1. Radiant Energy travels through space is energy that travels through space. light Is also known as light and electromagnetic radiation electromagnetic radiation. Major source is… THE SUN

2. Radiant Energy… Sun’s radiant energy is the result of nuclear fusion. Nuclear fusion – light nuclei combine to form heavier nuclei Fission vs. fusion?

3. The dual nature of light Particle??? Wave??? The dual nature – light can be viewed as a wave (continuous) OR a stream of extremely tiny, fast- moving particles (quantized)

4. The dual nature of light Particle??? Wave??? - Wave (continuous) as it travels through space - Particle (quantized) as it interacts with matter

5. Electromagnetic Spectrum ordered sequence light electromagnetic radiation The ordered sequence of all types of light or electromagnetic radiation.

6. The part of the electromagnetic spectrum that humans can see is called the … radio, micro, radar, IR, vis. Electromagnetic Spectrum vis., UV, X-ray, gamma, cosmic Low, low & Long! High, high & short!

7. Violet- High energy, bends more, inside of rainbow, 400 nm R O Y G B I V Red- Low energy, bends less, outside of rainbow, 700 nm

8. Sir Isaac Newton …in the 1670’s, diffracted light with a prism and… Concluded that each color of light has a unique wavelength… energy

9. Electromagnetic Spectrum h a

10. amplitude (a) - affects brightness, half the height Frequency (ν) = number of crests passing a point in a period of time h = height, from crest to trough speed (c) = distance per unit time wavelength (λ) = crest to crest

11. λ = c/ν wavelength speed of light frequency

12. If a light wave has a wavelength (λ) of 3.0 * 10 -7 m, what is its frequency? wavelength speed of light frequency C = 3 * 10 8 m/s 1.0 * 10 15 Hz

13. Quantum Theory Beginning of 20 th century – wave model is almost universally accepted. Problem = electromagnetic radiation emitted from hot objects

14. Can all light be described as continuous The energy of waves is continuous, or unbroken… When we look at an object as it is heated, what do we see? Quantum Theory … - quanta

15. Each color has its own energy And the energy changes with heating Quantum Theory … - quanta

16. Max Planck energy - related the frequency of light to its energy with the following: E = h Energyh=Planck’s const. E=Energy, h=Planck’s const. specific amounts His idea was that energy is absorbed and released in specific amounts. Quantum Theory 6.626e-34 J*s

17. specific amounts quantum quanta atoms His idea was that energy is absorbed and released in specific amounts. He called one piece, package, or bundle of energy one quantum. Bundles of energy were called quanta. He applied his quanta ideas to energy changes in atoms: Quantum Theory

18. atoms quantized continuous (wave-like) He applied his quanta ideas to energy changes in atoms: The energy of atoms is quantized. Formerly, scientists had thought that all energy was continuous (wave-like). Quantum Theory

19. Democritus’ atom Quantized Matter:Energy:Plank

20. Continuous: VS. Quantized: Height by step or rung Height on slide or ramp

21. Continuous: VS. Quantized: Ice cream scoops Soft-serve ice cream

22. Quantized: VS. Continuous: escalator stairs

23. Continuous: VS. Quantized: Cheese singles Cheese block

24. Quantum Theory Come up with your own: Continuous VS. Quantized: AND put it in your notes!

25. Albert Einstein - Imagined that light energy traveled in bundles - photons. 18 years later, Arthur Compton experimentally demonstrated that light is comprised of tiny particles, or photons, Quantum Theory

26. TODAY - Planck’s term quantum and Einstein’s term photon are used interchangeably. demonstrated that light is comprised of tiny particles, or photons, that can collide with electrons and cause them to move. Quantum Theory

27. TODAY - Planck’s term quantum and Einstein’s term photon are used interchangeably. Scientists also believe that light has properties of both waves and particles. Quantum Theory … - Dual nature of light

28. The relationship between frequency , wavelength ( ), and color to the energy of light:  , Color: red = low E, violet = high E E  : many photon punches E  : big gaps between consecutive photons.

29. Color: red = low E, violet = high E each color has its own energy color = type of light

30. The Effects of Different Photons: Microwaves: can’t feel Infrared: feel with skin, warms or burns. Visible light: see with eyes, heats when absorbed. Ultraviolet: can’t feel or see, affects cells – freckles, tan, burn, cataracts.

31. X-rays: can’t feel or see, pass through body, but absorbed by bones and dense matter. Gamma rays: can’t see or feel, affects cells, causes mutations in cells and molecules. The Effects of Different Photons: Ultraviolet: can’t feel or see, affects cells – freckles, tan, burn, cataracts.

32. Neils Bohr (bright) line spectrum H Tried to explain why each element has its own unique (bright) line spectrum. He studied H. specific energy levels orbitshell nucleus Using previous discoveries- Bohr hypothesized that an atom’s electrons are located in specific energy levels. Each energy level, aka orbit or shell is a set distance from the atom’s nucleus. …

33. Bohr’s Hypothesis In the line spectrum of an atom, Bohr saw specific colors.In the line spectrum of an atom, Bohr saw specific colors. Each specific color has a specific energy.Each specific color has a specific energy. That specific amount of energy is related to a specific distance from the nucleus.That specific amount of energy is related to a specific distance from the nucleus.

34. Energy Source Absorbed Energy Energy Released

35. Ground vs. Excited States: ground closest stable An atom is in the ground state when its electrons fill the lowest possible energy levels that are closest to the nucleus. This is when the atom is most stable. gain jump exact amount An electron can gain energy and jump to a higher energy level. The electron must absorb an exact amount …

36. gain jump exact amount photon An electron can gain energy and jump to a higher energy level. The electron must absorb an exact amount … of energy to make a jump to a specific energy level. The energy that the electron gains comes from a photon. Ground vs. Excited States:

37. excited stable When an atom’s electrons are in higher energy levels, the atom is in an excited state and is less stable. fall released visibleor invisible light The atom prefers to be stable, so the electrons fall into lower energy levels that are not full. As the electrons fall, energy is released in the form of visible or invisible light.

38. atoms prefer… to be stable! to have low energy! to be in their ground state!

39. Energy within the atom? Increases away from the nucleus ENERGY

40. Quantum Mechanics Mr. Bohr was concerned with calculating and predicting the line spectra of elements. What happens when there is more than 1 electron?

41. Quantum Mechanics Mr. Bohr was concerned with calculating and predicting the line spectra of elements. He wondered how electrons move and where they can be found in atoms. Bohr’s ideas worked well for hydrogen with 1 electron. … What happens when there is more than 1 electron?

42. Quantum Mechanics Bohr’s ideas worked well for hydrogen with 1 electron. … He predicted the infrared and ultraviolet bands of hydrogen’s emission spectrum. The equations he used came from Classical Mechanics, a branch of physics that describes the movements and interactions that are large enough to see.

43. But… Alas.. Bohr could not predict the bright-line spectra. The laws of Classical Mechanics just don’t cut it for atoms and electrons.

44. Electrons are tricky… they and other subatomic particles like them have their own code of conduct… They behave differently than anything you may be able to see with your eyes or with any other object. New ideas needed to be looked into, and these new ideas became known as Quantum Mechanics.

45. Spectroscopes

46. Louis de Broglie wave properties Planck’slight One of the first to think that electrons possess wave properties. He reasoned that since waves can act as particles do (taken from Planck’s idea about light), then particles might behave as waves do.

47. Large moving objects Wavelengths are small and practically unnoticed.

48. For tiny subatomic particles… are important increases atom Wave properties are important. As the size of the moving object decreases, its wavelength increases. The wavelength for a tiny electron can be as large as an entire atom.

49. So how does an electron move in an atom? circularspherical Bohr (and maybe you too…) thought that they moved in circular or spherical orbits. matter-wave idea With de Broglie’s matter-wave idea, now we theorize that electrons vibrate around the nucleus in a.

50. The Elusive Electron Evades Subatomic State Trooper! Werner Heisenberg Uncertainty Principle In 1927, he proposed the Uncertainty Principle speedlocation This states that it is impossible to know both the speed and location of an electron at the same time.

51. Why is it so hard to pinpoint the electron? To determine the speed and the location of an object, you must be able to SEE the object… light is bounced off the object when you see it. Light is made up of quanta or photons.

52. When photons hit a speeding car, the car is unaffected. But when a photon hits a speeding electron, the electron will move or change direction. So, if a photon hits an electron and the light bounces off it into your eyes, you will see where the electron was, but you won’t know how fast it was going at the time.

53. Heisenberg Explain the Heisenberg Uncertainty Principle. speedlocation It is impossible to know both the speed and location of an electron at the same time. What, am I speeding?

54. The New Atom…. location speed If you cannot know the exact location and speed of the electron, what is a scientist to do? region areas/zones of probability With the use of calculus, the region where the electron is most likely to be can be determined. These regions are called areas/zones of probability

55. pattern overlap standing The most likely location of an electron is described by a wave of probability. This type of wave is actually a set pattern that forms a 3-D shape within the space of the atom. This wave pattern does not overlap itself and is known as a standing wave. http://www.youtube.com/watch?v=-gr7KmTOrx0 http://www.youtube.com/watch?v=3BN5-JSsu_4 http://www.youtube.com/watch?v=18BL7MKjtZM

56. Let’s now return to Bohr’s atomic model…

57. specific energy levels Bohr said that electrons are found in specific energy levels in an atom. Each energy level is a circle or sphere with a definite radius. …

58. average cloud Each energy level is a circle or sphere with a definite radius. … Bohr was close: What he thought as definite is actually the average radius. With quantum mechanics, the model of the energy level has expanded from a specific sphere to a region of probability that is like a cloud around the nucleus.

60. Revisiting the Electron Cloud Electron Position: probability orbitals uncertainty Electron DistributionProperties of element

61. How can the whereabouts of an electron compare to an apartment building? Floor/story of building Suite/apartment Bedroom Female / Male Energy level Sublevel Orbital for e - couple Spin

62. The Address of the Electrons Just like people in an apartment, electrons have an address. The most probable location of an electron is described using quantum numbers. Each electron has 4 quantum numbers which each relays a different piece of information about the electron’s possible whereabouts in the atom.

63. Pauli exclusion- no 2 electrons can have the same address!!! – same 4 quantum numbers The Address of the Electrons

64. Energy levels - break into - Describing Atomic Structure Sublevels - break into - Orbitals Energy level – regions of space where there is a high probability of finding electrons

65. 1 2 3 First sublevel – 2 electrons Second sublevel – 6 electrons Third sublevel – 10 electrons Describing Atomic Structure Sublevels

66. Describing Atomic Structure – Building an Address Energy levels – designated by - # Sublevels – designated by – letter (s, p, d, f) 1 st Sublevel = s 2 nd Sublevel = p 3 rd Sublevel = d 4 th Sublevel = f 1 st Level = 1, 2 nd Level = 2, etc. …

67. Describing Atomic Structure # of energy level = # of sublevels in that energy level Energy level 1 = 1 sublevel Energy level 2 = 2 sublevels Energy level 3 = 3 sublevels …

68. Describing Atomic Structure Sublevels Energy still Increases away from the nucleus E n e r g y fdpsfdps

69. The New Atom? Do you think Bohr was right or wrong? Pauli exclusion? Energy levels - break into - Sublevels - break into - Orbitals

70. How do we define the “address” of the electron within the atom?

72. 1 2 3 Energy level 1, sublevel s, orbital 1s Energy level 2, has two sublevels s and p, 2s orbital and 2p orbitals Energy level 3, 3 sublevels s, p, and d, 3s orbital, 3p orbitals, and 3d orbitals Electron Arrangement Page 9

73. Electron Arrangement Quantum #s – specify the properties of atomic orbitals and the properties of electrons in those orbitals Principle quantum # (n) – main energy level = 1, 2, 3… Angular momentum quantum # (l) – shape of the orbital = 0, 1, 2… (n-1) 0 = s, 1 = p, 2 = d, 3 = f Magnetic quantum # (m l ) – orientation of the orbital around the nucleus

74. Energy levels – designated by - # (n the principle quantum #) Sublevels – designated by – letter (s, p, d, f) (l the angular momentum quantum #) s = 0 p = 1 d = 2 f = 3 E n e r g y Energy still Increases away from the nucleus Electron Arrangement

75. 2 electrons fit in an orbital one spinning in the +1/2 orientation and one spinning in the -1/2 orientation Electron Arrangement How many electrons in an orbital?

76. How many orbitals in each sublevel? Electron Arrangement s = 1 p = 3 d = 5 f = 7

77. How many sublevels in each energy level? Electron Arrangement 1 = _____ (__) 2 = _____ (____) 3 = _____ (______) 4 = _____ (__________) 1 2 3

78. xx 43214321 Energy Level xx s p d f Each x = an electron Each xx = an orbital Electron Arrangement

79. Taking a Look at Orbitals S P D

80. Recall that atoms like to stay in their most stable (lowest energy) state = electron configuration Electron Arrangement electron configuration – notation used to show electron placement within sublevels Skip ahead to page 11

81. Recall that atoms like to stay in their most stable (lowest energy) state. 1 2 3 - Sublevels fill from the nucleus outward Handouts Electron Arrangement

82. electron configuration – notation used to show electron placement within orbitals electron Configuration for: Si C 1s 2 2s 2 2p 6 3s 2 3p 2 1s 2 2s 2 2p 2 Valence electrons?

83. Electron Configurations Practice: S F 1s 2 2s 2 2p 6 3s 2 3p 4 Valence electrons? 1s 2 2s 2 2p 5

84. Taking a Look at Orbitals

85. Why does the ring model work? p – orbitals always occur in 3’s (one for each dimension) 3 rd and 6 th stopped

86. Taking a Look at Orbitals 3 rd and 6 th start

87. Why do exceptions exist? How can you explain the exceptions?

88. Electron Configurations using noble gas notation for: K Fe [Ar]4s 1 [Ar]4s 2 3d 6 Cd Sn [Kr]4d 10 5s 2 [Kr]4d 10 5s 2 5p 2

89. 1 2 3 Energy level 1, sublevel s, orbital 1s Energy level 2, has two sublevels s and p, 2s orbital and 2p orbitals Energy level 3, 3 sublevels s, p, and d, 3s orbital, 3p orbitals, and 3d orbitals

90. State Is energy released or absorbed?

91. xx 43214321 Energy Level xx s p d f Each x = an electron Each xx = an orbital Electron Arrangement In energy level 1 there is one s sublevel which contains one orbital and therefore 2 electrons. In energy level 2 there is an s sublevel and a p sublevel, the s sublevel contains one orbital (2 electrons) and the p sublevel contains 3 orbitals (6 electrons) for a total of 8 electrons. In energy level 3 …

92. electron Configuration for: Si C 1s 2 2s 2 2p 6 3s 2 3p 2 1s 2 2s 2 2p 2 Valence electrons?

93. Light energy hits the electrons in metal- the light must be powerful enough. The electrons become excited, and they jump out of the metal. Electrons in the metal absorb the energy. Quantum Theory Quanta’s Ability:

94. The electrons fall down again, and create a spark or current. Examples: The luster of a shiny metal, Photoelectric cells (solar power) The electrons become excited, and they jump out of the metal. Quantum Theory