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37 Cards in this Set
- Front
- Back
List 5 examples of electromagnetic radiation |
Radiowaves, microwaves, infared, visible light, ultraviolet light |
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Photoelectric effect |
When Light or a certain frequency shines on a metal, the metal glows |
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Wavelength and frequency are _________ related. |
inversely |
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Energy and frequency are _________ related. |
Directly |
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Energy and Wavelength are ________ related. |
Inversely |
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Wave-particle nature |
It behaves like a wave but can be measured like a particle. |
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2 differences between red and green light on the electromagnetic spectrum |
Green has a greater frequency Red has a greater wavelength |
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Put in order of increasing wavelength: ultraviolet light, microwaves, radio waves, X-rays |
X-rays, ultraviolet light, microwaves, radiowaves |
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Electromagnetic wave relationship formula |
c = λv
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c = λv c= λ= v=
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c= constant speed of a wave =3.00x10^8 λ= wavelength in meters (m) v= Frequency in Hertz (Hz) or 1/s |
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Formula for energy of a photon |
E=hv |
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E=hv E= h= v= |
E= Energy in Joules (J) h= Planck's constant, =6.626x10^-34Js (Joule seconds) v= frequency in Hertz (Hz) or 1/s |
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Wavelength |
the distance between 2 corresponding points on a wave |
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Frequency (v) |
the occurrance of a wave |
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Electromagnetic radiation |
A form of energy that travels as a wave |
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Electromagnetic spectrum |
Includes all forms of EM radiation (radiowaves, microwaves, infared, visible light, ultraviolet, X-ray, gamma rays) |
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The speed of light |
3.0x10^8 m/s |
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Quantum |
a minimum amount of energy that can be lost or gained by an atom |
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Photons |
Massless particles that carry a quantum of energy |
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Planck's constant |
relates energy with frequency h=6.626x10^-34js |
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Electron configuration |
the use of orbitals to tell the exact location of every electron of an atom (assume ground state), arrangement of electrons |
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Aufbau Principle |
electrons fill an atom by lowest energy levels and orbital first |
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Diagonal rule |
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f |
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Hund's rule |
orbitals of equal energy fill with 1 electrons before 2 (doesn't apply to s) |
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electron capacity |
s=2e- p=6e- d=10e- f=14e- |
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Pauli Exclusion Principle |
only 2 electrons can occupy each orbital variation and must be opposite in their spin |
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highest occupied level |
electrons in the outermost ring (from config) |
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valence electrons |
electrons in the highest occupied level used for chemical bonding |
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Bohr |
-Presented the "Planetary Model" of the atom -Electrons stay in orbits as close as possible to the nucleus with minimum energy (ground state) -Electrons could jump orbits when energy is absorbed by the atom (excited state) -When electrons absorb this energy they quickly release it in the form of light (back to ground state)
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Planetary model |
electrons travel in specific orbits around the nucleus |
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Problems with Bohr's Planetary Model |
-Only explained hydrogen atom (with 1 electron) -Didn't account for other chemical behavior of atoms -Conclusion: there really are no circular orbits of electrons |
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Heisenberg |
Tried and used methods to detect electrons Idea: to detect electrons which behave like light using light (photons) Impossible to take a measurement of an object without disturbing the object (light interferes with electron movement)
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Heisenberg Uncertainty Principle |
It is impossible to determine the location and speed of an electron simultaneously -the more you know about one quantity, the less you can know about the other
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Schrodinger |
Derived a mathematical equation that treats electrons as waves; the "Wave Mechanical Model" -Plots 3-D regions around the nucleus where electrons are probably located at any given point in time -3-D region called an Orbital (a section of empty space)
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Orbital |
Section of empty space |
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Main Energy Level (Ring or shell) |
Theoretically levels go from 1-infinity Actuall levels 1-7
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Orbital shape (subshell) |
Sphere (s), propeller (p), double propeller (d), flower petals (f) |