1. Chapter 10 Modern Atomic Theory

2. Greek Idea l Democritus and Leucippus l Matter is made up of indivisible particles l Dalton - one type of atom for each element

3. Thomson’s Model l Discovered electrons l Atoms were made of positive stuff l Negative electron floating around l “Plum-Pudding” model

4. Rutherford’s Model l Discovered dense positive piece at the center of the atom l Nucleus l Electrons moved around l Mostly empty space

5. Bohr’s Model l Why don’t the electrons fall into the nucleus? l Move like planets around the sun. l In circular orbits at different levels. l Amounts of energy separate one level from another.

6. Bohr’s Model Nucleus Electron Orbit Energy Levels

7. Bohr’s Model Increasing energy Nucleus First Second Third Fourth Fifth } l Further away from the nucleus means more energy. l There is no “in between” energy l Energy Levels

8. The Quantum Mechanical Model l Energy is quantized. It comes in chunks. l A quanta is the amount of energy needed to move from one energy level to another. l Since the energy of an atom is never “in between” there must be a quantum leap in energy. l Schrodinger derived an equation that described the energy and position of the electrons in an atom

9. l Things that are very small behave differently from things big enough to see. l The quantum mechanical model is a mathematical solution l It is not like anything you can see. The Quantum Mechanical Model

10. l Has energy levels for electrons. l Orbits are not circular. l It can only tell us the probability of finding an electron a certain distance from the nucleus. The Quantum Mechanical Model

11. l The atom is found inside a blurry “electron cloud” l A area where there is a chance of finding an electron. l Draw a line at 90 % The Quantum Mechanical Model

12. Atomic Orbitals l Principal Quantum Number (n) = the energy level of the electron. l Within each energy level the complex math of Schrodinger’s equation describes several shapes. l These are called atomic orbitals l Regions where there is a high probability of finding an electron.

13. l 1 s orbital for every energy level l Spherical shaped l Each s orbital can hold 2 electrons l Called the 1s, 2s, 3s, etc.. orbitals. S orbitals

14. P orbitals l Start at the second energy level l 3 different directions l 3 different shapes l Each can hold 2 electrons

15. P Orbitals

16. D orbitals l Start at the second energy level l 5 different shapes l Each can hold 2 electrons

17. F orbitals l Start at the fourth energy level l Have seven different shapes l 2 electrons per shape

18. F orbitals

19. Summary s p d f # of shapes Max electrons Starts at energy level 121 362 5103 7144

20. By Energy Level l First Energy Level l only s orbital l only 2 electrons l 1s 2 l Second Energy Level l s and p orbitals are available l 2 in s, 6 in p l 2s 2 2p 6 l 8 total electrons

21. By Energy Level l Third energy level l s, p, and d orbitals l 2 in s, 6 in p, and 10 in d l 3s 2 3p 6 3d 10 l 18 total electrons l Fourth energy level l s,p,d, and f orbitals l 2 in s, 6 in p, 10 in d, ahd 14 in f l 4s 2 4p 6 4d 10 4f 14 l 32 total electrons

22. By Energy Level l Any more than the fourth and not all the orbitals will fill up. l You simply run out of electrons l The orbitals do not fill up in a neat order. l The energy levels overlap l Lowest energy fill first.

23. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

24. Electron Configurations l The way electrons are arranged in atoms. l Aufbau principle- electrons enter the lowest energy first. l This causes difficulties because of the overlap of orbitals of different energies. l Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

25. Electron Configuration l Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to. l Let’s determine the electron configuration for Phosporus l Need to account for 15 electrons

26. l The first to electrons go into the 1s orbital l Notice the opposite spins l only 13 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

27. l The next electrons go into the 2s orbital l only 11 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

28. The next electrons go into the 2p orbital only 5 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

29. The next electrons go into the 3s orbital only 3 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

30. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into seperate shapes 3 upaired electrons 1s 2 2s 2 2p 6 3s 2 3p 3

31. The easy way to remember 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2 electrons

32. Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 4 electrons

33. Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 12 electrons

34. Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 20 electrons

35. Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 38 electrons

36. Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 56 electrons

37. Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 88 electrons

38. Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6 108 electrons

39. Exceptions to Electron Configuration

40. Orbitals fill in order l Lowest energy to higher energy. l Adding electrons can change the energy of the orbital. l Half filled orbitals have a lower energy. l Makes them more stable. l Changes the filling order

41. Write these electron configurations l Titanium - 22 electrons l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 l Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 l Chromium - 24 electrons l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 is expected l But this is wrong!!

42. Chromium is actually l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 l Why? l This gives us two half filled orbitals. l Slightly lower in energy. l The same principal applies to copper.

43. Copper’s electron configuration l Copper has 29 electrons so we expect l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 l But the actual configuration is l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 l This gives one filled orbital and one half filled orbital. l Remember these exceptions

44. Light l The study of light led to the development of the quantum mechanical model. l Light is a kind of electromagnetic radiation. l Electromagnetic radiation includes many kinds of waves l All move at 3.00 x 10 8 m/s ( c)

45. Parts of a wave Wavelength Amplitude Orgin Crest Trough

46. Parts of Wave l Orgin - the base line of the energy. l Crest - high point on a wave l Trough - Low point on a wave l Amplitude - distance from origin to crest l Wavelength - distance from crest to crest Wavelength - is abbreviated  Greek letter lambda.

47. Frequency l The number of waves that pass a given point per second. l Units are cycles/sec or hertz (hz) Abbreviated  the Greek letter nu c =

48. Frequency and wavelength l Are inversely related l As one goes up the other goes down. l Different frequencies of light is different colors of light. l There is a wide variety of frequencies l The whole range is called a spectrum

49. Radio waves Micro waves Infrared. Ultra- violet X- Rays Gamma Rays Low energy High energy Low Frequency High Frequency Long Wavelength Short Wavelength Visible Light

50. Atomic Spectrum How color tells us about atoms

51. Prism l White light is made up of all the colors of the visible spectrum. l Passing it through a prism separates it.

52. If the light is not white l By heating a gas with electricity we can get it to give off colors. l Passing this light through a prism does something different.

53. Atomic Spectrum l Each element gives off its own characteristic colors. l Can be used to identify the atom. l How we know what stars are made of.

54. These are called discontinuous spectra Or line spectra unique to each element. These are emission spectra The light is emitted given off.

55. Light is a Particle l Energy is quantized. l Light is energy l Light must be quantized l These smallest pieces of light are called photons. l Energy and frequency are directly related.

56. Energy and frequency E = h x n l E is the energy of the photon n is the frequency l h is Planck’s constant l h = 6.6262 x 10 -34 Joules sec. l joule is the metric unit of Energy

57. The Math in Chapter 11 l Only 2 equations c = ln E = h n l Plug and chug.

58. Examples l What is the wavelength of blue light with a frequency of 8.3 x 10 15 hz? l What is the frequency of red light with a wavelength of 4.2 x 10 -5 m? l What is the energy of a photon of each of the above?

59. An explanation of Atomic Spectra

60. Where the electron starts l When we write electron configurations we are writing the lowest energy. l The energy level and electron starts from is called its ground state.

61. Changing the energy l Let’s look at a hydrogen atom

62. Changing the energy l Heat or electricity or light can move the electron up energy levels

63. Changing the energy l As the electron falls back to ground state it gives the energy back as light

64. l May fall down in steps l Each with a different energy Changing the energy

65. { { {

66. l Further they fall, more energy, higher frequency. l This is simplified l the orbitals also have different energies inside energy levels l All the electrons can move around. Ultraviolet Visible Infrared

67. What is light l Light is a particle - it comes in chunks. l Light is a wave- we can measure its wave length and it behaves as a wave If we combine E=mc 2, c=ln, E = 1/2 mv 2 and E = hn We can get l = h/mv l The wavelength of a particle.

68. Matter is a Wave l Does not apply to large objects l Things bigger that an atom l A baseball has a wavelength of about 10 - 32 m when moving 30 m/s l An electron at the same speed has a wavelength of 10 - 3 cm l Big enough to measure.

69. The physics of the very small l Quantum mechanics explains how the very small behaves. l Classic physics is what you get when you add up the effects of millions of packages. l Quantum mechanics is based on probability because

70. Heisenberg Uncertainty Principle l It is impossible to know exactly the speed and velocity of a particle. l The better we know one, the less we know the other. l The act of measuring changes the properties.

71. More obvious with the very small l To measure where a electron is, we use light. l But the light moves the electron l And hitting the electron changes the frequency of the light.

72. Moving Electron Photon Before Electron Changes velocity Photon changes wavelength After