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Chapter 5
 Nuclear
atomic model did not explain
how the atom’s electrons are arranged
in the space around the nucleus.
 Did not explain the differences in
chemical behavior among the
elements.
 Scientists discovered that an element’s
chemical behavior is related to the
arrangement of the electrons in its
atom.
 Electromagnetic
radiation is a form of energy
that exhibits wavelike behavior as it travels
through space. (Ex: visible light, radio
waves, X rays)
 Wavelength is the shortest distance between
equivalent points on a continuous wave
(units: meters, centimeters, nanometers)
 Frequency is the number of waves that pass
a given point per second. (units: hertz; one
hertz (Hz) = one wave per second = s^ -1)
 Amplitude of a wave is the wave’s height
from the middle to the top or from the
middle to the bottom
 The
speed of light (c) is the product of its
wavelength (λ) and its frequency (ν).
 C = λν
 As wavelength increases, frequency
decreases and as frequency increases,
wavelength decreases.
 All electromagnetic waves have the same
speed, but have different wavelengths
and frequencies.
Higher Frequency
Lower Frequency
 Sunlight
is an example of white light.
 When sunlight passes through a prism, you
see a continuous spectrum of colors. (visible
spectrum)
 Rainbows form when drops of water in the
air scatters white light from the sun into the
spectrum of colors that you see.
 The electromagnetic spectrum includes all
forms of electromagnetic radiation.
 The
sequence of the visible spectrum is red,
orange, yellow, green, blue, indigo, violet
(Roy G. Biv)
 As frequency increases, energy increases.
Energy increases     
 What
is the wavelength of a microwave
having a frequency of 3.44 X 109 Hz (or 1/s)?



C = λν
3. 00 X 108 m/s = λ (3.44 X 109 Hz)
λ = 8.72 X 10-2 m
 What
is the frequency of green light, which
has a wavelength of 4.90 X 10-7 m?
 An X-ray has a wavelength of 1.15 X 10-10 m.
What is its frequency?
 What is the wavelength of an
electromagnetic wave that has a frequency
of 7.8 X 106 Hz?
 The
temperature of an object is a measure of
the average kinetic energy of its particles.
 When objects are heated up, they gain more
energy.
 Max Planck studied why objects gave off light
when heated.
 He found that matter can gain or lose energy
only in small, specific amounts called
quanta.
 A quantum is the minimum amount of energy
that can be gained or lost by an atom.
 Planck
proved that the energy of a
quantum is related to the frequency of the
emitted radiation.
 E = hν
 E = energy; h = Planck’s constant;
ν = frequency
 Planck’s constant = 6.626 X 10-34 J·s
 Unit for energy = joules (J)
 Matter can only have certain amounts of energy
 No amounts of energy between these values
exist (Child building wall of blocks)
 In
the photoelectric effect, electrons called
photoelectrons are given off from a metal’s
surface when light of a certain frequency
shines on the surface.
 Einstein said electromagnetic radiation has
both wavelike and particle-like natures.
 So a beam of light has many wavelike
characteristics, but also is a stream of tiny
particles, or bundles of energy, called
photons.
 A photon is a particle of electromagnetic
radiation with no mass that carries a
quantum of energy.
 What
is the energy of a photon from the
violet portion of the rainbow if it has a
frequency of 7.23 X 1014 s-1 ?



E = hν
E = (6.626 X 10-34 J·s) x (7.23 X 1014 s-1)
E = 4.79 X 10-19 J
 What
is the energy of each of the following
types of radiation?



6.32 X 1020 s-1
9.50 X 1013 Hz
1.05 X 10 s-1
 The
atomic emission spectrum of an element
is the set of frequencies of the
electromagnetic waves emitted by atoms of
the element.
 Only certain colors appear in a certain
element’s atomic emission spectrum so only
certain frequencies of light are emitted.
 Bohr
studied the hydrogen atom to learn
about energy states.
 The lowest energy state possible of an atom
is called its ground state.
 When an atom gains energy, it is said to be in
an excited state.
 Bohr also suggested that electrons move
around the nucleus only in certain circular
orbits.
 The smaller the electron’s orbit, the lower
energy level the atom is in.
 The larger the electron’s orbit, the higher
energy level the atom is in.
Bohr
assigned a quantum number to
each orbit.
The quantum numbers are n=1, n=2,
n=3, and so on.
Bohr’s model was the foundation for
atomic models that came later, but
his model was actually wrong
because electrons do not move
around the nucleus in orbits.
De
Broglie studied to see if electrons
(particles) could behave like waves.
His equation is λ = h/mv
Wavelength = Planck’s constant/
mass x volume
His equation predicts that all moving
particles have wave characteristics
 The
Heisenberg uncertainty principle says
that it is impossible to know both the
velocity and position of a particle at the
same time.
 Schrodinger’s quantum mechanical model
of the atom describes atomic orbitals,
which are regions around the nucleus
where electrons can be.
 Atomic orbitals do not have an exact size.
 Principal
quantum numbers indicate the size
and energies of atomic orbitals.
 The letter ‘n’ represents an atom’s major
energy levels called principal energy levels.
 Principal energy levels contain energy
sublevels.
 Principal energy level 1 has 1 sublevel.
 Principal energy level 2 has 2 sublevels.
 Principal energy level 3 has 3 sublevels and
so on.