Download teacher version filled in

EMR info
Waves, light, and energy: Where
chemistry and physics collide
http://imagers.gsfc.nasa.gov/ems/waves3.html
Before we get started….
1. What is light?


2.
List as many interactions of
light and matter as you can.

3.
Is it matter?
What forms of light exist?
think how light changes matter,
and how matter changes light
What are some uses of light?
First things first: Waves
a and b represent wavelength (λ)- the
distance of a wave from crest to
successive crest; measured in meters
http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec.html
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 (ν [nu]; I know some of you have
used f, move on and get with chemistry!):

the number of cycles 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

Violet
Indigo
Blue
Green
Yellow
Orange
Red






wavelength(nm)
f(*1014 Hz)
Energy (*10-19 J)
400---460
7.5--6.5
5.0--4.3
460---475
6.5--6.3
4.3--4.2
475---490
6.3--6.1
4.2--4.1
490---565
6.1--5.3
4.1--3.5
565---575
5.3--5.2
3.5--3.45
575---600
5.2--5.0
3.45--3.3
600---800
5.0--3.7
3.3--2.5
Some equations you need to
know




= c / ν
and
E = hν
So….
E = hc / 
And…
 = h / mv*
When
• = wavelength in m
• c = speed of light, 3.00E8 m/s
• ν (nu)= frequency in Hz
•(cycles/sec or s-1 or 1/s)
•E= energy in J
•h= Plank’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 = hν

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 eis 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.ht
m
http://www.xmission.com/~locutus/applet
s/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/chemistr
y/essentialchemistry/flash/linesp16.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
a Nitrogen
(z=7)
atom
Bohr Model and moving
electrons

http://www.colorad
o.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.colorad
o.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)



Where n=1, 2, 3, 4….
And RH = 2.18E-18J,
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=hν, so ΔE=hν
if ΔE=[(-2.18E-18J)(Z/n2)]f- [(-2.18E-18J)(Z/n2)]i
E changes continued


*Remember E=hν, so ΔE=hν 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 a excited state
If ΔE is negative



since Ef < Ei
E is released
The e- was going from excited state to a ground state
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

therefore  = h / p
p = mv



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
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 rampwe 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
..
Schrodinger’s Wave Equation


Continued wave-particle theory
Treated Hydrogen’s electron as a wave


Also worked for other elements’ electrons
(not just hydrogen like Bohr’s model).
Known as the quantum mechanical model


Limits electrons to certain energy levels (values)
Does not try to describe an electron’s path around
the nucleus
..
Schrodinger’s Wave Equation

Very complex mathematical equation


Solutions to equations are know as wave
functions
Is the probability of finding an electron in a
certain space around the nucleus



Predicts a 3 dimensional region known as an
atomic orbital
Looks like a fuzzy cloud
Density of cloud is determined by probability of
finding the e- there

high probability = dense cloud (dark)
Sublevels and Orbitals

The electrons are spread out in orbitals that
have varying



Shapes
Energy (distance from nucleus)
The orbitals are described in regards to their
quantum numbers


Descriptions that are descriptive and hierarchical
There are 4 numbers that describe an orbital

Written as follows: (#, #, #, ±#)
Principal quantum number (n)
The first number (1, #, #,±#)
Describe the


distance from the nucleus of the orbital
The energy of the orbital
Values

As
for n are integers
The smallest possible value is 1
the distance from the nucleus (and
therefore energy) increases, the number
increases
Quantum numbers
There periodic table and n


The 7 periods on the periodic table
correspond to n values
Each period has a unique n value



For the 1st period, n=1
For the 2nd period, n=2
And so on….
Angular Momentum (l)
(this is a script l, as in llama)



Is the shape of the sublevel
It is the second number in the description
(#,1, #, ±#)
Range from _____________(although we
never deal with anything above l=3)




s =0
p =1
d =2
f=3
The s sublevel l = 0
http://www.sfu.ca/~nbranda/28xweb/images/s_orbital.gif
p sublevel l = 1
d sublevel
l=2
Another look at d sublevel
f sublevel
l=3
General tutorials for electron
configuration stuff



some slides in this PowerPoint are from this
site already
http://www.wwnorton.com/chemistry/overvie
w/ch3.htm
See key equations and concepts (select
from menu on the left), as well as the
looking through the overview where to the
tutorials are listed (links for just those are on
the left, too)
Magnetic number (ml)


Denote the orbital sublevel that is filled
It is the third number in the description
(#,#,1, ±#)




s sublevesl have one orbital; a sphere has one
orientation in space
p sublevels have three orbitals; 3 orientations in
space
d sublevels have five orbitals; 5 orientations in
space
f sublevels have seven orbitals; 7 orientations in
space
“Flavors” of orbitals (ml)

s sublevels
have one
orbital; a
sphere has
one orientation
in space
“Flavors” of orbitals (ml)

p sublevels have three orbitals; 3
orientations in space
“Flavors” of orbitals (ml)

d sublevels
have five
orbitals; 5
orientations in
space
“Flavors” of orbitals (ml)

f sublevels
have seven
orbitals; 7
orientations
in space
Spin

Spin is +½ or -½

Up or down (could say clock wise and
counter-clock wise)
Summary, excluding spin
How we use this….

There is a specific order to how the efill the orbitals; it is not random

Although there are exceptions to the
rules (last thing we do)
“The aufbau diagram shows the energy of each
sublevel relative to the energy of other
sublevels. Each box on the diagram represents
an atomic orbital.” Excerpt From: Thandi Buthelezi, Laurel
Dingrando, Nicholas Hainen & Cheryl Wistrom. “Chemistry.” McGraw-Hill
Education, 2013. iBooks. https://itun.es/us/I0WrD.n
The principles of econfiguration

The Aufbau (next) Principle:


The Pauli Exclusion Principle:


That e- fill the lowest energy sublevel before
going to the next sublevel
That e-s are paired according to opposite spins
Hund’s Rule:

e-s spread out in equal energy orbitals before
pairing electrons

The first level to fill is the 1s level


It is the lowest energy sublevel
It holds two electrons


They are oppositely paired (up and down- ↑↓)
Each sublevel (each __) holds 2 electrons
Next…



The second sublevel is the 2s sublevel
It also holds 2 electrons (because s
holds 2, not because of the number),
also oppositely paired ↑↓
1s2, 2s2,then comes 2p6

So, as it states
above

1s fills, 2s fills ,then
comes 2p


It holds up to six
electrons
Because p orbitals
hold 6 electrons
Next…

From 2p,







3s fills with 2e-, then onto
3p, with 6e- then
4s with 2e- followed by
3d with 10e- (because d
holds 10e-)
Then 4p with 6e-
Notice, you follow the
arrows
Remember, the number
of electrons comes from
the letter (the orbital’s
momentum, m)



The sublevels of the
orbitals are first filled,
then you continue onto
the next level (Aufbau)
Also be sure to place
one electron in each
sublevel prior to filling
the level (↑ ↑ ↑ and
not ↑↓ ↑ _) (Hund)
e-s must be paired with
e-s of opposite spin
(↑↓, not ↑↑ or ↓↓)
(Pauli)
Putting it all together…


Carbon (neutral, so 6 electrons)
What this would look like:
↑↓ ↑↓ ↑ ↑ _
1s 2s 2p
(notice there are 6 arrows for 6 electrons)


This can also be written as 1s2 2s2 2p2
Notice the superscripts add up to 6
Nobel gases


Last column of P.T.
All end the same way with s2 p6



Neon (8 electrons) would look like:
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s
2p
Argon (18 electrons) would look like:
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s
2p
3s
3p
Can use this trend as a shortcut for
writing out long e- configurations
Nobel gas shortcut

Non-Nobel gas elements have inner econfigurations identical to certain Nobel gases



Write Nobel gas in brackets


Ne is 1s2 2s2 2p6
Na is 1s2 2s2 2p6 3s1
[Ne] instead of 1s2 2s2 2p6
After the brackets, write the electrons left over

[Ne] 3s1
* When using the shortcut, use the Nobel gas from the period (row)
before the element in question*
There are some exceptions…

This is because some energy levels are very close
together



electrons are able to move between close orbitals in order
to minimize repulsion
Example: the 4s and 3d orbitals are very close in
energy
So exceptions for some period 4 d block elements
occur



Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4
Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Because it takes less energy to split the electrons between
the 5 sublevels than it does to put them together in the 4s
and 3d
Valence electrons

The electrons in the outermost orbital (usually
highest energy level)




Be is 1s2 2s1 has 1 valence electrons (2s1)
N is 1s2 2s2 2p3 has 5 valence electrons (2s2 + 2p3 )
Determine the chemical properties of elements
Used in forming chemical bonds
Valence electrons

Can represent valence electrons visually using
electron dot structure


Dots stand for valence electrons
Placed one at a time on the four sides of the symbol
(they may be placed in any sequence) and then
paired up until all are shown