The electronic structures of atoms
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| Electromagnetic radiation | The wavelength of EM radiation has the symbol .
Wavelength is the distance from the top (crest) of one wave to the top of the next wave.
Measured in units of distance such as m, cm, Å.
1 Å = 1 x 10-10 m = 1 x 10-8 cm
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| Frequency | symbol V
Frequency is the number of crests or troughs that pass a given point per second
Measured in units of 1/time – 1/s - s-1
lamda*v = speed of propagation of the wave or
lamda*v = c (for light)
c = speed of light (3.00 x 108m/s)
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| Wavelength and frequency | are inversely proportional to each other, for the same wave, shorter wavelengths correspond to higher frequency
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| UV vs IR | UV: shorter λ, higher ν
IR: longer λ, low ν
High Energy; High Frequency; Short λ or Low Energy; Low Frequency; Long λ
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| Electromagnetic Radiation | Molecules interact with electromagnetic radiation.
Molecules can absorb and emit light.
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| Once a molecule has absorbed light (energy), the molecule can | Rotate
Translate
Vibrate
Electronic transition
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| Planks constant | E=hV or E=(hc)/lamda
h=6.626e-34j*s
energy is directly proportional to frequency
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| The Photoelectric Effect | e- r ejected only if light is of sufficiently short wavelength/high energy no matter how long or bright
when photon energy of light is high enough to start photoelectric effect the # of e- emitted per second (current) increases with brightness/intensity
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| Intensity of light | The intensity of the light is the brightness of the light. In wave terms, it is related to the amplitude of the light waves. In photon terms, it is the # of photons hitting the target.
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| emission spectrum | formed by an electric current passing through a gas in a vacuum tube (at very low pressure) which causes the gas to emit light.
Sometimes called a bright line spectrum
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| absorption spectrum | is formed by shining a beam of white light through a sample of gas.
Absorption spectra indicate the wavelengths of light that have been absorbed.
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| Every element has a unique spectrum. | Can use spectra to identify elements.
Can serve as fingerprints that allow us to identify different elements present in a sample, even in trace amounts.
Atomic and molecular spectra are important indicators of the underlying structure of the species.
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| Rydberg equation | is an empirical equation that relates the wavelengths of the lines in the hydrogen spectrum. Derived from numerous observations, not from theory
LOOK AT SLIDE
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| When an electron is excited from a lower energy level to a higher one | it absorbs a definite (or quantized) amount of energy. When the electron falls back to the original energy level, it emits exactly the same amount of energy it absorbed in moving from the lower to the higher energy level
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| Orbits | Orbits get farther apart as n increases, but their energies get closer together.
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| Postulates of Bohr’s theory | Atom has a number of definite and discrete energy levels (orbits) in which an electron may exist without emitting or absorbing electromagnetic radiation.
As the orbital radius increases so does the energy
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| Postulates of Bohr's theory | The larger the value of n, the farther from the nucleus is the orbit being described and the radius of this orbit increases as the square of n increases.
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| Postulates of Bohr's theory | e- may move from one energy level/orbit to another, but, in so doing, monochromatic radiation is emitted or absorbed in accordance with the following equation.
Energy is absorbed when e- jump to higher orbits
Energy is emitted when e- fall to lower orbi
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| Postulates of Bohr's theory | e- moves in a circular orbit about nucleus & motion is governed by ordinary laws of mechanics & electrostatics, with the restriction that the angular momentum of the e- is quantized (can only have certain discrete values).angular momentum = mvr = nh/2pi
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| origin of emission spectra | Light of a characteristic wavelength (and frequency) is emitted when electrons move from higher E (orbit, n = 4) to lower E (orbit, n = 1).
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| origin of absorption spectra | Light of a characteristic wavelength (and frequency) is absorbed when electrons jump from lower E (orbit, n = 2) to higher E (orbit, n= 4)
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| Electron wavelength | =h/(mv)
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