Energy Levels & Photons Atomic & Nuclear Lesson 2.
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Transcript of Energy Levels & Photons Atomic & Nuclear Lesson 2.
![Page 1: Energy Levels & Photons Atomic & Nuclear Lesson 2.](https://reader036.fdocuments.net/reader036/viewer/2022082518/56649eef5503460f94bfeeba/html5/thumbnails/1.jpg)
Energy Levels & Photons
Atomic & Nuclear Lesson 2
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Homework
• Revise for the skills test.
• After the test, complete worksheet & read about emission spectra and absorption spectra in the textbook.
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Learning Objectives
• To outline evidence for the existence of atomic energy levels.
• To know what a photon is.
• To know how to calculate the energy of a photon.
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Part of E-M Spectrum
Wavelength Range
Visible 400-700 nm
Electromagnetic Spectrum
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Part of E-M Spectrum
Wavelength Range
Radio > 0.1mMicrowave 0.1m to 1mmInfrared 1 mm to 700
nmVisible 400-700 nmUltraviolet 400 nm – 1 nmX-rays < 1 nmGamma-rays < 1 nm
Electromagnetic Spectrum
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Mnemonic?
• Rabbits Radio• Mate Microwaves• In Infrared• Very Visible• Unusual Ultraviolet• X-rated X-rays• Gardens Gamma Rays
• Highest energy? Highest frequency?
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Looking at Emission Spectra
• Look at the spectra of a white light and a set of standard discharge lamps: sodium, neon, hydrogen and helium.
• Make notes of your observations. What are the differences between the spectra?
(Diagram: resourcefulphysics.org)
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Emission spectra
• An energy input raises the electrons to higher energy levels. This energy input can be by either electrical, heat, radiation or particle collision.
• When the electrons fall back to a lower level there is an energy output. This occurs by the emission of a quantum of radiation.
• When ever possible, electrons occupy the lowest energy level called the ground state.
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Emission Spectra
Each element has its own specific set of lines.
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How do we get emission spectra?
Bohr proposed that electrons moving between energy levels caused the line spectra.
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The energy levels and spectra series
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Photons
Light is not a continuous wave but is emitted as “packets”.
These “packets” of energy are called “photons”.
The different colours of light correspond to different photon energies only certain energies are allowed.
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Absorption spectra• When light of all frequencies is passed
through a gas then the gas absorbs light of the same frequency as it would emit.
• The light is radiated in all directions causing a reduction of intensity in the direction of the observer (dark lines).
• And so is seen when emitted energy is absorbed by a medium the Sun
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Wave Equation (GCSE)
• Recall from GCSE:-
• Electromagnetic radiation travels at the speed of light which is 3.00 × 108 m s-1 in a vacuum.
(m)Wavelength Hz)Frequency( )s (m Speed Wave -1
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Wave Equation (AS Physics)
• In AS we write this equation using symbols:-
• where c = the speed of light in a vacuum (m s-1)
• f = the frequency of e-m radiation (Hz)• λ = the wavelength of e-m radiation (m)
f c
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Photon Energy• We can work out the energy of an
incoming photon using this equation:-
• Where E = Energy of Photon in Joules (J)• f = Frequency of the Radiation in Hertz
(Hz)• h = Planck’s constant = 6.63 x 10-34 JsOr in words:-
Radiation ofFrequency ConstantsPlanck'EnergyPhoton
hfE
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Photon Energy• Recall from GCSE that f = c/λ so we
can substitute this into the photon energy equation E=hf to get:
• Or in words:-
(m)RadiationofWavelength
)(msLight of Speed(Js)ConstantsPlanck'(J)EnergyPhoton
1
hc
E
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Worked Example
• Q: What is the photon energy for UV radiation with a wavelength 400 nm?
λ = 400 nm = 400 × 10-9 mE = ?h = 6.63 × 10-34 J sc = 3 × 108 ms-1
J19-9-
-18-34
10 97.4m 10 400
ms 10 x 3 Js 10 6.63
hcE