 We
have previously described atoms using
three different models.
 As a group,
1) Build a replica of the model of the atom
as described by the scientist your group is
assigned.
2) Discuss in your group the experimental
evidence that led to your model.
 Dalton
gave us the basis of the modern
atomic theory.
- All matter are composed of tiny indivisible
particles called atoms.
- Atoms of the same element are identical.
Atoms of different elements are different.
- Atoms combine in simple, whole number
ratios to form compounds.
- Atoms are separated, rearranged, and
recombined during chemical changes.
 J.J.
Thomson discovered the electron, a
negatively charged particle, through his
cathode ray tube experiment.
 This discovery resulted in a revised model of
the atom called the plum pudding model.
- Atom was made of a positively charged
material.
-Negatively charged electrons were
distributed throughout the atom.
 After
Rutherford’s discovery of the positively
charged nucleus through his gold foil
experiment, another revision of the model of
the atom occurred.
 He proposed that the protons were
concentrated in a small dense nucleus and
the electrons were located in the large space
surrounding the nucleus.
 Bohr,
a student of Rutherford, questioned
how the electrons, being oppositely charged
from the protons, did not fall into the
nucleus due to the attraction that forms
between opposite charges.
 Bohr proposed that the negatively charged
electrons must be moving in circular orbits
around the nucleus, much like planets orbit
the sun.
 This model was referred to as the planetary
model.
 Bohr’s
model did not address the differences
and similarities in chemical behavior among
the various elements.
 In the early 1900’s, chemists observed that
certain elements emitted visible light when
heated in a flame or exposed to other energy
sources.
 Analysis of the emitted light revealed that
this behavior was related to the electron
arrangement within its atoms.
 In order to understand this relationship, we
must first review the nature of light.
 Visible
light is a type of electromagnetic
radiation-a form of energy that exhibits
wavelike behavior as it travels through
space.
 Other forms of electromagnetic radiation
include TV/radio waves, microwaves,
infrared waves, ultraviolet waves, x-rays,
and gamma rays.
 All of these types of waves together make up
the electromagnetic spectrum.
 The
only differences among the different
types of waves making up the
electromagnetic spectrum are the
frequencies and wavelengths.
1) wavelength (represented by the Greek
letter lambda, λ) is the shortest distance
between equivalent points on a continuous
wave. Wavelength is usually expressed in
meters or nanometers (1 nm = 1 x 10-9 m)
The wavelength of visible light is between
400 and 700 nm.
2) frequency (represented by the Greek
letter nu, ν) is the number of waves that
pass a given point per second. Frequency is
measured in hertz, Hz. One hertz equals one
wave per second.
What do you notice about the
relationship between the frequency
and the wavelength of a wave?
It is an inverse relationship: as the
frequency increases the wavelength
decreases.
3) Amplitude of a wave is the wave’s height
from the origin to a crest, or from the origin
to a trough. Wavelength and frequency do
not affect the amplitude of a wave.
4) All electromagnetic waves, including visible
light, travel at a speed of 3.00 x 108 m/s in a
vacuum. This value is often referred to as the
speed of light. Because this is such an
important and universal value, it is given its
own symbol, c.
 The speed of light = wavelength x frequency.
 Calculate
the wavelength of the yellow light
emitted by the sodium in the pickle if the
frequency of the radiation is 5.10 x 1014 Hz.
 c = νλ
3.00 x 108 = 5.10 x 1014 (λ)
λ = 5.88 x 10-7 m or 588 nm
 The red-colored light given off by the
compound held in the flame has a
wavelength of 6.50 x 10-7 m. What is the
frequency of such red light?
 ν = 4.62 x 1014 Hz
 Sunlight,
our most common form of light is
continuous and covers all of the wavelengths
in the visible spectrum.
 It can be separated by passing the light
through a prism.
 This model cannot explain, however, why
heated elements emit only certain
frequencies of light at a given temperature.
 Scientists realized that a new model or
revision of the wave model of light was
needed.
 When
objects and elements are heated or
exposed to energy, they emit glowing light of
different colors.
 These colors correspond to different
frequencies and wavelengths.
 A German scientist, Max Planck, studied this
light phenomenon and came to a conclusion:
matter can gain or lose energy in only small,
specific amounts called quanta.
 A quantum of energy is the minimum amount
of energy that can be gained or lost by an
atom.
Prior to this discovery, physicists thought that
energy could be absorbed and emitted in varying
quantities, with no minimum limit to the
amount.
 Example: think about heating a cup of water on
a hot plate. The temperature on the
thermometer seems to rise continuously, not in
small jumps.
 Actually, the water’s temperature is increasing
in very tiny steps as the water molecules absorb
quanta of energy.
 Because these steps are so small, the
temperature just appears to be continuous.

 Planck
demonstrated mathematically that a
relationship exists between the energy of a
quantum and the frequency of the radiation
that is emitted.
 E = hν
 h = 6.626 x 10-34 J s (the joule, J, is a unit of
energy)
 In 1905, Einstein described these quanta of
energy as particles that he called photons.
 What
is the energy of a photon from the
violet portion of the Sun’s light if it has a
frequency of 7.23 x 1014 Hz.
 E = hν
 E = 6.626x10-34 (7.23 x 1014)
 E = 4.79 x 10-19J





Recall that Bohr stated that electrons exist on energy
levels.
When enough energy is added to an atom, electrons
jump to a higher energy level.
(This is sometimes referred to as a “quantum leap” or
the electrons are said to be in an excited state)
This is an unstable situation and the electron will
drop back to its former energy level, the ground
state, releasing the absorbed energy in the form of
light.
The color of light emitted depends on what energy
level an electron is dropping to.
A revision in Bohr’s model of the atom was necessary
to explain the relationship among atomic structure,
electrons, and the unique frequencies of light
emitted.