Unit 3 - Atomic Structure and
Chapter 4 pp. 83-91
Chapter 12 pp. 323 - 328
• “Atomists” - thought that
matter was composed of
tiny indivisible particles
• Democritus - first
suggested the idea of
• Felt that atoms were
fundamental units of
• Democritus had a good idea, but did
not prove his theory with hard facts
• Early Greek and Roman philosophers
often theorized with little or no
• Not very useful for scientific
• English school teacher
determine the ratio
elements combine in
• Created “The Atomic
Postulates of Dalton’s
1) All elements are composed of substances of
submicroscopic indivisible particles called atoms
2) Atoms of the same element are identical. Each
atom of an element is different from another
3) Atoms of different elements can physically mix or
chemically combine in whole-number ratios to
4)Atoms of one element can never change into atoms
of another element
• The smallest particle of an element that
retains the properties of that element
• Consider the Aluminum foil experiment...
Many of you found anywhere from 1 x 104
to 1 x 106 atoms of Al in the width of the
• Imagine how many atoms of Aluminum
would be in a soda can...
• Most of Dalton’s Atomic Theory is
• However, there is one revision…
• Atoms can be broken down into even
• Electrons, Protons, and Neutrons
Sir J.J. Thomson
• 1897 - discovered
• Used a cathode ray
tube to separate
particles of different
Cathode Ray Tube Experiment
• Found that cathode
rays were attracted to
metal plates that carry
a positive charge
• Similar to magnets
• The rays were repelled
by plates with a
• Thomson was able to prove that all cathode rays
are composed of e-
• Realized that electrons must be part of atoms
• Found that the electron carries one unit of
• 1/2000 the mass of H
Ray is bending due to negatively
• Each side of the tube charged item here
has a metal disc that is
either positive or
• The ray is negatively
charged and moves
from the cathode
toward the anode
• Approximately calculated
the charge and mass of an
• An electron has a mass of
1/1840 the mass of a
• In terms of mass that is
9.11 x 10–28 g
Protons and Neutrons
– Positively charged particle
– Has a mass of 1 amu (atomic mass unit) or
1.67 x 10-24 g
– Neutral or no charge
– Also has a mass of 1 amu (atomic mass unit) or
1.67 x 10-24 g
Ernest Rutherford and the
Gold Foil Experiment
• 1911 – Manchester,
• Rutherford initially
believed that the atom was
• Tested his theory of atomic
structure by creating an
experiment that proved
• He got a big surprise!
The Gold Foil Experiment
The Gold Foil Experiment
experiment proves that
the atom has a
• He finds that some of
the alpha particles
• Let’s take a look at
this on a smaller
• Rutherford was so
amazed by the outcome
of the experiment he was
quoted as saying,
“It was almost as if you
fired a 15-inch shell into
a piece of tissue paper
and it came back and hit
• The central core of an atom, composed of
protons and neutrons.
• Very dense
• Occupies a very small space compared to
the rest of the atom
• Most of the small alpha particles fired by
Rutherford were able to pass through the
Just how big is the nucleus?
• Most of the mass of an
atom is located in a very
• In fact it takes up about
1/10,000 of the radius of
the whole atom.
• If we could compare an
atom to a football
stadium, the nucleus
would be the size of a
marble in the middle of
the stadium field.
What about the rest of the atom?
• Electrons are located in the space around
the atom, but were so tiny they did not crash
into the alpha particles.
• Rutherford realized that because most of the
alpha particles passed through the gold foil,
much of the atom must be empty space.
Models of the Atom
• After discovering the electron, Sir J.J. Thomson
realized that the atom was not a hard ball of
matter, as was previously thought.
• Up to this point, atoms were considered indivisible
using Dalton’s Atomic Theory.
• Thomson reasoned that electrons were located in
the atom in a kind of “Plumb Pudding”
The Plumb Pudding Model
• Thomson believed that
the atom was made up
of electrons that were
“stuck” into a kind of
lump of protons
• This model explained
some of the electrical
properties of the atom,
but nothing really
about it’s arrangement
• After his gold foil
proposed that electrons
surrounded a dense
nucleus that was
actually in the center of
• The rest of the atom was
• Later, it was shown that
neutrons also composed
• The Rutherford model was one of the best
in it’s time.
• However, scientists realized that positively
and negatively charged particles would tend
to be attracted to one another.
• What kept the particles from crashing into
each other in a “magnetic” attraction?
The Bohr Model
• In 1913, one of Rutherford’s students, a
physicist named Niels Bohr, proposed that
electrons were arranged in circular paths
around the nucleus.
• Also known as the “Planetary Model” Bohr
proposed that the electrons orbited around
the nucleus like planets around the sun.
• Each electron had a fixed amount of energy
that kept it in a path around the nucleus.
The Bohr Model cont’d.
• Each electron moved
within an energy level.
• Energy levels were
regions around the
nucleus where an
electron was likely to
• Bohr compared energy levels to rungs on a ladder.
• The lowest energy levels were at the bottom of the
• Electrons could move up and down the energy
levels just like a person on a ladder.
• Electrons could not be located between energy
levels, just like a person cannot stand in the empty
space between rungs on a ladder.
• To move from one energy level to another, an
electron had to gain or lose just the right amount
of energy; a person climbing a ladder must move
his/her feet just the right distance.
Ladders and Energy Levels
• Notice that the bottom
of the ladder, which is
the lowest energy
level, is closest to the
• The more energy an
electron gains, the
higher up it moves on
• The opposite can be
said if it moves down
The “Quantum” Idea
• Scientists have labeled the amount of
energy needed to move from one energy
level to another as “QUANTA”
• When an electron gains or loses quanta
(packets) of energy, the electrons move up
or down the ladder
• Electrons that have energy are said to be
The difference between
Energy Levels and Ladders
• One main difference between energy levels
and ladders is that energy levels are not
• Energy levels get closer together as the
electrons get farther away from the nucleus;
kind of like an escalator rather than a ladder.
• It’s easy to get off at the top of an escalator.
A person can just step off...
• Electrons also have an easier time escaping
from an atom at the top most energy levels.
The Quantum Mechanical Model
• 1926 - Developed by
• Uses the quantum theory
to write a mathematical
equation to describe the
location and energy of an
electron in a hydrogen
The Quantum Mechanical Model
• Considered the most modern description of
• Abstract and mathematical, this new model
restricts the energy of electrons to specific
values like Bohr.
• Differs from Bohr by estimating the
location of an electron rather than assuming
that the e- is following a given path.
• Schrödinger uses the idea of probability to
estimate the electron (e-) location.
• Example: If you have 3 red balls and
1 green ball in a bag. You have a 25%
chance of picking the green ball and a 75%
chance of picking a red ball.
• The same principle can be applied to
locating an electron in a given space.
Calculating the Probability
of finding an Electron
• The probability of
finding an e- in a
certain atomic volume
is based upon the
density of e- in the
• The higher the e-
density, the higher the
probability of finding
• The opposite applies to
low e- density
• No identifiable line around the nucleus
• Fuzzy cloud indicates the likelihood of
finding an electron 90% of the time.
– If you could put a bag around the area an e- is
located, you would find the e- 90% of the time
in the bag.
The Principle Quantum Number
& Energy Sublevels
• The P.Q.N. (n) – identifies the energy level
that an electron is located; kind of like a
section of a movie theater where an e- is
• As the distance from the nucleus increases,
so also the Energy level where n = 1,2,3,4,5,
• In each Energy Level, an e- occupies an
energy sublevel, like a row in a section of a
• Electrons are limited by the Quantum
Mechanical Model to cloud shapes as we
• We call these cloud shapes orbitals
• Orbitals are the most likely location of e-
• We denote orbitals by letters instead of
• We will be looking at
two different types of
orbitals (s & p)
• s - orbital types are
spherical shaped and
can hold only 2 e-.
• These orbital types are
filled first with e- on
every energy level
• p – orbitals are filled after s – orbitals and can
hold up to 8 e-
• There are many shapes for the p-orbital
• We call these shapes dumbbells
Basic Nomenclature for
• To understand where e- are located we write
e- configurations using a special
2s 2 Number of e-
We can use the Periodic Table