Light
Equations
Regular
Chemistry
Quantum Mechanical Model of Atom
 By 1900, light considered wave – like.
 Light consists of Electromagnetic Radiation
 All electromagnetic radiation travels at “c”
SPEED OF LIGHT 2.998 x 108 m/s
 A range of wavelengths from Radio Waves to
Gamma Rays .
Electromagnetic Radiation
c=
(lambda) = symbol of wavelength in meters
(nu) = symbol for frequency in Hz (Hertz) or s-1
The amount of wave cycles to pass a given point
per unit of time. Frequency & Wavelength
inversely related.
c=
As “ ” increases “ ” decreases, or vice versa
because “c” is constant.
Example: Light with a wavelength of 525 nm is
green. Find the frequency
First : Convert 525nm to m
525 nm | 10-9 m = 525 x 10-9 m
| 1 nm
Second: Find the frequency.
Since c = λ
then
=c
λ
2.998 x 108 m/s = = 5.71 x 1014 Hz
525 x 10-9 m
Wavelength (λ) is the distance between two
successive wave crests.
Amplitude (φ) is the height of a wave. Other
terms are intensity, brightness, etc..
Electromagnetic Radiation
 Sunlight consists of light with a continuous
range of wavelengths and frequencies.
 Each color of light has a range of
wavelengths or frequencies.

e.g. Red ~ 700 nm
Blue ~ 380 nm
Atomic Emission Spectra
 Every element emits light when it is excited
by electricity through its vapor (e.g. J.J.
Thomson and CRT experiment where he
discovered electrons).
 Each element has a characteristic Atomic
Emission Spectrum.
 Each spectral line corresponds to a set “ ” or
“ ” which is really directly related to a certain
“quantity of energy”.
Atomic Emission Spectra
 Unique spectra for each element allows you
to identify unknown compounds.
 Flame Test
 Classic Physics cannot explain the emission
spectra of atoms which consists of specific
lines.
Energy Quanta
 The amount of energy emitted by a body is
proportional to the frequency “ ” of the radiation.

E=h
Planck’s constant
h = 6.6262 x 10-34 J.s
The Photoelectric Effect
Photoelectric Effect
 Classical physics said any frequency of light
should generate electrons, but …
 Lower frequency red light will not cause
ejection of electrons from alkali metals.
 Higher frequency blue light will generate
photoelectric electrons.
Photoelectric Effect
 Einstein - There is a threshold value of
energy below which the photoelectric effect
does not occur.
 If frequency of photon too low, then no
photoelectrons ejected.
 Practical applications today in photovoltaics
(turning sunlight into electricity).
e.g.
calculators, cars, planes…
Example
Calculate Energy of a quantum of radiant
energy with a frequency of 5.0 x 1015 s-1
E=h
= (6.6262 x 10-34 J s)( 5.0 x 1015 s-1)
= 3.3 X 10-18 J
Light is pure energy!
Explanation of Atomic Spectra
 Bohr’s application of quantum theory to
electron energy levels resulted in explanation
of Hydrogen Spectrum.
 e.g. Excited electron goes back to ground
state and gives off photon of light of set “ ”

E=h
Emission
 Only electrons in transition from higher
energy to lower energy levels lose energy
and emit light.


e.g.
n=2
E=h
n=1
Hydrogen Spectra
 Various electromagnetic radiations emitted by
hydrogen:

Lyman (uv)
decay to n = 1

Balmer (visible)
decay to n = 2

Paschen (IR)
decay to n = 3
Hydrogen Spectral Emissions