Chapter 5
Section 5.1
Electromagnetic Radiation
Section 5.2 Atomic Spectra and
Energy Levels
Problems with the Rutherford
Nuclear Model:
Where are the electrons?
 Why don’t the electrons get pulled into the
nucleus?
 No explanation for spectral lines.

*Light behaves as both a wave and a
particle.
Sometimes light behaves like a
wave.

Visible light is only a small part of the
electromagnetic spectrum.
Electromagnetic Spectrum:
All the forms of energy that exhibits wavelike behavior as it
travels through space.
Examples:
 Gamma rays (short wavelength, high frequency)
 X-rays
 Ultra violet
 Visible (from short to long wavelength - Violet,I,B,G,Y,O,Red)
 Infrared
 Microwave
 TV
 Radio (long wavelength, low frequency)
Speed of all electromagnetic radiation (light): c = 3.0 x
108m/s)
Wavelength (λ):
The distance between corresponding parts
of adjacent waves, e.g. crest to crest.
 Usually in units of nm
Frequency (ν):
Number of waves that pass a given point in
a given time
 Usually in waves/second or Hertz (Hz)
Speed of Light (c):
Speed of light = wavelength x frequency
c = λv
What is the relationship between the wavelength
and frequency?
Inverse relationship!
- the shorter the wavelength the higher the
frequency (as λ ↑ , v ↓)
- the longer the wavelength the lower the
frequency (as λ ↓, v ↑)
What is the frequency of a light wave with a
wavelength of 106 nm?
Solution:
1. Convert nm to m:
106 nm x (1m/1,000,000,000nm) = 1.06 x 10-7m
2. Solve the light equation for v.
c = λv
v =c/λ
= 3.0 x 108 m/s / 1.06 x 10-7m
= 2.83 x 1015 Hz
Sometimes light behaves like a
particle.
Photoelectric effect: the emission of
electrons from a metal when light shines
on the metal.
 http://wps.prenhall.com/wps/media/objects
/439/449969/Media_Portfolio/Chapter_05/
Photoelectric_Effect.MOV

Max Planck
Light energy exists in specific small
amounts called quanta (quantum –
singular)
 quantum: the minimum quantity of energy
that can be lost or gained by an atom.
E=hv
 Planck’s constant, h = 6.626 x 10-34 J•s

Determine the energy in joules of a photon
whose frequency is 3.55 x 1017 Hz.
Solution:
E = hv
E = (6.626 x 10-34 J•s) (3.55 x 1017 Hz)
= 2.35 x 10-16 J

Einstein:
Won the Nobel Prize for explanation of
photoelectric effect.
 Einstein explained the photoelectric effect
by proposing that electromagnetic
radiation is absorbed by matter only in
whole numbers of photons.

Photon:
Massless particle of light with a quantum
of energy.
 The energy of the photon depends on the
frequency of radiation. Only certain
frequencies are energetic enough to knock
electrons off the surface of metals.

Flame Test:

Metal ions produce characteristic colors in
a flame.

NaCl
SrCl2
BaSO4
Line Emission Spectrum:

Produced when the light given off by an
element is separated into a set of colored
lines that are a unique “finger print” of the
element.
Bohr Model of the Atom
Niels Bohr applied quantum theory to the
Rutherford model.
1. Electrons exist only in definite orbits.
Each orbit has a definite fixed energy.
Labeled n = 1,2,3,…etc. Lowest energy
level is called the ground state, closest to
the nucleus.
2. As long as electrons stay in one energy
level they don’t gain or lose energy.
Bohr Cont.
3. Adding energy can put atoms in an excited
state, the electrons move to a higher energy
level, farther from the nucleus.
4. Electrons can “fall” back from higher (excited) to
lower energy levels. The difference in energy is
released as electromagnetic radiation (light).
Explains the color of the lines in the hydrogen
spectrum.