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Concept #4
“Electrons in the Atom”
Honors Chemistry 1
The story continues,
• Rutherford discovered the nucleus
and proposed the “plum-pudding”
model of the atom.
• Scientists knew that the model was
flawed and worked on improving it.
• Their investigations shed “light” on
the problem.
Different elements give
different colors.
EM Waves
• Electromagnetic radiation is a
form of energy that exhibits
wavelike behavior as it travels
through space.
• Together, all the forms of
electromagnetic radiation form
the electromagnetic spectrum.
The Nature of Light
• Electromagnetic radiation,
including light, is described
using wave terms.
Crest
origin
Trough
Amplitude
Wavelength (l)
Frequency and
Wavelength
• Frequency (n) is the number of waves that
pass a point in a given amount of time.
• Hertz (Hz) are the units and 1 Hz is one
wavelength/second.
• Wavelength (l) is the distance from
corresponding points on adjacent waves.
• Light has a speed of 3.00 x 108 m/s (c).
• For light, c = l n, this also applies to all
electromagnetic radiation.
 l and n are inversely proportional.
Particle Nature of Light
• The photoelectric effect refers
to the emission of electrons
from a metal when light shines
on the metal.
• This phenomena did not fit with
the wave description of light.
Light can be a particle
also.
• Max Planck explained a “problem”
that scientists had with glowing hot
objects by suggesting that light
could be a particle.
• Matter could lose or gain energy in small
packets of energy called “quanta”.
• A quantum is the smallest amount of
energy that an atom can lose or
gain.
Planck’s Equation
• By applying a constant (Planck’s
constant) Planck found the
relationship between the energy of a
photon (a quantum of light) was:
• E = hn , where h is Planck’s
constant and E is the energy
measured in Joules.
• h = 6.626 x 10-34J.s
Planck’s Equation
• A photon is a particle of
electromagnetic radiation
having zero mass and carrying a
quantum of energy.
• The energy of a particular
photon depends on the
frequency of the radiation.
• Ephoton = hv
Rainbows are
“continuous spectra”.
• The refraction
and reflection of
sunlight by
water droplets
produces light in
the full or
continuous
spectrum.
Atomic Emission
Spectra
• Because energy must be released
only in certain size packages, the
light given off by exciting gases can
be resolved into a series of lines
with only certain colors of light.
Each acts like a “fingerprint”.
Different gases produce
different lights.
The Bohr Model
• Niels Bohr attempted to “fix”
Rutherford’s model by having the
electrons move in energy levels
around the nucleus.
• Each energy level could only hold a
certain number of electrons.
• The larger the energy level, the
further from the nucleus it was.
Bohr Model of the Atom
e- e-
e-
ee-
4
ee-
e3
e-
e-
e-
e-
1
ee-
ee-
e-
e-
2
e-
e-
ee-
e-
e-
e-
ee-
e-
The numbers represent the principal quantum
numbers.
Ground State Electrons
• When an electron is in the
lowest energy level possible it
is said to be in the ground state.
• When an electron doesn’t
occupy the lowest energy level
possible due to outside energy
it is said to be in the excited
state.
Changing Energy Levels
• When ever an electron moves to
a higher or lower energy level
an energy change is required.
• If the right amount of energy is
added, the electron moves up.
• To move down, a certain
amount of energy must be
released.
The Quantum
Mechanical Model
• Louis De Broglie combined
Planck’s equation E=hn and
Einstein’s E=mc2 to produce an
equation that explains why the
electrons can only occupy
certain energy levels.
l = h/mn
Heisenberg
• Werner Heisenberg proposed
that we could never tell the
position and momentum (speed)
of an electron at the same time.
• This uncertainty is the
“Heisenberg Uncertainty
Principle”.
Schrodinger
• Erwin Schrodinger derived a
formula that described the
space that an electron of a
certain energy would occupy.
Instead of an orbit, he
described an “electron cloud”
where you would have the
greatest probability finding the
electron.
Four Quantum numbers
for electrons in an atom.
• The principal quantum number (n)
describes the size and energy of the
electron orbital.
• Angular momentum quantum number (l)
describe the shape of orbitals. The
number of sublevels = n
• The magnetic quantum number(m)
describes orientation of the sublevels.
• Spin quantum number (s) refers to how an
electron spins.
Shapes of Orbitals
Quantum numbers
Electron configuration
• The arrangement of electrons in an atom
is known as the atom’s electron
configuration.
• The lowest energy levels are filled first,
this is the “Aufbau Principle”.
• Once an energy level is full electrons can
then fill the next highest energy level.
• “Pauli Exclusion Principle” states that no 2
electrons in the same atom can have the
same 4 quantum numbers.
Electron Sublevels
• s sublevels can have up to 2 e- and is
spherical.
• p sublevels can have up to 6 e- and
looks like 3 “dumbbells” in the x, y
and z axis
• d sublevels can have up to 10 e- and
the shape is more complicated
• f sublevels can have up to 14 e-.
p and d orbitals
Electron orbitals
• Electron orbitals can contain 0,
1 or 2 e-.
• s has 1 orbital, p has 3, d has 5,
and f has 7.
• We show orbitals as a box with
electrons represented as
arrows.
• Spin is represented by the
direction of the arrows.
Hund’s Rule
• Electrons will seek an
unoccupied orbital within a
sublevel before it will pair up
with another electron in an
orbital.
P sublevel
NO
P sublevel
YES
Orbital Filling
• The e- must fill the lowest
energy level first and only 2 ecan be placed per orbital.
• The first e- goes into the lowest
principle energy level, n=1.
• Level 1 only has 1 sublevel, s
• An s sublevel can only have 1
orbital
Orbital filling continued
• When the first energy level is
filled, the e- can occupy the
second energy level, n=2.
• Level 2 can have 2 sublevels, s
and p
• After sublevel s is filled then p
is filled according to Hund’s
rule.
The Flow Chart of
Science
• Electron filling follows the
following pattern:
• 1s 2sp 3sp 4s 3d 4p 5s 4d 5p 6s
4f 5d 6p 7s 5f 6d 7p
• Notice that each p is followed
by the next higher level’s s.
Electron configuration
• An easy way to represent the ein an atom is through electron
configuration.
•
2p
5
• The principle energy level is 2
and there are 5 e- in the p
sublevel.
Example:
• The electron configuration
for the sulfur (Z= 16) atom
is: 1s2 2s2p6 3s2p4
• This says that the total
number of e- is 16 and the
highest energy level is n=3.
Valence Electrons
• The chemical characteristics of
atoms are based upon the number of
e- that are in the outer energy levels.
• These are called valance electrons.
They are only s and p sublevel e-.
• For the sulfur atom, 1s2 2s2p6
3s2p4, there are 6 valence e-.
Electron Configuration
• Noble-Gas Notation
• The Group 18 elements (helium,
neon, argon, krypton, xenon, and
radon) are called the noble gases.
• A noble-gas configuration refers to
an outer main energy level occupied,
in most cases, by eight electrons.