Properties of Light

1. Properties of Light

2. http://www-03.ibm.com/ibm/history/exhibits/vintage/vintage_4506VV1003.htmlhttp://www-03.ibm.com/ibm/history/exhibits/vintage/vintage_4506VV1003.html

3. The lowest energy state of an electron is the ground state. A state in which that atom has a higher energy potential is called an excited state. As the excited electron falls back to its ground state, it releases EM radiation of an energy that corresponds to the amount of energy gained to reach the excited state. Light

5. How Microwaves Work http://portal.acs.org/portal/acs/corg/content?_nfpb=true&_pageLabel=PP_SUPERARTICLE&node_id=569&use_sec=false&sec_url_var=region1&__uuid=350b3935-ebfe-486c-b3ce-faf11f2847bc

6. Dual Nature of Light Light can exhibit wave-like properties. (electromagnetic spectrum, diffraction and interference) Light can also behave like a particle (light travels in a straight line, when shined on a piece of metal it can cause electrons to be released: photoelectric effect)

7. EM wave- a form of energy that exhibits wavelike behavior as it travels through space All the forms of EM radiation form the electromagnetic spectrum Light is an electromagnetic wave

8. Speed- all forms of EM radiation travel at 3.00 x 108 m/s in a vacuum Wavelength-the distance between 2 consecutive waves Frequency- the # of waves that pass a stationary pt. in one second Amplitude- the height of a wave measured from the origin to its crest Properties of waves

9. Light exhibits many wavelike properties but can also be thought of as a stream of particles called photons. Photon- a particle of EM radiation having zero mass and carrying a quantum of energy The Dual Nature of Light

10. Photoelectric Effect • Refers to the emission of electrons from a metal when light shines on the metal http://phet.colorado.edu/en/simulation/photoelectric

11. Quantum- the minimum quantity of energy that can be lost or gained by an atom This relationship is expressed as: E = hv Where E = energy h= Planck’s constant v= frequency of light Planck’s Constant= 6.626 x 10-34 J s Particle Description of Light

12. Calculating Quantities Related to Light Energy of a photon of light E = hv Where E = energy (Joules) h= Planck’s constant v= frequency of light(Hz or 1/s) Planck’s Constant= 6.626 x 10-34 J s Speed of wave = wavelength x frequency c = λv This is an __________ relationship. Practice using the equations pg 126 #47 and #50

13. When the EM radiation released is passed through a prism or diffraction grating- it is separated into a series of specific frequencies of visible light. This is referred to as a line emission spectra http://jersey.uoregon.edu/vlab/elements/Elements.html you can see emission or absorption for any element

14. 1. Name the type of spectrum you see from the gas tubes. 2. Look through the spectroscope at sunlight. What type of spectrum do you see? 3. How is the sunlight spectrum different from the gas tube spectrum? 4. Why are the two spectra different?

15. energy of emitted photon = (atom energy before) - (atom energy after)

17. Bohr’s Model(particle idea of light) • The electron can circle the nucleus only in allowed paths or orbits • In an orbit, an electron has a fixed energy • The lowest energy state is closest to the nucleus • An electron can move to a higher orbital if it gains the amount of energy equal to the difference in energy between the initial orbit and the higher energy orbit

18. Suggested that electrons could also have a wave nature much like light. This was based on the fact that: Electrons can be diffracted (bending) http://www.youtube.com/watch?v=p6A0bzlm0YA http://www.youtube.com/watch?v=KT7xJ0tjB4A Electrons exhibit interference (overlapping that results in a reduction of energy) DeBroglie’s Idea(wave description of light)

19. Heisenberg Uncertainty Principle • It is impossible to determine simultaneously both the position and velocity of an electron.

20. Laid the foundation for modern quantum theory. Quantum Theory describes mathematically the wave properties of electrons. The solutions to this equation give only the probability of finding an electron at a given location. Schrodinger’s Wave Equation

21. Conclusion • Electrons do not travel in neat orbits • They exist in three-dimensional regions called orbitals that indicate the probable location of an electron. http://www.youtube.com/watch?v=DZGINaRUEkU (song about quantum theory)

22. Quantum Numbers The location of electrons within the atom can be described using quantum numbers: • Principle • Orbital (Angular Momentum) • Magnetic • Spin

23. Gives the principle energy level n= 1, 2, 3, etc. Maximum # of Electrons for Orbitals: 1st- 2 2nd – 8 3rd – 18 4th - 32 Principle Quantum Number

24. Orbital Quantum Number • Tells the shape or type of orbital http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/index.html http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao4.html

25. Designates specific regions of space w/in each energy sublevel (s, p, d, f) Ex. p sublevel has 3 orbitals (px, py, pz) Magnetic Quantum Number

26. Designates direction of electron spin Electrons w/in an orbital spin in opposite directions Spin Quantum Number

28. Governing Rules and Principles • Pauli Exclusion Principle- only 2 electrons in each orbital • Aufbau Principle- electrons must occupy lower energy orbitals first • Hunds Rule- a second electron can not be added to an orbital until each orbital in a sublevel contains an electron

29. Summary: