1. Chapter 5 “Electrons in Atoms” Pre-AP Chemistry Charles Page High School Stephen L. Cotton

2. Ernest Rutherford’s Model l Discovered dense positive piece at the center of the atom- “nucleus” l Electrons would surround and move around it, like planets around the sun l Atom is mostly empty space l It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

3. Niels Bohr’s Model l Why don’t the electrons fall into the nucleus? l Move like planets around the sun.  In specific circular paths, or orbits, at different levels.  An amount of fixed energy separates one level from another.

4. The Bohr Model of the Atom Niels Bohr I pictured the electrons orbiting the nucleus much like planets orbiting the sun. However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another.

5. Bohr’s model l Energy level of an electron analogous to the rungs of a ladder l The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder l A quantum of energy is the amount of energy required to move an electron from one energy level to another

6. The Quantum Mechanical Model l Energy is “quantized” - It comes in chunks. l A quantum 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 In 1926, Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

7. Schrodinger’s Wave Equation probability Equation for the probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger

8. 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 (like plum pudding!) The Quantum Mechanical Model

9. 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

10. l The atom is found inside a blurry “electron cloud” l An area where there is a chance of finding an electron. l Think of fan blades The Quantum Mechanical Model

11. Quantum Mechanical Model Quantum mechanics was developed by Erwin Schrodinger Estimates the probability of finding an e - in a certain position Electrons are found in an “electron cloud” or orbital

12. Radial Distribution Curve Orbital Orbital (“electron cloud”) – Region in space where there is 90% probability of finding an e -

13. Each orbital letter has a different shape.

14. “s” orbital spherical shaped, and holds up to 2e-

15. “p” orbital Dumbbell shaped Arranged x, y, z axes, and can hold up to 6e-

16. “d” orbital clover shaped, and can hold up to 10e-

17. “f” orbital f Orbitals combine to form a spherical shape. This orbital can hold up to 14e- 2s 2p z 2p y 2p x

18. Atomic Orbitals l Principal Quantum Number (n) = the energy level of the electron: 1, 2, 3, etc. l Within each energy level, the complex math of Schrodinger’s equation describes several shapes. l These are called atomic orbitals (coined by scientists in 1932) - regions where there is a high probability of finding an electron. l Sublevels- like theater seats arranged in sections: letters s, p, d, and f

19. Principal Quantum Number Generally symbolized by “n”, it denotes the shell (energy level) in which the electron is located. Maximum number of electrons that can fit in an energy level is: 2n 2 How many e - in level 2? 3?

20. Summary s p d f # of shapes (orbitals) Maximum electrons Starts at energy level 1 2 1 3 6 2 5 10 3 7 14 4

21. Hog Hilton You are the manager of a prestigious new hotel in downtown Midland—the “Hog Hilton”. It’s just the “snort of the town” and you want to keep its reputation a cut above all the other hotels. Your problem is your clientele. They are hogs in the truest sense. Your major task is to fill rooms in your hotel. The Hog Hilton only has stairs. You must fill up your hotel keeping the following rules in mind: 1) Hogs are lazy, they don’t want to walk up stairs! 2) Hogs want to room by themselves, but they would rather room with another hog than walk up more stairs. 3) If hogs are in the same room they will face in opposite directions. 4) They stink, so you can’t put more than two hogs in each room.

22. Hog Hilton Your hotel looks like the diagram below: 6th floor ______ 5th floor ______ ______ ______ 4th floor ______ 3rd floor ______ ______ ______ 2nd floor ______ 1st floor ______ Book 7 hogs into the rooms.

23. Hog Hilton Your hotel looks like the diagram below: 6th floor ______ 5th floor ______ ______ ______ 4th floor ______ 3rd floor ______ ______ ______ 2nd floor ______ 1st floor ______ Book 14 hogs into the rooms.

24. Let’s play Hog Hilton!!

25. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f aufbau diagram - page 133 Aufbau is German for “building up”

26. Rules for e - configurations Aufbau principle 1. Aufbau principle: e - enter orbitals of lowest energy level ( Hogs are lazy, they don’t want to walk up stairs!) Pauli exclusion principle 2. Pauli exclusion principle: an atomic orbital may have at most 2 e -, e - in the same orbital will spin in opposite directions ( They stink, so you can’t put more than two hogs in each room. & If hogs are in the same room they will face in opposite directions.) Hund’s rule 3. Hund’s rule: when e - occupy orbitals of = energy, 1 enters each orbital until all the orbitals contain 1 e - w/parallel spins ( Hogs want to room by themselves, but they would rather room with another hog than walk up more stairs.)

27. B. Within the energy level are sublevels, designated by letters. Principle energy level (n) Number of sublevels Type of Orbital 1 st energy level1 sublevel“s” (1 orbital) 2 nd 2 sublevels“s” (1) & “p” (3 orbitals) 3 rd 3 sublevels“s”(1), “p” (3) & “d” (5 orbitals) 4 th 4 sublevels“s”(1), “p”(3), “d”(5), and “f” (7)

28. 1s 2s 2p 3p 3s 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p7s 7s 7p 6s 6p 6d 6f 6g 5s 5p 5d 5f 5g 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s

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

30. 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

31. 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

32. 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

33. 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 separate shapes (Hund’s) 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3 Orbital notation

34. l An internet program about electron configurations is: Electron Configurations (Just click on the above link)

35. Orbitals fill in an order l Lowest energy to higher energy. l Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. l However, half filled orbitals have a lower energy, and are next best Makes them more stable. Changes the filling order

36. Write the electron configurations for these elements: lTlTitanium - 22 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 lVlVanadium - 23 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 lClChromium - 24 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 (expected) BBut this is not what happens!!

37. 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 (the others are all still full) l Half full is slightly lower in energy. l The same principal applies to copper.

38. Copper’s electron configuration l Copper has 29 electrons so we expect: 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 change gives one more filled orbital and one that is half filled. l Remember these exceptions: d 4, d 9

39. Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel

40. Section 3

41. 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 types: gamma rays, x-rays, radio waves… l Speed of light = 2.998 x 10 8 m/s, and is abbreviated “c” l All electromagnetic radiation travels at this same rate when measured in a vacuum

42. - Page 139 “R O Y G B I V” Frequency Increases Wavelength Longer

43. Parts of a wave Wavelength Amplitude Origin Crest Trough

44. Equation: c = c = speed of light, a constant (2.998 x 10 8 m/s) (nu) = frequency, in units of hertz (hz or sec -1 ) (lambda) = wavelength, in meters Electromagnetic radiation propagates through space as a wave moving at the speed of light.

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

46. - Page 140 Use Equation: c =

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

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

49. Atomic Spectra l White light is made up of all the colors of the visible spectrum. l Passing it through a prism separates it.

50. 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.

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

52. These are called the atomic emission spectrum Unique to each element, like fingerprints! Very useful for identifying elements

53. Light is a Particle? l Energy is quantized. l Light is a form of energy. l Therefore, light must be quantized l These smallest pieces of light are called photons. l Photoelectric effect? Albert Einstein l Energy & frequency: directly related.

54. Equation: E = h E = Energy, in units of Joules (kg·m 2 /s 2 ) (Joule is the metric unit of energy) (Joule is the metric unit of energy) 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.

55. The Math in Chapter 5 l There are 2 equations: 1) c = 2) E = h Know these!

56. Examples 1) What is the wavelength of blue light with a frequency of 8.3 x 10 15 hz? 2) What is the frequency of red light with a wavelength of 4.2 x 10 -5 m? 3) What is the energy of a photon of each of the above?

57. Explanation of atomic spectra l When we write electron configurations, we are writing the lowest energy. l The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level.

58. Changing the energy l Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

59. Changing the energy l Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be “ excited ”

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

61. Experiment #6, page 49-

62. l They may fall down in specific steps l Each step has a different energy Changing the energy

63. { { {

64. l The further they fall, more energy is released and the higher the frequency. l This is a simplified explanation! l The orbitals also have different energies inside energy levels l All the electrons can move around. Ultraviolet Visible Infrared

65. What is light? l Light is a particle - it comes in chunks. l Light is a wave - we can measure its wavelength and it behaves as a wave If we combine E=mc 2, c=, E = 1/2 mv 2 and E = h, then we can get: = h/mv (from Louis de Broglie) l called de Broglie’s equation l Calculates the wavelength of a particle.

66. Wave-Particle Duality J.J. Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

67. Confused? You’ve Got Company! “No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.” Physicist Sir Arthur Eddington The Nature of the Physical World 1934

68. The physics of the very small l Quantum mechanics explains how very small particles behave Quantum mechanics is an explanation for subatomic particles and atoms as waves l Classical mechanics describes the motions of bodies much larger than atoms

69. Heisenberg Uncertainty Principle l It is impossible to know exactly the location and velocity of a particle. l The better we know one, the less we know the other. l Measuring changes the properties. l True in quantum mechanics, but not classical mechanics

70. 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

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

72. Moving Electron Photon Before Electron velocity changes Photon wavelength changes After Fig. 5.16, p. 145