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Chapter 5
Electrons in Atoms
Section 5.1—Light &
Quantized Energy
 In chapter 4, we discussed Rutherford’s
wonderful scientific discovery, the
nuclear model of the atom. It lacked one
thing—details of how electrons occupy
the space surrounding the nucleus.
 Electromagnetic radiation is a form of
energy that exhibits wavelike behavior
as it travels through space.
Section 5.1—Light &
Quantized Energy
 Wavelength () is measured from one
crest to the next crest, or from trough to
trough. Wavelength is usually expressed
in meters, centimeters, or nanometers
(1nm = 1x10-9 m).
 Frequency () is the number of waves
that pass a point per second. One hertz
(Hz) is the SI unit of frequency and is
expressed as s-1.
Section 5.1—Light &
Quantized Energy
 The amplitude is the wave’s height from
the origin to a crest, or from origin to a
trough.
 The speed of light (c) in a vacuum is
3.00 x 108 m/s. It has the formula:
c = 
 Wavelength & frequency are inversely
related, which means that as one
increases, the other decreases.
Section 5.1—Light &
Quantized Energy
 When white light is passed through a
prism, it is separated into a continuous
spectrum of colors.
 These visible colors only make-up a
small portion of the electromagnetic
spectrum (EM Spectrum)—includes all
the forms of electromagnetic radiation,
noting their differences in wavelength &
frequency.
Electromagnetic spectrum
 ***TRANSPARENCY
Section 5.1—Light &
Quantized Energy
 The sequence of the visible colors can
be remembered by the character Roy G.
Biv.
 In 1900, a German physicist named Max
Planck concluded that matter can gain or
lose energy only in small, specific
amounts called quanta—Quantum
concept.
Section 5.1—Light &
Quantized Energy
 That is, a quantum is the minimum
amount of energy that can be gained or
lost by an atom. This is demonstrated
by the equation:
Equantum = h
 Where E is energy, h is Planck’s
constant (6.626 x 10-34 J.s), &  is
frequency.
Planck’s Constant
• The frequency of an e- is proportional to
its energy.
Section 5.1—Light &
Quantized Energy
 In the photoelectric effect, electrons are
emitted from a metal’s surface when light
of a certain frequency shines on the
surface.
 What does this sound like?
 A calculator powered by photoelectric
cells converts energy from incident light
into electrical energy.
Photoelectric effect
Section 5.1—Light &
Quantized Energy
 In 1905, Einstein said that EM radiation
has both wavelike & particle-like
properties. That is, light can also be
described as a stream of particles called
photons.
 Photons are radiation with no mass that
carries a quantum of energy. Einstein
calculated that a photon’s energy
depends on its frequency:
Ephoton = h
Section 5.1—Light &
Quantized Energy
 The atomic emission spectrum of an
element is the set of frequencies of the
EM waves emitted by atoms of the
element. An atom’s emission spectrum
can be used to identify that particular
element.
Section 5.2—Quantum Theory
& the Atom
 When an atom is in the lowest energy
state, it is in the ground state.
 When it gains energy, it is in an excited
state.
 The smaller the electron’s orbit, the
lower the atom’s energy level.
Section 5.2—Quantum Theory
& the Atom
 Niels Bohr, a Danish physicist who worked
alongside Rutherford, in 1913, assigned a
quantum number (n) to each orbit & calculated
its radius.
 Bohr’s model explained the spectral lines of
hydrogen.
 Bohr’s idea of quantized energy levels laid the
groundwork for atomic models to come, but was
later disproved because it did not fully account
for the chemical behavior of atoms.
Section 5.2—Quantum Theory
& the Atom
 The movements of electrons in atoms
are not completely understood even
now; however, substantial evidence
indicates that electrons do not move
around the nucleus in circular orbits.
 In 1924, a French physicist named Louis
de Broglie derived an equation that
predicts that all moving particles have
wave characteristics.
 = h/m
Section 5.2—Quantum Theory
& the Atom
 In 1926, Austrian physicist Erwin Schrodinger
stated that electrons act as waves which led to
the quantum mechanical model of the atom.
Very similar to Bohr’s model, but is different in
that it makes no attempt to describe the
electron’s path around the nucleus.
 The quantum mechanical model of the atom
predicts a 3-dimensional region around the
nucleus called an atomic orbital describes the
electron’s probable location.
Section 5.2—Quantum Theory
& the Atom
 Because the atomic orbital is “fuzzy,” it
does not have an exactly defined size.
The quantum mechanical model assigns
principal quantum numbers (n) that
indicate the relative sizes & energies of
atomic orbitals.
 Therefore, n specifies the atom’s major
energy levels called principal energy
levels. Principal energy levels contain
energy sublevels named s, p, d, or f.
Section 5.2—Quantum Theory
& the Atom
 All s sublevels are spherical in shape
and all p sublevels are dumbbell shaped.
Not all d & f orbitals have the same
shape.
 Each orbital may have a maximum of 2
electrons. So the maximum number of
electrons related to each principal
energy level equals 2n2.
“S” sublevels
2nd energy level, “S” sublevel
“P” sublevels
5.3 Electron Configuration