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

Line spectrum

Ne

I2

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