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Light, Energy, And More


chapter 5 notes

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  • 1. Light, Energy, and More October 23, 2007 Chemistry
  • 2. Recap…
    • Electromagnetic Spectrum
    • High Energy
    • Low Energy
    • Wave Nature of Light
  • 3.  
  • 4. What’s Going On Here?
  • 5. When we heat metal what happens?
  • 6.
    • Does the wave model of light explain these changes?
      • Does not explain different wavelengths and frequencies at different temperatures
    • What is light?
      • Radiation….what is radiation?
        • Particles or rays of energy
    • What is temperature anyways?
      • The measure of the average kinetic energy of the particles in an object
      • Kinetic Energy vs. Potential Energy
    • Too many questions….
  • 7. Max Planck (1900)
    • German Physicist
    • Began to look for answers
    • Matter can only gain or lose energy in small quantized amounts
    • What’s quantized?
  • 8. Vocab Word!!!
      • Minimum amount of energy that can be gained or lost by an atom
      • The emitted light from a glowing metal is a ENERGY…this energy is quantized
  • 9. If energy is now quantized…how can we determine the amount of energy of a quantum?
    • What is energy measured in?
    • What are we observing?
    • What happens to the color when we increase the temperature (energy)?
      • Proportional or inversely proportional?
    • Now we need a constant…
      • Planck’s constant, h=6.626 x 10^-34 J*s
  • 10.  
  • 11. Time to put these words into action!
    • What is the frequency and wavelength electromagnetic radiation that emits 1.68 x 10^-17 J of energy? What type of electromagnetic radiation is this?
    • Wavelength= 1.18 x 10^-8 m
    • Ultraviolet radiation
  • 12. Some questions to answer…
    • What is the color we see?
    • What happens to the energy of the radiation when we increase the frequency, v, of the radiation emitted?
    • Iron at room temp…color and E?
    • Iron with a little heat…color and E?
    • Iron with lots of heat…color and E?
  • 13. According to Planck’s Theory…
    • If we have a given v, matter can emit or absorb E only in whole number multiples of hv (1hv, 2hv, 3 hv…)
    • Matter can ONLY have specific amounts of energy
    • Wall of kids building blocks
      • We can only add or take away in increments of whole blocks…we cannot remove half a block
  • 14. The Big Mystery of the 1900’s…
    • The Photoelectric Effect…
      • What caused these color changes in metals???
  • 15. Photoelectric Effect
    • Electrons (photoelectrons) are emitted from a metal’s surface when a light of a certain frequency shines on the surface
    • Certain specific amounts of energy (what’s this called???) needed to knock out electrons from metal atoms.
  • 16.  
  • 17. Albert Einstein (1905)
    • Added onto Planck’s Theory…
    • Called the electron’s emitted, PHOTONS (the little energy packets Planck called quantums)
    • Now… E photon = hv
  • 18.
    • Planck paved the way for the explanation behind the mystery
    • But some one else came into the picture…
  • 19. Now light is not just a wave…
    • Einstein’s Dual Nature of Light
      • Particle and wave characteristics
      • Light is a beam of tiny particles, called photons, acting like a wave
  • 20. NEW WORD!!!
    • Photon
    • A particle of electromagnetic radiation with no mass that carries a quantum of energy
  • 21.  
  • 22. What Einstein added…
    • Energy of a photon has a minimum or threshold value to eject photoelectrons
    • What must happen for the photoelectric effect to occur?
      • Energy of a photon (particle of EM radiation) must have the minimum energy requirement to free the electron from the atom of metal
  • 23. Mystery Solved!
    • No matter how long a light of a certain frequency is shone on metal (intensity), electrons will not be ejected unless the minimum amount of energy is shone.
    • Silver metal
      • Photoelectrons ejected when a light with a frequency of at least 1.14 x 10^15 Hz or greater is used
    • Sodium metal
      • Red light
      • Violet light
  • 24. Revised Planck’s Work…
    • Einstein piggy-backed off of Planck’s Theory and we now have…..
  • 25. Time to do a little work….
    • Tiny water drops in the air disperse the white light of the sun into a rainbow. What is the Energy oa a photon from the violet portion of the rainbow if it has a frequency of 7.23x10^14 Hz?
    • E=4.79 x 10^-19 J
      • Energy in a photon of violet light
  • 26. A couple more… 
    • A photon has an energy of 2.93 x 10^-25 J. What is its frequency? What type of electromagnetic radiation is the photon?
    • V=4.42 x 10^8 Hz
    • TV or FM waves
  • 27. Practice makes perfect… 
    • What is the energy of each photon in the following types of radiation?
      • 6.32 x 10^20 Hz
      • 9.50 x 10^13 Hz
      • 1.05 x 10^16 Hz
    • What types of radiation are each?
    • 4.19 x 10^-13 J gamma or x-ray
    • 6.29 x 10^20 J infrared
    • 6.96 x 10^-18 J ultraviolet
  • 28. How does this work? (Neon Signs)
  • 29. What do we know about neon signs?
    • Electricity is passed through tube full of neon gas
    • Neon atoms in tube absorb this energy
      • What happens when something absorbs energy?
    • Neon atoms in tube become excited
      • Stable of Unstable?
      • What happens when something is unstable?
      • What do we see released energy as?
        • Electromagnetic radiation…visible light!!!
  • 30. EM spectrum
    • What happens when we pass sunlight through a prism?
      • Continuous spectrum of colors
      • ROYGBIV
  • 31. What happens when we pass light from neon gas or hydrogen gas through prism?
    • Separation of colors
    • Discontinuous spectrum
    • This is called…
    • AES of an element is the set of frequencies of the electromagnetic radiation emitted by the atoms of that element
    • Individual lines of color
    • Only certain lines of color appear for certain elements…
      • What does this mean…????
    • Every element has a unique AES
    • Why is this important?
  • 33.  
  • 34.  
  • 35. Hydrogen Atom
    • Why did scientists want to use hydrogen?
      • How many protons?
      • How many electrons?
      • Do you think it is easy to use?
      • Check out the AES of hydrogen gas…
  • 36.  
  • 37. Neils Bohr (1913)
    • Danish Physicist
    • Worked with Rutherford
    • Quantum Model of Hydrogen atom
      • Predicted lines of Hydrogen AES
  • 38.
    • Hydrogen has only one electron but why do we get different colored lines on AES???
      • We get hydrogen atoms excited…
      • Electrons move to excited levels
    • H has certain allowable energy states….
      • The lowest energy state is called the GROUND STATE
  • 39. Bohr’s Hydrogen Orbits…
    • He related H’s energy states to the motion of an electron in an atom
    • Single electron in moves around nucleus in circular orbits
    • Smaller orbit, smaller radius, closer to nucleus means…?
      • Lower energy level
    • Larger orbit, larger radius, farther from the nucleus means…?
      • Higher energy level
  • 40.  
  • 41. Bohr’s Quantum Model
    • Assigned quantum numbers, n, to each orbit
    • Calculated orbits radius
      • Chart on page 127
    • 1 st orbit  n=1 (first energy level)
    • 2 nd orbit  n=2 (second energy level)
    • 3 rd orbit  n=3 (third energy level)
  • 42.  
  • 43.  
  • 44. When we add energy, what happens to electron?
    • Electron excited
    • Moves to next energy level
    • Excited=?
      • unstable
    • What happens when something is unstable?
      • Wants to get back to being stable
      • Releases energy
      • Goes back down to lower energy level
    • Photon is emitted corresponding to the 2 different energy levels associated with the 2 orbits
  • 45.  
    • /_ E= E higher e- orbit - E lower e- orbit =E photon =hv
    • Only certain energies are possible so only certain frequencies, v, of EM radiation are emitted
    • Lets look at the AES of Hydrogen…
  • 47.
    • How many lines are there?
    • So how many different types of radiations are we seeing?
    • There are 4 electron transitions account for lines in the hydrogen spectrum
    • Going from 3 rd orbital to 2 nd orbital…
    • Going from 4 th orbital to 2 nd orbital…
    • Going from 5 th orbital to 2 nd orbital…
    • Going from 6 th orbital to 2 nd orbital…
  • 48. Names for these lines…
    • Balmer Series
      • The 4 visible color lines
      • Electrons that drop into n=2
    • Other electrons transitions not visible
      • Lyman series
        • Ultraviolet light
        • Electrons drop into n=1
      • Paschen series
        • Infrared
        • Electrons drop into n=3
  • 49.  
  • 50. Problems with Bohr’s Model
    • Predicted AES lines of H but not any other elements
    • Did not account for all chemical behavior
    • Big problem…
      • Electrons don’t move in circular orbits
      • Time for a new model…
  • 51.  
  • 52. Louis De Broglie (1924)
    • French physics graduate student
    • Proposed idea that accounted for the fixed energy levels in Bohr’s model
  • 53. If waves can have particle like characteristics, then can particles, such as electrons, have wave like characteristics???
  • 54. What he knew…
    • Electrons have wavelike motion (because it’s a particle)
    • An electron had restricted orbits
    • Each orbit had a fixed radius from the nucleus
    • Are a wide variety of wavelengths, frequencies, and energies possible?
  • 55.  
  • 56.
    • No…there could only be allowed certain possible frequencies, wavelengths, and energies in an atom
    • De Broglie came up with an equation for the wavelength of a particle of mass (m) moving at velocity (v).
  • 57. De Broglie’s Equation
  • 58. What does this equation do?
    • What are we using?
      • Wavelength
      • Planck’s constant
      • Mass of the particle
      • Velocity
    • Tells us that all moving particles have wave-like characteristics
  • 59. Food for thought…
    • Cars?
    • Baseball?
    • Do these have wavelike characteristics? Why or why not?
  • 60.
    • Yes…let’s look at the equation…
    • λ = h
    • mv
    • The car and the baseball do have a velocity and a mass…
    • Using De Broglie’s equation we do get a wavelength for the movement of a baseball and a car…
    • Let’s try the calculation…
  • 61. Problem time…
    • Mass of car= 910 kg
    • Velocity of car= 25m/s
    • What is the wavelength of the moving car?
      • 2.9 x 10^-38 m
    • How big is this?
    • Can we see or measure this wavelength?
      • No, much to small to be detected, even with the most sophisticated equipment
  • 62. Another one…
    • Electron speed= 25 m/s
    • Electron mass= 9.11 x 10^-28 g
    • What is the wavelength of the moving electron?
      • 2.9 x 10^-5 m
    • Do you think we can measure this wavelength and see it?
      • Yes, with the right equipment
  • 63. Practice makes perfect 
    • What is the wavelength of an electron of mass 9.11 x 10-28 kg traveling at a velocity of 2.00 x 108 m/s? (Planck's constant = 6.63 x 10-34 J/Hz.
    • 3.64 x 10-15m.
  • 64.  
  • 65. Werner Heisenberg (1901-1976)
    • German theoretical physicist
    • Drew conclusion from Rutherford, Bohr, and De Broglie’s models
  • 66. Problem with finding the position of an electron
    • Helium balloon in a dark room
    • How would you determine the location of this balloon?
    • Is the balloon going to stay in the same position?
    • Energy transfer
    • What if I gave you a flashlight?
      • What happens when we shine a beam of light on the balloon?
  • 67.
    • Photons from light that reflect off of the balloon reach our eyes and tell us where the balloon is
    • Is there a transfer of energy?
      • How big is the balloon compared to the photons?
    • Can we do the same thing with finding the location of an electron in an atom?
    • Heisenberg focused on the interactions between photons and electrons…
  • 68. Heisenberg Uncertainty Principle
    • It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time
  • 69. Erwin Schrodinger (1926)
    • Austrian physicist
    • Furthered De Broglie’s wave-particle theory
    • Derived equation that treated hydrogen’s electron as a wave
    • Unlike Bohr’s, his fit well with atoms of different elements
  • 70. Quantum Mechanical Model of the atom
  • 71. The Quantum Mechanical Model
    • Similar to Bohr’s…
      • Limits an electron’s energy to certain values
    • Unlike Bohr’s…
      • What did Bohr say about the orbit of an electron around the nucleus?
      • The Quantum Mechanic Model makes no attempt to describe the electron’s path
  • 72.
    • Schrodinger’s wave equation
      • Solutions to equation called wave function
        • Don’t worry about the equation its self…just know the basics….
      • Wave function  probability of finding the electron within a particular volume of space around the nucleus
      • High probability  more likely to occur
      • Low probability  less likely to occur
  • 73. What the wave function tells us
    • The atomic orbital of the electron
      • Atomic orbital  3-D region around nucleus
    • Fuzzy Cloud
    • Density of the cloud at a given point is proportional to the probability of finding the electron at that point
  • 74. New Word
    • Orbital  region of space where there is a 90% probability of finding an electron of a given energy
    • “electron cloud”
  • 75. What did Bohr assign to electron orbitals?
    • Quantum numbers
    • Quantum Mechanical Model does the same…
  • 76.
    • Four Quantum Numbers:
      • Specify the “address” (zip code) of each electron in an atom
  • 77.
    • First number…Principal Quantum Number ( n)
    • Energy level (associated with the electron)
    • Size if orbital
      • Lowest energy level is assigned principle quantum number of 1 (n=1)
        • Ground state
      • What do you think happens as we increase n?
        • Orbital becomes larger
        • Electron spends more time farther away from the nucleus  atom’s energy increases
  • 78. Principle energy levels contain… Energy Sublevels
  • 79.
    • Principle energy level 1  single sublevel
    • Principle energy level 2  two sublevels
    • Principle energy level 3  three sublevels
    • What pattern do you see in the number of sublevels as we move further away from the nucleus?
      • They increase as n increases (the further we get from the nucleus)
  • 80.
    • Electron’s are labeled according to n value
    • In atom’s with more than one electron, two or more electron’s may have the same n value
      • They are in the same “electron shell”
  • 81. Second quantum number Angular Momentum Quantum Number (l)
  • 82.
    • Each value of l corresponds to a different type of orbital with a different shape
    • Value of n controls l (subshells possible)
    • Angular momentum numbers can equal 0, 1, 2, 3…
    • l=n-1
      • When n=1, l=0  only one possible subshell
      • When n=2, l=0,1  two possible subshells
  • 83. What the number of l means…
    • Corresponds to the name of the subshell
      • L=0  subshell s
      • L=1  subshell p
      • L=2  subshell d
      • L=3  subshell f
    • Each of these 4 sublevels has a unique shape
    • Each orbital may contain at most, 2 electrons
      • S  sharp…spherical
      • P  principal…dumbbell shaped
      • D  diffuse…not all the same shape
      • F  fundamental…not all the same shape
  • 85.  
  • 86.  
  • 87.
    • When principle energy level n=1, then l=0, which means there is only a single sublevel (one orbital) which is the small, spherical 1s
    • When principle energy level n=2, then l can equal 0 or 1, which means that there are two sublevels (orbitals) 2s and 2p
      • 2s sublevel  bigger than 1s, still sphere
      • 2p sublevel  three dumbbell shaped p orbitals of equal energy called 2px, 2py, and 2pz
        • The letters are just there to tell you what axis the electrons go on: x,y, or z axis
    • When the principle energy level n=3, then l can equal 0,1, or 2, which means that there are 3 possible sublevels:
      • 3s, sphere, bigger than 1s and 2s
      • 3p, dumbbells
      • 3d
        • Each d sublevel consists 5 orbitals of equal energy
        • Four d orbitals have same shape but different orientations
        • Fifth d orbital, 3d z2 is shaped and oriented different from the other four
    • When the principle energy level n=4, then 1 can equal 0,1,2, or 3 which means l=n-1=4 possible sublevels:
      • Seven f orbitals of equal energy ( 2 electrons in each orbital)
      • 4s, sphere
      • 4p, dumbbells
      • 4d,
      • 4f
  • 88. n = # of sublevels per level n 2 = # of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f
  • 89.  
  • 90.  
  • 91. Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x
  • 92. Remember…
      • 1. Principal #  energy level
      • 2. Ang. Mom. #  sublevel (s,p,d,f)
    There are two more quantum numbers (3 and 4) we will discuss next class
  • 93. Third Quantum Number
    • M l  specifies the orientation of the orbital in space containing the electron
    • Tells us whether the orbital is on the x, y, or z axis
  • 94. Fourth Quantum Number
    • M s  related to the direction of the electron spin
    • Tells us if electron has a clockwise spin or counter clockwise spin
    • Specifies orientation of electrons spin axis
  • 95. Recap…
    • Bohr?
      • Orbits explained hydrogen’s quantized energy states
    • De Broglie?
      • Dual particle and wave nature of electrons
    • Schrodinger?
      • Wave equation predicted existence of atomic orbitals containing electrons
  • 96. Electron Configuration
    • Definition: arrangement of electrons in an atom
    • Basic rules for filling up orbital's with electrons
    • Which is more stable, low energy or high energy?
      • So which orbitals are going to be filled up first?
      • We are going to want an arrangement that gives us the lowest possible energy
  • 97. Ground state electron configuration
    • The most stable, lowest energy electron arrangement of an atom
    • Each element has a ground-state electron configuration
  • 98. Three Rules for Electron Arrangement
    • Aufbau Principle
    • Pauli Exclusion Principle
    • Hund’s Rule
  • 99. Aufbau Principle
    • Each electron occupies the lowest energy orbital available
    • In order to do this, you must learn the sequence of atomic orbitals from lowest to highest energy
    • Aufbau Diagram
      • Each box represents an orbital
      • Each arrow represents an electron
      • Only two arrows per box…
        • Only two electrons per orbital
  • 100.  
  • 101.  
  • 102. Some important things to remember about Aufbau…
    • All orbitals related to an energy sublevel are of equal energy
      • All three 2p orbitals have the same energy
    • In a multi-electron atom, the energy sublevels within a principle energy level have different energies
      • All three 2p orbitals are of higher energy than the one 2s orbital
  • 103.  
  • 104.
    • In order of increasing energies, the sequence of energy sublevels within a principle energy level is s, p, d, f
    • Orbitals related to energy sublevels within one principle energy level can overlap orbitals related to energy sublevels within another principle level
      • Ex. An orbital related to the atoms 4s sublevel has a lower energy than the five orbitals related to 3d sublevel.
  • 105.  
  • 106. Pauli Exclusion Principle
    • States that a maximum on 2 electrons can occupy a single atomic orbital but only if the electrons have opposite spins
    • Wolfgang Pauli
    • Austrian Physicist
    • Observed atoms in excited states
  • 107.
    • Each electron has a spin
    • Kinda like a spinning top
    • It can only spin in one of 2 directions
    • In order for electrons to be together in an orbital, they must have opposite spins
  • 108. Hund’s Rule
    • What kind of charge do electrons have?
    • Do they attract or repel each other?
    • So……..
    • Hund’s Rule states that single electrons with the same spin must occupy all each energy equal orbital before additional electrons with opposite spins can occupy the same orbital
  • 109. 2p orbitals
  • 110. Read section 5-3!