Transcript Electromagnetic Radiation

(Electron Configurations)
 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/sec = 1/sec = sec-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.
Waves or Particles
 Electromagnetic radiation has properties of
waves but also can be thought of as a stream of
particles.
 Example: Light
 Light as a wave: Light behaves as a transverse
wave which we can filter using polarized lenses.

Light as particles (photons)

When directed at a substance light can knock
electrons off of a substance (Photoelectric
effect)
c
= λ∙ν
λ = wavelength (m)
ν = frequency (Hz)
c
= speed of light= 3.0 x 108 m/sec (constant)
λ
and ν are _______________ related.
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 v= c
λ
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
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!
 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.
 Calculate
the frequency for the yelloworange light of sodium.
 Calculate
the frequency for violet light.
 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∙sec
 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.
 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.

 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)
 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.) Orbital Quantum Number
 Identifies the shape of the orbital





s (2 electrons)
P (6 electrons)
d (10 electrons)
f (14 electrons)
sphere
dumbbell
1 orbital
3 orbitals
4-4 leaf clovers & 1-dumbbell w/doughnut5 orbitals
very complex
7 orbitals
 3.)

Magnetic Quantum Number
Identifies the orientation in space (x, y, z)

s  1 orientation
p 3 orientations
d 5 orientations
f 7 orientations

4.) Spin Quantum Number





States the spin of the electron.
Each orbital can hold at most 2 electrons with
opposite spin.
 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.
