2. This slide contains classified material and cannot be shown to high school students. Please continue as if everything is normal.

3. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

4. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

5. Rutherford’s Gold Foil Experiment  Alpha particles are helium nuclei  Particles were fired at a thin sheet of gold foil  Particle hits on the detecting screen (film) are recorded

7. The Puzzle of the Atom  Protons and electrons are attracted to each other because of opposite charges  Electrically charged particles moving in a curved path give off energy  Despite these facts, atoms don’t collapse

8. c = C = speed of light, a constant (3.00 x 10 8 m/s) = frequency, in units of hertz (hz, sec -1 ) = wavelength, in meters Electromagnetic radiation propagates through space as a wave moving at the speed of light.

9. Types of electromagnetic radiation:

10. Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table

11. The Wave-like Electron Louis deBroglie The electron propagates through space on an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Toupee?

12. Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital. Sizes of s orbitals

13. Orbitals in outer energy levels DO penetrate into lower energy levels. This is a probability Distribution for a 3s orbital. What parts of the diagram correspond to “nodes” – regions of zero probability? Penetration #1

14. The Great The Great Niels Bohr (1885 - 1962)

15. …produces all of the colors in a continuous spectrum Spectroscopic analysis of the visible spectrum…

16. …produces a “bright line” spectrum Spectroscopic analysis of the hydrogen spectrum…

17. This produces bands of light with definite wavelengths. Electron transitions involve jumps of definite amounts of energy.

18. Bohr Model Energy Levels

19. Schrodinger Wave Equation probability Equation for probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger

20. Heisenberg Uncertainty Principle You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg

21. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it.  Principal quantum number  Angular momentum quantum number  Magnetic quantum number  Spin quantum number (n) (l) (m) (s)

22. Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located.  The principal quantum number (n) cannot be zero.  n must be 1, 2, 3, etc. Number of electrons that can fit in a shell: 2n 2

23. Angular Momentum Quantum Number The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located. The angular momentum quantum number (l ) can be any integer between 0 and n - 1. l =3 f

24. Orbital shapes are defined as the surface that contains 90% of the total electron probability. An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…

25. Magnetic Quantum Number The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space. The magnetic quantum number (m l ) can be any integer between -l and +l.

26. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli

27. Spin Quantum Number Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin:

28. Assigning the Numbers  The three quantum numbers (n, l, and m) are integers.  The principal quantum number (n) cannot be zero.  n must be 1, 2, 3, etc.  The angular momentum quantum number (l ) can be any integer between 0 and n - 1.  For n = 3, l can be either 0, 1, or 2.  The magnetic quantum number (m l ) can be any integer between -l and +l.  For l = 2, m can be either -2, -1, 0, +1, +2.

29. Principle, angular momentum, and magnetic quantum numbers: n, l, and m l

30. Aufbau

31. 1. energy is emitted 2. energy is absorbed 3. no change in energy occurs 4. light is emitted 5. none of these

32. 1. gamma rays 2. microwaves 3. radio waves 4. infrared radiation 5. x-rays

33. 1. 2 2. 5 3. 10 4. 18 5. 6

34. nlms 1. 110½ 2. 300–½ 3. 21–1½ 4. 43–2–½ 5. 420 ½

35. 1. 3.00 x 10 13 s–1 2. 4.12 x 10 5 s–1 3. 8.50 x 10 20 s–1 4. 9.12 x 10 12 s–1 5. 3.20 x 10 9 s–1

36. 1. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 5d 10 2. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4d 10 4p 1 3. 1s 2 3s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 6 5s 2 5d 10 5p 1 4. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 5. none of these

37. Orbital filling table

38. Yet Another Way to Look at Ionization Energy

39. ElementConfiguration notation Orbital notationNoble gas notation Lithium1s 2 2s 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 1 Beryllium1s 2 2s 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 Boron1s 2 2s 2 p 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 1 Carbon1s 2 2s 2 p 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 2 Nitrogen1s 2 2s 2 p 3 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 3 Oxygen1s 2 2s 2 p 4 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 4 Fluorine1s 2 2s 2 p 5 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 5 Neon1s 2 2s 2 p 6 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 6

41. The s orbital has a spherical shape centered around the origin of the three axes in space. s orbital shape

42. There are three peanut-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space. P orbital shape

43. Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of: …and a “peanut with a donut”! d orbital shapes “double peanut”

44. Shape of f orbitals Things get even more complicated with the seven f orbitals that are found in the f sublevels beginning with n = 4. To remember the shapes, think of: Flower

45. Electron configuration of the elements of the first three series

46. Periodicity

47.  Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

48. Influenced by three factors. Energy Level Higher energy level is further away. Charge on nucleus More charge pulls electrons in closer. Shielding Layers of electrons shield from nuclear pull.

49. The electron on the outside energy level has to look through all the other energy levels to see the nucleus

50. The electron on the outside energy level has to look through all the other energy levels to see the nucleus. A second electron has the same shielding.

51. As we go down a group Each atom has another energy level, So the atoms get bigger. H Li Na K Rb

52. As you go across a period the radius gets smaller. Same energy level. More nuclear charge. Outermost electrons are closer. NaMgAlSiPSClAr

53. Table of Atomic Radii

54. Cations form by losing electrons. Cations are smaller that the atom they come from. Metals form cations. Cations of representative elements have noble gas configuration.

55. Anions form by gaining electrons. Anions are bigger that the atom they come from. Nonmetals form anions. Anions of representative elements have noble gas configuration.

57. Atomic Number Atomic Radius (nm) H Li Ne Ar 10 Na K Kr Rb

59. The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a +1 ion. The energy required is called the first ionization energy.

60. The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron.

61. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

62. The greater the nuclear charge the greater IE. Distance from nucleus increases IE Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. Shielding

63. As you go down a group first IE decreases because The electron is further away. More shielding.

64. All the atoms in the same period have the same energy level. Same shielding. Increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 fill orbitals.

65. First Ionization energy Atomic number H He Li Be B C N O F Ne l Na has a lower IE than Li l Both are s 1 l Na has more shielding l Greater distance Na

66.  Affinity tends to increase across a period  Affinity tends to decrease as you go down in a group Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals Electron Affinity - the energy change associated with the addition of an electron

68. The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How fair it shares. Big electronegativity means it pulls the electron toward it. Atoms with large negative electron affinity have larger electronegativity.

69. The further down a group the farther the electron is away and the more electrons an atom has. More willing to share. Low electronegativity.

70. Metals are at the left end. They let their electrons go easily Low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away. High electronegativity.

71. Ionization energy, electronegativity Electron affinity INCREASE

72. Atomic size increases, shielding constant Ionic size increases

73. Another Way to Look at Ionization Energy

74. Yet Another Way to Look at Ionization Energy

76. Summary of Periodic Trends

77. Put the following in order of Decreasing atomic radius: a)Cl,Ar,K b)b) O, O -, O 2- c) Co, Rh, Ni Now put them in order of Decreasing ionization energy:

79. Prepare yourself to ^ C

80. c = C = speed of light, a constant (3.00 x 10 8 m/s) = frequency, in units of hertz (hz, sec -1 ) = wavelength, in meters Electromagnetic radiation propagates through space as a wave moving at the speed of light.

81. Types of electromagnetic radiation:

83. E = h E = Energy, in units of Joules (kg·m 2 /s 2 ) h = Planck’s constant (6.626 x 10-34 J·s) = frequency, in units of hertz (hz, sec -1 ) = frequency, in units of hertz (hz, sec -1 ) The energy (E ) of electromagnetic radiation is directly proportional to the frequency ( ) of the radiation.

84. Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table

85. Relating Frequency, Wavelength and Energy Common re-arrangements:

87. Types of electromagnetic radiation:

88. PLANCK’S PRACTICE PROBLEMS 1.When we see light from a neon sign, we are observing radiation from excited neon atoms. If this radiation has a wavelength of 640 nm, what is the energy of the photon being emitted? 2. Light with a wavelength of 614.5 nm looks orange. What is the energy, in joules, of a photon of this orange light? 3. A photon of light produced by a surgical laser has an energy of 3.027 x 10 -19 J. Calculate the frequency and the wavelength of the photon.

91. method that provides information on all the occupied energy levels of an atom (that is, the ionization energies of all electrons in the atom) is known as photoelectron spectroscopy; this method uses a photon (a packet of light energy) to knock an electron out of an atom.

92. The photoelectron spectrum is a plot of the number of electrons emitted versus their kinetic energy. In the diagram below, the “X” axis is labeled high to low energies so that you think about the XY intersect as being the nucleus.   http://www.chem.arizona.edu/chemt/Flash/photoelectron.html

94. Interpretations from the data: 1. There are no values on the y axis in the tables above. Using the Periodic Table and Table 1, put numbers on the y axis. 2. Label each peak on the graphs above with s, p, d, or f to indicate the suborbital they represent.. 3. What is the total number of electrons in a neutral potassium atom? Orbital names s, p, d, and f stand for names given to groups of lines in the spectra of the alkali metals. Early chemists called the line groups sharp, principal, diffuse, and fundamental. 1- 2- 6- 1s2s3s 4s 2p3p

95. If a certain element being studied by an X-ray PES displays an emission spectrum with 5 distinct kinetic energies. What are all the possible elements that could produce this spectrum? Determine the orbitals that the spectral lines are originating from and then determine the elements that have electrons in only these orbitals.