved 8 Classification of Electromagnetic Radiation

9. Section 7.2 The Nature of Matter Return to TOC Copyright © Cengage Learning. All rights reserved 9 Pickle Light

10. Section 7.2 The Nature of Matter Return to TOC Copyright © Cengage Learning. All rights reserved 10 Energy can be gained or lost only in integer multiples of. A system can transfer energy only in whole quanta (or “packets”). Energy seems to have particulate properties too.

11. Section 7.2 The Nature of Matter Return to TOC Copyright © Cengage Learning. All rights reserved 11 Energy is quantized. Electromagnetic radiation is a stream of “particles” called photons.

12. Section 7.2 The Nature of Matter Return to TOC Copyright © Cengage Learning. All rights reserved 12 The Photoelectric Effect

13. Section 7.2 The Nature of Matter Return to TOC Copyright © Cengage Learning. All rights reserved 13 Energy has mass E = mc 2 Dual nature of light:  Electromagnetic radiation (and all matter) exhibits wave properties and particulate properties.

14. Section 7.3 The Atomic Spectrum of Hydrogen Return to TOC Copyright © Cengage Learning. All rights reserved 14 Continuous spectrum (when white light is passed through a prism) – contains all the wavelengths of visible light Line spectrum – each line corresponds to a discrete wavelength:  Hydrogen emission spectrum

15. Section 7.3 The Atomic Spectrum of Hydrogen Return to TOC Copyright © Cengage Learning. All rights reserved 15 Refraction of White Light

16. Section 7.3 The Atomic Spectrum of Hydrogen Return to TOC Copyright © Cengage Learning. All rights reserved 16 The Line Spectrum of Hydrogen

17. Section 7.3 The Atomic Spectrum of Hydrogen Return to TOC Copyright © Cengage Learning. All rights reserved 17 Significance Only certain energies are allowed for the electron in the hydrogen atom. Energy of the electron in the hydrogen atom is quantized.

18. Section 7.3 The Atomic Spectrum of Hydrogen Return to TOC Copyright © Cengage Learning. All rights reserved 18 Concept Check Why is it significant that the color emitted from the hydrogen emission spectrum is not white? How does the emission spectrum support the idea of quantized energy levels?

19. Section 7.4 The Bohr Model Return to TOC Copyright © Cengage Learning. All rights reserved 19 Electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. Bohr’s model gave hydrogen atom energy levels consistent with the hydrogen emission spectrum. Ground state – lowest possible energy state (n = 1)

20. Section 7.4 The Bohr Model Return to TOC Copyright © Cengage Learning. All rights reserved 20 Electronic Transitions in the Bohr Model for the Hydrogen Atom a) An Energy-Level Diagram for Electronic Transitions

21. Section 7.4 The Bohr Model Return to TOC Copyright © Cengage Learning. All rights reserved 21 Electronic Transitions in the Bohr Model for the Hydrogen Atom b) An Orbit-Transition Diagram, Which Accounts for the Experimental Spectrum

22. Section 7.4 The Bohr Model Return to TOC Copyright © Cengage Learning. All rights reserved 22 For a single electron transition from one energy level to another: ΔE = change in energy of the atom (energy of the emitted photon) n final = integer; final distance from the nucleus n initial = integer; initial distance from the nucleus

23. Section 7.4 The Bohr Model Return to TOC Copyright © Cengage Learning. All rights reserved 23 The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electron. As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state (free electron). As the electron is brought closer to the nucleus, energy is released from the system.

24. Section 7.4 The Bohr Model Return to TOC Copyright © Cengage Learning. All rights reserved 24 Bohr’s model is incorrect. This model only works for hydrogen. Electrons do not move around the nucleus in circular orbits.

25. Section 7.4 The Bohr Model Return to TOC Copyright © Cengage Learning. All rights reserved 25 Exercise What color of light is emitted when an excited electron in the hydrogen atom falls from: a)n = 5 to n = 2 b)n = 4 to n = 2 c)n = 3 to n = 2 Which transition results in the longest wavelength of light? blue, λ = 434 nm green, λ = 486 nm orange/red, λ = 657 nm

26. Section 7.5 The Quantum Mechanical Model of the Atom Return to TOC Copyright © Cengage Learning. All rights reserved 26 We do not know the detailed pathway of an electron. Heisenberg uncertainty principle:  There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. Δx = uncertainty in a particle’s position Δ(mν) = uncertainty in a particle’s momentum h = Planck’s constant

27. Section 7.5 The Quantum Mechanical Model of the Atom Return to TOC Copyright © Cengage Learning. All rights reserved 27 Physical Meaning of a Wave Function The square of the function indicates the probability of finding an electron near a particular point in space.  Probability distribution – intensity of color is used to indicate the probability value near a given point in space.

28. Section 7.5 The Quantum Mechanical Model of the Atom Return to TOC Copyright © Cengage Learning. All rights reserved 28 Probability Distribution for the 1s Wave Function

29. Section 7.5 The Quantum Mechanical Model of the Atom Return to TOC Copyright © Cengage Learning. All rights reserved 29 Radial Probability Distribution

30. Section 7.5 The Quantum Mechanical Model of the Atom Return to TOC Copyright © Cengage Learning. All rights reserved 30 Relative Orbital Size Difficult to define precisely. Orbital is a wave function. Picture an orbital as a three-dimensional electron density map. Hydrogen 1s orbital:  Radius of the sphere that encloses 90% of the total electron probability.

31. Section 7.6 Quantum Numbers Return to TOC Copyright © Cengage Learning. All rights reserved 31 Principal quantum number (n) – size and energy of the orbital. Angular momentum quantum number (l) – shape of atomic orbitals (sometimes called a subshell). Magnetic quantum number (m l ) – orientation of the orbital in space relative to the other orbitals in the atom.

32. Section 7.6 Quantum Numbers Return to TOC Copyright © Cengage Learning. All rights reserved 32 Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom

33. Section 7.6 Quantum Numbers Return to TOC Copyright © Cengage Learning. All rights reserved 33 Exercise For principal quantum level n = 3, determine the number of allowed subshells (different values of l), and give the designation of each. # of allowed subshells = 3 l = 0, 3s l = 1, 3p l = 2, 3d

34. Section 7.6 Quantum Numbers Return to TOC Copyright © Cengage Learning. All rights reserved 34 Exercise For l = 2, determine the magnetic quantum numbers (m l ) and the number of orbitals. magnetic quantum numbers = –2, – 1, 0, 1, 2 number of orbitals = 5

35. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 35 1s Orbital

36. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 36 Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals

37. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 37 2p x Orbital

38. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 38 2p y Orbital

39. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 39 2p z Orbital

40. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 40 The Boundary Surface Representations of All Three 2p Orbitals

41. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 41 3d x 2 -y 2 Orbital

42. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 42 3d xy Orbital

43. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 43 3d xz Orbital

44. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 44 3d yz Orbital

45. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 45 Orbital

46. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 46 The Boundary Surfaces of All of the 3d Orbitals

47. Section 7.7 Orbital Shapes and Energies Return to TOC Copyright © Cengage Learning. All rights reserved 47 Representation of the 4f Orbitals in Terms of Their Boundary Surfaces

48. Section 7.8 Electron Spin and the Pauli Principle Return to TOC Copyright © Cengage Learning. All rights reserved 48 Electron Spin Electron spin quantum number (m s ) – can be +½ or -½. Pauli exclusion principle - in a given atom no two electrons can have the same set of four quantum numbers. An orbital can hold only two electrons, and they must have opposite spins.

49. Section 7.9 Polyelectronic Atoms Return to TOC Copyright © Cengage Learning. All rights reserved 49 Atoms with more than one electron. Electron correlation problem:  Since the electron pathways are unknown, the electron repulsions cannot be calculated exactly. When electrons are placed in a particular quantum level, they “prefer” the orbitals in the order s, p, d, and then f.

50. Section 7.9 Polyelectronic Atoms Return to TOC Copyright © Cengage Learning. All rights reserved 50 Penetration Effect A 2s electron penetrates to the nucleus more than one in the 2p orbital. This causes an electron in a 2s orbital to be attracted to the nucleus more strongly than an electron in a 2p orbital. Thus, the 2s orbital is lower in energy than the 2p orbitals in a polyelectronic atom.

51. Section 7.9 Polyelectronic Atoms Return to TOC Copyright © Cengage Learning. All rights reserved 51 Orbital Energies

52. Section 7.9 Polyelectronic Atoms Return to TOC Copyright © Cengage Learning. All rights reserved 52 A Comparison of the Radial Probability Distributions of the 2s and 2p Orbitals

53. Section 7.9 Polyelectronic Atoms Return to TOC Copyright © Cengage Learning. All rights reserved 53 The Radial Probability Distribution of the 3s Orbital

54. Section 7.9 Polyelectronic Atoms Return to TOC Copyright © Cengage Learning. All rights reserved 54 A Comparison of the Radial Probability Distributions of the 3s, 3p, and 3d Orbitals

55. Section 7.10 The History of the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 55 Originally constructed to represent the patterns observed in the chemical properties of the elements. Mendeleev is given the most credit for the current version of the periodic table.

56. Section 7.11 The Aufbau Principle and the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 56 Aufbau Principle As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen–like orbitals. An oxygen atom as an electron arrangement of two electrons in the 1s subshell, two electrons in the 2s subshell, and four electrons in the 2p subshell. Oxygen: 1s 2 2s 2 2p 4

57. Section 7.11 The Aufbau Principle and the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 57 Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals.

58. Section 7.11 The Aufbau Principle and the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 58 Orbital Diagram A notation that shows how many electrons an atom has in each of its occupied electron orbitals. Oxygen: 1s 2 2s 2 2p 4 Oxygen: 1s 2s 2p

59. Section 7.11 The Aufbau Principle and the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 59 Valence Electrons The electrons in the outermost principal quantum level of an atom. 1s 2 2s 2 2p 6 (valence electrons = 8) The elements in the same group on the periodic table have the same valence electron configuration.

60. Section 7.11 The Aufbau Principle and the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 60 The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

61. Section 7.11 The Aufbau Principle and the Periodic Table Return to TOC Copyright © Cengage Learning. All rights reserved 61 Exercise Determine the expected electron configurations for each of the following. a) S 1s 2 2s 2 2p 6 3s 2 3p 4 or [Ne]3s 2 3p 4 b) Ba [Xe]6s 2 c) Eu [Xe]6s 2 4f 7

62. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Periodic Trends Ionization Energy Electron Affinity Atomic Radius

63. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Ionization Energy Energy required to remove an electron from a gaseous atom or ion.  X(g) → X + (g) + e – Mg → Mg + + e – I 1 = 735 kJ/mol(1 st IE) Mg + → Mg 2+ + e – I 2 = 1445 kJ/mol(2 nd IE) Mg 2+ → Mg 3+ + e – I 3 = 7730 kJ/mol*(3 rd IE) *Core electrons are bound much more tightly than valence electrons.

64. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Ionization Energy In general, as we go across a period from left to right, the first ionization energy increases. Why?  Electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons.  Electrons in the same principal quantum level are generally more strongly bound from left to right on the periodic table.

65. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Ionization Energy In general, as we go down a group from top to bottom, the first ionization energy decreases. Why?  The electrons being removed are, on average, farther from the nucleus.

66. Section 7.12 Periodic Trends in Atomic Properties Return to TOC The Values of First Ionization Energy for the Elements in the First Six Periods

67. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Concept Check Explain why the graph of ionization energy versus atomic number (across a row) is not linear. electron repulsions Where are the exceptions? some include from Be to B and N to O

68. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Concept Check Which atom would require more energy to remove an electron? Why? Na Cl

69. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Concept Check Which atom would require more energy to remove an electron? Why? Li Cs

70. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Concept Check Which has the larger second ionization energy? Why? Lithium or Beryllium

71. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Successive Ionization Energies (KJ per Mole) for the Elements in Period 3

72. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Electron Affinity Energy change associated with the addition of an electron to a gaseous atom.  X(g) + e – → X – (g) In general as we go across a period from left to right, the electron affinities become more negative. In general electron affinity becomes more positive in going down a group.

73. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Photoelectron Spectroscopy Technique that is used to gather information about the electrons in an atom. An atom is bombarded with photons. Some of the photons are absorbed and electrons are emitted. The electrons are collected and their energy is analyzed. Since we can know the energy of the photons, and we know that energy is conserved we know that the difference in energy between the photons sent into the atom and the energy of the electrons emitted will be the potential energy of the electrons when they are attached to the atom. Remember that the potential energy of the electron in the atom is the work needed to remove the electron from the atom. Energy of emitted electron = energy of photon - work needed to remove electron from atom Copyright © Cengage Learning. All rights reserved 73

74. Section 7.12 Periodic Trends in Atomic Properties Return to TOC PES apparatus: iramis.cea.fr

75. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Photoelectron Spectroscopy How it works: 1.Sample is exposed to EM radiation 2.Electrons jump out of sample and go through analyzer http://chemwiki.ucdavis.edu

76. Section 7.12 Periodic Trends in Atomic Properties Return to TOC PES Data Energy to remove an electron (binding energy) (increases to the left!)  Number of electrons Electrons generally farther from the nucleus Electrons generally closer to the nucleus The bigger the peak – the more electrons Each peak represents the electrons in a single sublevel in the atom

77. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Hydrogen vs. Helium Hydrogen Helium The helium peak is twice as tall because there are twice as many electrons in the 1s sublevel 1 electron in 1s 2 electrons in 1s #e-  energy #e-  energy

78. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Hydrogen vs. Helium Hydrogen Helium The helium peak is farther to the left (higher energy) thus more energy is needed to remove the 1s electrons in helium. They must be held more tightly because there is a higher effective nuclear charge. (Helium has 2 protons pulling on 1s but hydrogen only has 1) 1 electron in 1s 2 electrons in 1s #e-  energy #e-  energy

79. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Oxygen (1s 2 2s 2 2p 4 ) Energy to remove an electron (binding energy) (increases to the left!)  Number of electrons 2 electrons in 1s 2 electrons in 2s4 electrons in 2p

80. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Scandium (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 ) Energy to remove an electron (binding energy) (increases to the left!)  Number of electrons 2 in1s 2 in 2s 2 in 4s 6 in 2p 2 in 3s 6 in 2p 1 in 3d *Notice that it takes more energy to remove an electron from 3d than from 4s. This is because as electrons are added to 3d they shield 4s thus it’s easier (takes less energy) to remove 4s electrons compared to 3d electrons. Remember when transition metals make positive ions - it’s the s electrons that are lost first!

81. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Example 1: Energy  Number of electrons Identify the element whose PES data is shown to the right. Sodium Why is one peak much larger Than the others? This peak represents 6 electrons In the 2p sublevel the other Peaks represent only 1 or 2 electrons In which sublevel are the electrons Represented by peak A 3s A

82. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Example 2: Nitrogen Oxygen #e-  energy #e-  energy The PES data above shows only the peak for the 1s electrons. Why is the peak for Nitrogen farther to the left? It takes less energy to remove a 1s electron from nitrogen because it has a lower Effective nuclear charge (less protons) than oxygen

83. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Example 3: Energy  Number of electrons Draw the expected PES Spectrum for the element boron

84. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Atomic Radius In general as we go across a period from left to right, the atomic radius decreases.  Effective nuclear charge increases, therefore the valence electrons are drawn closer to the nucleus, decreasing the size of the atom. In general atomic radius increases in going down a group.  Orbital sizes increase in successive principal quantum levels.

85. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Atomic Radii for Selected Atoms

86. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Concept Check Which should be the larger atom? Why? Na Cl

87. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Concept Check Which should be the larger atom? Why? Li Cs

88. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Concept Check Which is larger? The hydrogen 1s orbital The lithium 1s orbital Which is lower in energy? The hydrogen 1s orbital The lithium 1s orbital

89. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Atomic Radius of a Metal

90. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Atomic Radius of a Nonmetal

91. Section 7.12 Periodic Trends in Atomic Properties Return to TOC Exercise Arrange the elements oxygen, fluorine, and sulfur according to increasing:  Ionization energy S, O, F  Atomic size F, O, S

92. Section 7.13 The Properties of a Group: The Alkali Metals Return to TOC Copyright © Cengage Learning. All rights reserved 92 The Periodic Table – Final Thoughts 1.It is the number and type of valence electrons that primarily determine an atom’s chemistry. 2.Electron configurations can be determined from the organization of the periodic table. 3.Certain groups in the periodic table have special names.

93. Section 7.13 The Properties of a Group: The Alkali Metals Return to TOC Copyright © Cengage Learning. All rights reserved 93 Special Names for Groups in the Periodic Table

94. Section 7.13 The Properties of a Group: The Alkali Metals Return to TOC Copyright © Cengage Learning. All rights reserved 94 The Periodic Table – Final Thoughts 4.Basic division of the elements is into metals and nonmetals.

95. Section 7.13 The Properties of a Group: The Alkali Metals Return to TOC Copyright © Cengage Learning. All rights reserved 95 Metals Versus Nonmetals

96. Section 7.13 The Properties of a Group: The Alkali Metals Return to TOC Copyright © Cengage Learning. All rights reserved 96 The Alkali Metals Li, Na, K, Rb, Cs, and Fr  Most chemically reactive of the metals  React with nonmetals to form ionic solids  Going down group:  Ionization energy decreases  Atomic radius increases  Density increases  Melting and boiling points smoothly decrease