EMR info
http://imagers.gsfc.nasa.gov/ems/waves3.html
Waves, light, and energy: Where chemistry and physics collide
Before we get started….
1. What is light? 1. Is it matter?2. What forms of light exist?
2. List as many interactions of light and matter as you can. think how light changes matter,
and how matter changes light
3. What are some uses of light?
http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec.html
First things first: Waves
a and b represent different wavelengths (λ)- the distance of a wave from crest to successive crest; measured in meters
Waves: amplitude
The height of a wave from crest to midline or trough to midline; measured in meters
Terms you need to know:
Wavelength (λ) Amplitude Frequency (ⱱ) ; I know some of you have
used f, move on and get with chemistry! :) the number of cycles (oscillations) per second
measured in cycles per second (s-1) or Hz (Hertz)
Waves on a string
http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec2.html
http://micro.magnet.fsu.edu/primer/lightandcolor/images/electromagneticfigure1.jpg
http://lepus.physics.ualr.edu/~tahall/EXAM2/emspec.jpg
http://www.arpansa.gov.au/images/emsline2.gif
Visible Light
color wavelength(nm) ⱱ (*1014 Hz) Energy (*10-19 J)
Violet 400---460 7.5--6.5 5.0--4.3 Indigo 460---475 6.5--6.3 4.3--4.2 Blue 475---490 6.3--6.1 4.2--4.1 Green 490---565 6.1--5.3 4.1--3.5 Yellow 565---575 5.3--5.2 3.5--3.45 Orange 575---600 5.2--5.0 3.45--3.3 Red 600---800 5.0--3.7 3.3--2.5
Some equations you need to know
λ= c / ⱱ
and E = hⱱ So….
E = hc / λ And…
λ = h / mv*
When• λ = wavelength in m• c = speed of light, 3.00E8 m/s• ⱱ (nu)= frequency in Hz
• (cycles/sec or s-1 or 1/s)• E= energy in J• h= Planck’s constant, 6.626E-34
J*s [Joule(seconds)]• m= mass of particle in kg• V*= velocity in m/s
What the h? Planck’s Constant
When metals are heated, they glow 1800s- physicists were trying to determine
the relationship between the color (wavelength) and intensity of the glow
Max Planck (1900)- energy can be released or absorbed only in little chunks (packets) of energy “of some minimal size”
Max Planck and the h
The chunks of energy were dubbed “quantum” (“fixed amount”), which is the smallest amount that can be emitted or absorbed as EMR.
Proposed: E = hⱱ The energy (E) of a single quantum is
equal to its frequency (ν) times a constant
Planck and the Nobel (Physics)
Planck determined that h= 6.626E-34 J-s
Energy is always released in multiples of hv (1hv, 2hv, 3hv, etc)
h is so small that we cannot see the effects of this in our daily lives
Analogous to… Planck won the 1918 Nobel Prize in
physics for his work
Einstein & Bohr: Perfect Together
Einstein, left
Bohr, above
Einstein:The Photoelectric Effect
Einstein discovered that one could cause electrons to be ejected from the surface of a metal if the energy of the light wave was strong enough
He treated the light needed to do this as a piece of matter- a photon, if you will
This ejection of e- is the photoelectric effect
The Photoelectric Effect
Only light of a certain energy could knock off an electron from the metal Intense light of a weaker wavelength
would not work, but even a low intensity of the correct wavelength would work
(the energy of the light is transferred to the kinetic energy of the electron)
Hmmm… light acting as a particle and as a wave…..
The photoelectric effect…
Online animations PhET http://www.lewport.wnyric.org/mgagnon/P
hotoelectric_Effect/photoelectriceffect1.htm
http://www.xmission.com/~locutus/applets/Photoelectric.html
Getting to Bohr….
Light of a given wavelength is monochromatic (one color)
Most common EMR sources are polychromatic, but we see only one color
These can be reduced to a spectrum when the different wavelengths are separated out
Spectral Emissions
Continuous spectrum: shows all colors of the rainbow
Bright line spectrum: only certain wavelengths are visible (the rest do not appear at all)
Different elements have different bright line spectrum when they are heated Na is yellow Ne is orange-red
http://www.cartage.org.lb/en/themes/Sciences/Astronomy/Modenastronomy/Interactionoflight/AtomicAbsorption/AtomicAbsorption.htm
Hydrogen Spectra
Emission Spectra
Absorption Spectra
http://www.mhhe.com/physsci/astronomy/applets/Bohr/content_files/section1.html
http://www.cartage.org.lb/en/themes/Sciences/Astronomy/Modenastronomy/Interactionoflight/AtomicAbsorption/AtomicAbsorption.htm
Color and what you see:
Absorption: the wavelengths that are absorbed by an object are not available for us to see, as we see the wavelengths of light that are reflected off of an object
This is not the same as those wavelengths that are emitted by an object that is emitting radiant energy.
Color and what you see…
Chlorophyll absorption spectra
Perception of color
Line spectra formation- go to…..
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf
http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/pem1s3_1.swf
Bohr Model and Spectral Emissions
Bohr proposed that the emission of light energy from an (electrically or thermally) excited atom corresponds to the orbit of the electron around the nucleus of the atom That energy can only be achieved by
being a specific distance from the nucleus
What you’ve seen so far….
Model of an Iodine atom (atomic number =53)
Bohr Model and moving electrons
http://www.colorado.edu/physics/2000/quantumzone/bohr.html
Energy levels- Bohr Model Electrons travel within set
energy levels that have a particular energy associated with each level
After all, the e-s are moving around the nucleus think KE here
Each shell has a number Closest to the nucleus is n=1 For each successive level add
1 to n n=2, n=3, ect….
Energy increases as the distance from the nucleus increases
Bohr Model and moving electrons
http://www.colorado.edu/physics/2000/quantumzone/bohr.html
Electron config in energy level
SO…
We know that the e-’s are free to move around the nucleus
They also can move from one energy level to the next (and fall) back when energy is added Move from ground state (“home” level) to
a higher level (the “excited” state) Returning back to the ground state
releases energy
This emission is how we see colors: the wavelengths of EMR released from
an atom when it has been excited by Heat energy Electrical energy Chemical energy
Think glowing red hot metal, or fireworks
Determining Energy for n
To determine the energy for a given energy level, use the equation:
En=(-RH)(Z/n2) RH = 2.18E-18J, Z= the atomic number of the atom n=1, 2, 3, 4…. So En=(-2.18E-18J)(Z/n2)
To determine E emitted or absorbed:
To determine the change in energy for a given energy transition:
ΔE=Ef-Ei *Remember E=hⱱ, so ΔE=hⱱ
so ΔE=[(-2.18E-18J)(Z/n2)]f- [(-2.18E-18J)(Z/n2)]i Remember that + values mean E that
is absorbed, and – values mean released
E changes continued *Remember E=hν, so ΔE=hⱱ to get the
frequency of the light emitted or absorbed If ΔE is positive
since Ef >Ei E is absorbed The e- was going from ground state to an
excited state If ΔE is negative
since Ef < Ei E is released The e- was going to ground state from an
excited state
To determine E emitted or absorbed:
What is the change in energy associated with an electron dropping from n=5 to n=1 in a Hydrogen atom?
ΔE=Ef-Ei so ΔE=[(-2.18E-18J)(Z/n2)]f- [(-2.18E-18J)(Z/n2)]i ΔE=[(-2.18E-18J)(1/12)]f- [(-2.18E-18J)(1/52)]I
ΔE = -2.09E-18 J Which means 2.09E-18J are released
Makes sense; an e- is dropping from 5 to1, E is released when e- drop
Back to basics EMR calcs…
That released Energy can be used to determine the wavelength and frequency of the EMR emitted. Remember that you need to treat the
energy as positive to do this! The sign only gives direction of energy flow There is no negative energy, only energy
leaving If you used – energy, you’d get a - or -ⱱ
This isn’t possible!
Also…life after Einstein and Bohr
We know that electrons have characteristics of both light (waves) and matter, so we say that they have a dual nature
De Broglie
De Broglie proposed that an electron moving about the nucleus had a wave-like behavior, so it has a particular wavelength associated with it. This wavelength depends upon the mass and velocity of the electron. = h / mv mv = the momentum of the particle
Mass* velocity = p momentum = p so p = mv
therefore = h / p
This matter-wave idea applies to all matter, not just to electrons
However, the mass is so large, and the wavelength so small, that we cannot see it in macroscale objects
This matter-wave theory led to applications like the electron microscope
Scanning electron microscope image of a leaf from a Black Walnut tree. Image shows a cross-section of a cut leaf, itsupper epidermal layer, mesophyll layer with palisade cells and vascular bundles, and lower epidermal layer. The protrusion at center is just over 50 microns tall. (Dartmouth Electron Microscope Facility/Dartmouth College)#
Pollen from a variety of common plants: sunflower, morning glory, hollyhock, lily, primrose and caster bean. The largest one at center is nearly 100 microns wide. (Dartmouth Electron Microscope Facility/Dartmouth College) #
De Broglie wavelength
Heisenberg:The Uncertainty Principle
We can’t determine information about small scale objects the same way we can for large scale objects Case in point: a ball rolling down a ramp-
we can get position, direction, and speed at the same time
We can’t for electrons Hence, the uncertainty principle
Heisenberg, cont’d
It is inherently impossible for us to simultaneously know both the exact momentum and exact location of an electron
This is because anything we do to determine the location or momentum of the electron moves it from its original path and location; this can’t be reduced past a certain minimal level
We can know only momentum or location- not both We can talk probability of the location/ momentum
of an electron
Which brings us to this question:
What the heck does all of this have to do with electron configuration and how matter behaves? On to electron configuration, courtesy of
Schrödinger and company (enter math that we’ll skip) Quantum theory
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