Electron Configuration
and New Atomic Model
Chapter 4
The Key to the New Atomic Model
• A connection between light and electrons!
• Before the 20th century, it was believed that
light behaved only like a wave.
• To fully understand the connection between
light and electrons, we must review the
properties of light.
Electromagnetic Radiation
• A form of energy that exhibits wavelike behavior as
it travels through space is called electromagnetic
radiation.
• All forms of this radiation make up the
electromagnetic spectrum.
• All forms move at the speed of light
(c= 3.0 x108 m/s), through a vacuum and slightly
slower through matter.
• We can assume that it moves that fast through air
because it is mostly empty space.
Electromagnetic Spectrum
Features of Light
• Wavelength ()- is the distance between
corresponding points on adjacent waves.
• Units are either in meters, centimeters or
most commonly nanometers.
• 1 nm= 10-9 m
Features of Light
• Frequency ()- the number of waves that
pass a given point in a specific amount of
time, usually in one second
• Unit- Hertz (waves/second)
• Higher frequency equals shorter wavelength.
• Lower frequency equals longer wavelength.
• They are related by the following equation:
c=
The Photoelectric Effect
• One experiment completed in the early 1900s
challenged the wave theory of interaction between
light and matter.
• The photoelectric effect refers to the emission of
electrons from a metal when light shines on it.
• The wave theory of light predicted that any
frequency of light would supply enough energy to
eject an electron.
• However, in this experiment, electrons weren’t
emitted if the light’s frequency was below a certain
level.
Particle Description of Light
• Max Planck- a German physicist who studied
the emission of light from hot objects
• Hot objects do not continuously emit
electromagnetic radiation, as they would if
they were in the form of waves.
• Planck suggested that hot objects emitted
energy in small, specific amounts called
quanta.
Relationship between Quanta and
Frequency of Radiation
• A quantum is the minimum energy lost or
gained by an atom. (plural=quanta)
• Planck’s proposed relationship:
E=h
• E= energy in Joules of a quantum of
radiation,  is the frequency, h is a constant,
known as Planck’s constant: 6.626 x 10-34
J•s.
Wave-Particle Duality
• Einstein took Planck’s idea further and proposed that
electromagnetic radiation has a dual wave and
particle nature.
• Light has wave-like properties but can also be
thought of as a stream of particles.
• Each particle of light carries a quantum of energy.
• He called these particles photons.
• A photon is a particle of electromagnetic radiation
having zero mass and carrying a quantum of energy.
Einstein explains the Photoelectric
Effect
• Electromagnetic radiation is absorbed by matter only
in whole numbers.
• In order for an electron to be ejected from a metal
surface, it must be struck by a photon that has the
minimum required energy.
• This corresponds to a minimum frequency.
• Since electrons are bound more or less closely,
depending on the metal, the frequency required to
remove an electron is different for different metals.
Energy of Atoms
• The lowest energy state of an atom is called
its ground state.
• The state in which the atom has a higher
potential energy than its ground state is
called its excited state.
• When an excited state atom returns to the
ground state, it gives off energy in the form
of electromagnetic radiation, or light.
Emission Spectra
• Emission spectra are bands of light at
specific frequencies (and therefore
wavelengths) that result when the light
emitted from an element.
• This spectra can be observed when the
element is in gaseous form and has a current
running through it or when it is being
burned.
Emission Line Spectra of Hydrogen
• What was expected to
be observed from
hydrogen was a
continuous spectrum.
• Since this was not
observed, attempts to
explain this resulted in
a new theory of the
atom called quantum
theory.
Fixed Energy
• Because hydrogen
• When an excited
emits light at specific
hydrogen atom returns
frequencies, the energy
to its ground state, a
differences between
photon of radiation is
these states must be
emitted.
fixed.
• The energy of the
photon is equal to the • So a new model of
hydrogen must be
difference between the
created!
initial and final states.
Bohr Model of Hydrogen
• Niels Bohr proposed a
model of the atom that
linked it’s electron with
the photon emission.
• According to his model,
the electron could circle
the nucleus in allowed
paths called orbits.
• The electron is in
lowest orbit when
closest to the nucleus,
and therefore has the
lowest energy.
• When the energy of the
electron becomes
higher, it orbits further
from the nucleus.
Bohr’s Results
• He matched the
spectral lines produced
for hydrogen to the
energies that hydrogen
would allow for
electrons.
• The mathematics
supported the
experimental data!
• The discovery of the
new model for the
hydrogen atom was
thought to apply to all
other atoms.
• However, it was soon
discovered that his
model did not work for
atoms with more than
one electron.
Quantum Model of the Atom
• Scientists struggled
with the notion that
electrons could only
exist in certain orbits
with definite energies.
• They did not
understand why there
couldn’t be a limitless
number of orbits with
slightly different
energies.
• Louis de Broglie, a
French scientist,
investigated this very
question in 1924.
• What he found would
change our
understanding of
matter forever.
Quantum Model of the Atom
• De Broglie suggested
that electrons be
considered waves that
are confined to the
spaces around the
nucleus.
• His theory about
electrons behaving like
waves was soon proven
to be correct by
experimentation.
• Electrons can be bent
or diffracted.
• Diffraction refers to
the bending of a wave
as it passes by the
edge of an object.
• It was also shown that
interference occurs
when waves overlap,
causing a slight
decrease in energy.
Heisenberg Uncertainty
Principle
• Scientists wondered
where electrons were in
the atom, because they
could not be detected
with photons since the
energy of a photon
would send an electron
off its current path.
• The Heisenberg
Uncertainty principle
says that it is
impossible to determine
simultaneously both the
position and velocity of
an electron or any
other particle.
Schrödinger Wave Equation
• Schrödinger developed
an equation based on
the theory that
electrons have a dual
wave-particle nature.
• Together with the
Heisenberg uncertainty
principle, the new
equation led to modern
quantum theory.
• Quantum theory
mathematically
describes the wave
properties of electrons.
• Quantum theory
determined that wave
equations give only the
probability of finding an
electron at a given
place.
Quantum Numbers
• Numbers used to
specify the properties
of atomic orbitals and
properties of electrons
in those orbitals
• There are 4 quantum
numbers.
• First three are from
Schrödinger's equation
• 1) Principal Quantum
number- (n) indicates the
main energy level occupied
by an electron
-all are whole
numbers
- as number increases
the average distance
from the nucleus
increases
Quantum Numbers
• 2) Angular Momentum
Quantum Number- (l)
indicates the shape of an
orbital
• The number of orbital
shapes possible is equal to
n-1.
• Zero counts…
• So for a value of n=2, there
are a total of 2 l values, l=0
and l=1.
• The value of l is then
assigned a letter:
• L=0
s
• L=1
p
• L=2
d
• L=3
f
Orbital Shapes
• An s orbital is spherical.
• A p orbital is dumbbell
shaped.
• A d orbital can be
clover shaped.
• An f orbital is really
complex.
Quantum Numbers
• 3) Magnetic
Quantum numbers(m) indicates the
orientation of an orbital
around the nucleus
• In each s sublevel,
there is only 1
orientation because it’s
a sphere.
• In a p sublevel, there
are 3 orientations
because the dumbbell
can be aligned on the
x, y or z axis.
• In a d sublevel, there
are 5 orientations.
• In an f sublevel, there
are 7 orientations.
Spin Quantum Numbers
• Electrons can be
thought to spin on an
internal axis.
• It can spin in one of
two possible directions.
• The spin quantum
number can have only
2 possible values: +1/2
or -1/2, which indicate
direction of spin.
• A single orbital can hold
a maximum of 2
electrons which must
have opposite spins!!!