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(Electron Configurations)
Scientists
began to
understand how electrons
acted by observing the way
that light interacts with
matter.
 In
your notebook, answer the question:
How would you describe what light is to
someone else?
 Electromagnetic
Radiation-form of energy
that exhibits wave-like behavior as it travels
through space.
 Electromagnetic Spectrum-ordered
arrangement by wavelength or frequency for
all forms of electromagnetic radiation.
 Wavelength-lambda
(λ)
The distance between corresponding points
on adjacent waves. Units: m, nm, cm, or Å
 Frequency-nu (ν)
The number of waves passing a given point in
a definite amount of time. Units: hertz (Hz)
or cycles/s = 1/sec = s-1
 When
an electric field changes, so does the
magnetic field. The changing magnetic field causes
the electric field to change. When one field
vibrates—so does the other.
 RESULT-An electromagnetic wave.
EM waves do not require a medium to propagate
through
 Today



you will need:
Notes
Calculator
Periodic Table
Agenda:




Review Properties of Waves and EM Spectrum
Flame Test Demo
Frequency and Wavelength Calculations
Goal: Know the relationship between λ and v and
how to calculate them!
 Wavelength-lambda
(λ)
The distance between corresponding points
on adjacent waves. Units: m, nm, cm, or Å
 Frequency-nu (ν)
The number of waves passing a given point in
a definite amount of time. Units: hertz (Hz)
or cycles/s = 1/s = s-1
 Electromagnetic
Radiation-form of energy
that exhibits wave-like behavior as it travels
through space.
 Electromagnetic Spectrum-ordered
arrangement by wavelength or frequency for
all forms of electromagnetic radiation.
c
= λ∙ν
 λ = wavelength (m)
 ν = frequency (Hz)
c
= speed of light= 3.0 x 108 m/s (constant)
λ
and ν are _______________ related.
 Calculate
the frequency for violet light.
Truck-mounted helium-neon laser produces
red light whose wavelength (λ ) is 633
nanometers. Determine the frequency
(v).
*Remember that c=3.0x108m/s.
*Use the formula c= λ . v
c =3.0x108 m/s
c= λ . v
v=c / λ
λ = 633nm= 6.33x10-7m
v = 3.0x108 m/s = 0.47x 1015s-1 = 4.7x1014 s-1
6.33x10-7m
Frequency = 4.7x1014 Hz (cycles per second)
 EX:
Find the frequency of a photon with a
wavelength of 434 nm.
GIVEN:
WORK:
=c
=?

 = 434 nm
= 4.34  10-7 m  = 3.00  108 m/s
-7 m
8
4.34

10
c = 3.00  10 m/s
 = 6.91  1014 Hz
 Calculate
the frequency for the yelloworange light of sodium.
 Calculate
the frequency for violet light.
E
= h∙ν
 E = energy (joule)
 h = Planck’s constant = 6.63 x 10-34 J∙s
 ν = frequency (Hz)
 E and ν are ______________ related.
 Calculate
the energy for the yellow-orange
light for sodium.
 Calculate
the energy for the violet light.
 Welcome!
Please turn into the box:
-Chp. 4 Outline
-EM Spectrum Practice (if you didn’t Fri)
-Energy Calculations (if you didn’t Fri)
 Chp.
4 Reading Quiz
 Pass Back Chp. 3 Quiz
 Wavelength, Frequency, Energy Problem
 Evidence that light is a wave!
 If Time: Evidence that light is a particle!
 Pick
your favorite (or any, really) radio
station. Knowing that radio station
frequencies are measured in Mega Herz and
that 1 MHz = 106Hz, calculate:


A)The wavelength (in m) of your radio wave
B) The Energy associated with your station
Discuss: What do you know/notice about the
relationship between wavelength and energy?
 Today!
Understand that there is experimental
evidence that supports the idea that light is
both a wave AND a particle.
 Constructive
Interference:
Before
During
 Destructive
Interference:
Before
During
2
problems that could not be explained if
light only acted as a wave.
 1.)
Emission of Light by Hot bodies:
Characteristic color given off as bodies
are heated: red  yellow  white
If light were a wave, energy would be given
off continually in the infrared (IR) region of
the spectrum.
 2.)
Absorption of Light by Matter =
Photoelectric Effect
Light can only cause electrons to be ejected
from a metallic surface if that light is at
least a minimum threshold frequency . The
intensity is not important.
If light were only a wave intensity would be
the determining factor, not the frequency!
 Please
turn your photoelectric effect
assignment into the box.
 Tomorrow


is the Flame Test Lab.
We will meet in Rm. 126
To participate, you MUST be wearing proper lab
attire.
 When
an object loses
energy, it doesn’t happen
continuously but in small
packages called “quanta”.
“Quantum”-a definite
amount of energy either
lost or gained by an atom.
“Photon”-a quantum of
light or a particle of
radiation.
 Excited
State: Higher energy state than the atom
normally exists in.
 Ground State: Lowest energy state “happy state”
 Line Spectrum: Discrete wavelengths of light
emitted.
 2 Types:
1.) Emission Spectrum: All wavelengths of light
emitted by an atom.
 2.) Absorption Spectrum: All wavelengths of light
that are not absorbed by an atom. This is a continuous
spectrum with wavelengths removed that are absorbed
by the atom. These are shown as black lines for
absorbed light.
 Continuous Spectrum: All wavelengths of a region
of the spectrum are represented (i.e. visible light)

Hydrogen’s spectrum can be explained with the
wave-particle theory of light.
 Niel’s Bohr (1913)

1.) The electron travels in orbits (energy levels)
around the nucleus.
 2.) The orbits closest to the nucleus are lowest in
energy, those further out are higher in energy.
 3.) When energy is absorbed by the atom, the
electron moves into a higher energy orbit. This
energy is released when the electron falls back to a
lower energy orbit. A photon of light is emitted.

1.
2.
3.
Bohr models are used to predict reactivity
in elements.
Reactivity refers to how likely an element
is to form a compound with another
element.
When looking at Bohr models, we look at
its valence electrons (the electrons on the
last energy level) to determine reactivity.
1.
2.
3.
4.
5.
Draw the nucleus.
Write the number of neutrons and the
number of protons in the nucleus.
Draw the first energy level.
Draw the electrons in the energy levels
according to the rules (on the next
slide). Make sure you draw the
electrons in pairs.
Keep track of how many electrons are
put in each level and the number of
electrons left to use.
Each electron shell can
hold a certain number of
electrons
 Electron shells are filled
from the inside out
 Noble Gases have full
outer electron shells
 All other elements have
partially filled outer
electron shells

Electron
Shell
1
Number of
Electrons
2
2
8
3
8
4
18
5
18
6
32
7
32
In order to draw Bohr models of these elements, you
must first determine the number of protons, neutrons,
and electrons. Once you have found this information,
follow the directions to draw your model.
6
C
Carbon
12.011
6
6
6
Protons: _____
Neutrons: _____
Electrons: ______
2
How many energy shells will this have? ____
4
How many valence (outer) electrons does this element have? ____
Bohr Model:
16
S
Sulfur
32.066
16
16
16
Protons: _____
Neutrons: _____
Electrons: ______
3
How many energy shells will this have? ____
6
How many valence (outer) electrons does this element have? ____
Bohr Model:
3
Li
Lithium
6.941
3
4
3
Protons: _____
Neutrons: _____
Electrons: ______
2
How many energy shells will this have? ____
1
How many valence (outer) electrons does this element have? ____
Bohr Model:
10
Ne
Neon
20.180
10
10
10
Protons: _____ Neutrons: _____ Electrons: ______
2
How many energy shells will this have? ____
8
How many valence (outer) electrons does this element have? ____
Bohr Model:
15
P
Phosphorus
30.974
15
Protons: _____
16
15
Neutrons: _____ Electrons: ______
3
How many energy shells will this have? ____
5
How many valence (outer) electrons does this element have? ____
Draw the Bohr Model:
11
Na
Sodium
22.990
11
Protons: _____
12
11
Neutrons: _____
Electrons: ______
3
How many energy shells will this have? ____
1
How many valence (outer) electrons does this element have? ____
Draw the Bohr Model:
 Lyman
Series-electrons
falling to the 1st orbit,
these are highest energy,
_____ region.
 Balmer Series- electrons
falling to the 2nd orbit,
intermediate energy,
_______ region.
 Paschen Series-electrons
falling to the 3rd orbit,
smallest energy, ______
region.
 En
= (-RH) 1/n2
 En
= energy of an electron in an allowed orbit
(n=1, n=2, n=3, etc.)
 n = principal quantum number (1-7)
 RH = Rydberg constant (2.18 x 10-18 J)
 When an electron jumps between energy
levels: ΔE =Ef – Ei
 By


substitution: ΔE = hν = RH(1/ni2 - 1/nf2)
When nf > ni then ΔE = (+)
When nf < ni then ΔE = (-)
 DeBroglie
(1924)-Wave properties of the
electron was observed from the diffraction
pattern created by a stream of electrons.
 Schrodinger (1926)-Developed an equation
that correctly accounts for the wave
property of the electron and all spectra of
atoms. (very complex)
 DO
NOT WRITE THIS IN YOUR NOTES!
 Rather
than orbits  we refer to orbitals.
These are 3-dimensional regions of space
where there is a high probability of locating
the electron.
 Heisenberg Uncertainty Principle-it is not
possible to know the exact location and
momentum (speed) of an electron at the
same time.
 Quantum Numbers-4 numbers that are used
to identify the highest probability location
for the electron.
 1.)


Principal Quantum Number (n)
States the main energy level of the electron and
also identifies the number of sublevels that are
possible.
n=1, n=2, n=3, etc. to n=7
2.) Azimuthal Quantum Number (l)
 Values from 0 to n-1
 Identifies the shape of the orbital





l=0
l=1
l=2
l=3
s
p
d
f
sphere
dumbbell
1 orbital
3 orbitals
4-4 leaf clovers & 1-dumbbell w/doughnut5 orbitals
very complex
7 orbitals
 3.)


Magnetic Quantum Number (ml)
Values from –l  l
States the orientation in space (x, y, z)





ml = 0
ml = -1, 0, +1
ml = -2,-1,0,+1,+2
ml = -3,-2,-1,0,+1+2,+3
s
p
d
f
only 1 orientation
3 orientations
5 orientations
7 orientations
4.) Spin Quantum Number (ms)
Values of +1/2 to -1/2
 States the spin of the electron.
 Each orbital can hold at most 2 electrons with
opposite spin.
