Chapter 5: Periodicity and Atomic Structure

1. Chapter 5: Periodicity and Atomic Structure

2. Chapter Outline Introduction: Periodicity and the Periodic Table Part 1:History and Bohr Light & Electromagnetic Radiation Particle-like Nature of Light Bohr Model of the Atom Wave-like Nature of Matter Part 2:Quantum Mechanical Model of the Atom orbitals, quantum numbers, energy diagrams Electron Configurations and the periodic table Part 3: Periodic Properties of the Elements Z effective Atomic radii Ionization Energy (Chapter 6) Electron Affinity (Chapter 6)

3. Periodicity

4. Development of the Periodic Table

5. The Electromagnetic Spectrum

6. The Wave Nature of Light Classically, light consists of oscillating electric and magnetic fields.

7. Electromagnetic Radiation The number of crests that pass a given point per second is the frequency, ?.

8. Wavelength and Frequency CO2 absorbs light with a wavelength of 0.018 mm. Which frequency range is this?

9. Blackbody Radiation

10. Blackbody Radiation

11. Blackbody Radiation

12. Ultraviolet Catastrophe Ultraviolet radiation has a wavelength that is shorter than visible light Classical Physics: energy density of electromagnetic radiation increases with decreasing wavelength. Question 1: Why would this lead to a catastrophe? Question 2: What�s going on? Why do you think electromagnetic radiation does not increase with decreasing wavelength according to the classical laws of Physics?

13. Quantized Energy and Photons Planck: energy can only be absorbed or released from atoms in certain amounts called quanta. The relationship between energy and frequency is E = h ? where h is Planck�s (6.626 ? 10-34 J.s). To understand quantization consider walking up a ramp versus walking up stairs: For the ramp, there is a continuous change in height whereas up stairs there is a quantized change in height.

14. Questions for You 1. Infrared or UV light� Which has the higher frequency? Which has the longer wavelength? Which has the greatest energy? 2. What is the frequency of a microwave with ? = 4.33 x 10-3 m

15. The Photoelectric Effect Many metals emit electrons when irradiated by light (see Animation) Again the results of these experiments deviate from classical physics

16. Energy of a Light Wave Classically, the energy of a wave depends on its amplitude. The photoelectric effect however shows that the energy of light depends on its frequency.

17. Conceptual Question From the ChemQuizzes CD-ROM by David Laws and Alexander Pines � 2003 W. W. Norton & Company Inc. All rights reserved.From the ChemQuizzes CD-ROM by David Laws and Alexander Pines � 2003 W. W. Norton & Company Inc. All rights reserved.

18. Answer: AAnswer: A

19. The Photoelectric Effect The minimum energy for ejection of electrons from cesium is 3.05?10-19 J. Will green light with ? = 505 nm cause electrons to be ejected?

20. Concept Building Check Wave Nature of Light Particle Nature of Light Plank-energy is quantized Einstein�The Photoelectric Effect

21. Bohr Model of the Hydrogen Atom Bohr Proposed: Energies around atomic nuclei are also quantized The Electron in the hydrogen atom must reside in one of an array of discrete energy states, or �allowed� energy levels. Animation Used Line Spectrum of Hydrogen to support his argument

22. Bohr Model of the Hydrogen Atom

23. Bohr Model of the Hydrogen Atom

24. Bohr Model of the Hydrogen Atom

25. Conceptual Question From the ChemQuizzes CD-ROM by David Laws and Alexander Pines � 2003 W. W. Norton & Company Inc. All rights reserved.From the ChemQuizzes CD-ROM by David Laws and Alexander Pines � 2003 W. W. Norton & Company Inc. All rights reserved.

26. Answer: CAnswer: C

27. Conceptual Question From the ChemQuizzes CD-ROM by David Laws and Alexander Pines � 2003 W. W. Norton & Company Inc. All rights reserved.From the ChemQuizzes CD-ROM by David Laws and Alexander Pines � 2003 W. W. Norton & Company Inc. All rights reserved.

28. The Bohr Theory-Calculations 1. The energy of a hydrogen atom is quantized with energies given by

29. The Bohr Theory What wavelength of light is emitted during a transition from n = 4 to n = 2?

30. What frequency of light is absorbed during a transition from n=3 to n=5? More Problem Solving

31. How does the QM Model Differ from Bohr�s Model? Animation

32. Quantum Mechanics De Broglie�s hypothesis and Heisenberg�s Uncertainty Principle set the stage for a new theory of atomic structure. In the new model, the wave nature of the electron is recognized and its behavior is described in terms of waves This new model precisely describes the energy of the electron while describing its location in terms of probabilities

33. Quantum Mechanics In 1925, Schr�dinder described the hydrogen atom in terms of a wave function ?: H? = E? Solving this equation yields a set of wave functions {?n} with energies {En}. Born interpreted ?n2 as the probability of locating the electron in state n at a given point in space. See Orbital Viewer

34. The Hydrogen Atom The Quantum Mechanical Model gives us a set of three quantum numbers to describe an orbital (n, l, ml) ?n,l,m2 maps out the probability of locating the electron. We refer to these functions as orbitals.

35. Atomic Quantum Numbers Principle quantum number n = 1,2,3,.... The energy of an orbital depends only on n:

36. From the Bohr Model

37. Atomic Quantum Numbers Azimuthal quantum number l = 0,1,2,..., n-1 l determines the shape of an orbital. The value of l is designated by a letter: l = 0 , 1 , 2 , 3 , � s p d f 1. n = 1 l = ? 2. n = 2 l = ?

38. Atomic Quantum Numbers Magnetic quantum number ml = -l, -l+1, ... , l-1, l ml determines the orientation of an orbital.

39. Shells and Subshells n l ml orbital energy 1 0 0 1s -RH 2 0 0 2s -RH/4 2 1 -1,0,1 2p -RH/4 3 0 0 3s -RH/9 3 1 -1,0,1 3p -RH/9 3 2 -2,-1,0,1,2 3d -RH/9

40. Orbital Representations Electron density plots show the probability of locating the electron.

41. Orbital Representations A contour diagram is a surface that encloses most (say 90%) of the probability density. The s orbitals are spherical:

42. Orbital Representations Electron density plots show the probability of locating the electron.

43. Orbital Representations A contour diagram is a surface that encloses most (say 90%) of the probability density. The s orbitals are spherical:

44. p Orbitals p orbitals (l = 1) have two lobes lying along the x, y, or z axis. Rather than ml = -1,0,1, the orbitals are labelled px, py, and pz.

45. d and f Orbitals These orbitals have complicated shapes, but the electron density at the nucleus is always zero.

46. Animation n = 2, l = 1, ml = -1 n = 4, l = 1, ml = 0

47. Electron Spin In 1928, it was discovered that an electron has an intrinsic angular momentum, or spin. In a magnetic field, the rotation axis has only two possible orientations.

48. Concept Question: Which of the following sets of quantum numbers are allowed for an electron in an atom? n l ml ms1) 2 1 0 +1/22) 3 0 +1 -1/23) 3 2 -2 -1/24) 1 1 0 +1/25) 2 1 0 0 a) � 2, 4b) � 1, 3c) � 3, 4d) � 1, 2, 3e) � 2, 4, 5

49. Multi-electron Atoms

50. The Helium Atom Consider a two-electron atom with nuclear charge Z.

51. Effective Nuclear Charge We account for electron repulsion by assuming that the electrons shield each other from the nuclear charge. The net nuclear charge experienced by an electron is the effective nuclear charge, Zeff. If S is the average number of screening electrons: Zeff = Z - S

52. Electron Energies Due to screening, different subshells have different energies, increasing in the order: s < p < d < f

53. Conceptual Question Identify the subshell in which electrons with the quantum numbers n = 6, l = 1 may be found. a) � 5pb) � 6dc) � 6pd) � 6fe) � 3d

54. Conceptual Question What is the lowest-numbered principal shell in which f orbitals are found? 2 1 4 5 3

55. The Exclusion Principle How many electrons can fit in, or �occupy� an orbital? The Pauli Exclusion principle states: No two electrons in an atom can have the same four quantum numbers. The ground state of helium has two electrons in the 1s orbital, but with opposite spins. n l ml ms electron 1 1 0 0 +� electron 2 1 0 0 -�

56. Many-Electron Atoms The Aufbau principle: Electrons are assigned, one at a time, to hydrogen orbitals with lowest possible energy. An orbital diagram shows the number of electrons in each occupied orbital.

57. Hund�s Rule Hund�s rule: The lowest energy state has the most unpaired electrons. Carbon:

58. Write Energy diagrams for each atom in first two rows

59. Conceptual Question Which of the following electron diagrams represents a correct ground state?�

60. Electron Configurations and the Periodic Table The electron configuration of an atom can be estimated from the Periodic table. The actual configuration must be determined by experiment.

61. Electron Configurations and the Periodic Table Write electron configurations for:

62. Valence Shell/Electrons Valence shell: outer most energy level/shell that contains electrons in an atom Valence Electrons: electrons in the outermost shell Core electrons: Inner electrons (non-valence electrons) Examples C Rb Br

63. Activity Create a model of an O atom. Use all that you�ve learned about modern atomic structure including the location of the subatomic particles, orbital drawings (3D), orbital energy diagram, electron configuration, quantum numbers, valence electrons, core electrons, etc� What other atoms in the periodic table have a similar electron configuration to O?

64. Electron Configurations and the Periodic Table The electron configuration of an atom can be estimated from the Periodic table. The actual configuration must be determined by experiment.

65. Elements within a group have similar electron configurations in their valence electron shells Alkali metals 1s1 2s1 3s1 4s1

66. Nobel Gases Filled valence shell � HAPPY!

68. Periodic Properties of the Elements Early versions of the Periodic table were constructed by Mendeleev and Meyer. We now know that the periodic properties are due to the electronic structure of atoms. Electronic structure explains the observed trends in Atomic size Ionization Energy Electron Affinity �

69. Effective Nuclear Charge Consider a two-electron atom with nuclear charge Z.

70. Effective Nuclear Charge We account for electron repulsion by assuming that the electrons shield each other from the nuclear charge. The net nuclear charge experienced by an electron is the effective nuclear charge, Zeff. If S is the average number of screening electrons: Zeff = Z - S

71. Effective Nuclear Charge As the distance from the nucleus increases, S increases and Zeff decreases. As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases. � The effective nuclear charge experienced by the outer electrons is determined primarily by the difference between the charge on the nucleus and the charge of the core electrons.

74. Atomic Size The radius of an atom is found from the distance between nuclei in a molecule.

75. Trends in Atomic Size Size increases going down a column of the periodic table. Size decreases from left to right in a row. There are two factors at work: principal quantum number, n, and the effective nuclear charge, Zeff.

76. Trends in Atomic Size In a column, size increases with the addition of successive electron shells.

78. You Try� �Which set of elements is not in order of increasing atomic radius (smallest one first, etc.)?� Be, Al, Mg C, N, O Br, Se, As F, S, As Na, K, Rb

79. You Try� True or false? The stronger the nucleus attracts the outer orbital electrons, the larger the atom.