Atomic Structure
CHAPTER 4
4.1 STUDYING ATOMS

If you cut a piece of Aluminum foil in half and continue
to cut the resulting piece in half, what will happen?
Greek Philosophers pondered this 2500 years ago.

Democritus- believed all matter consisted of extremely small
particles that could not be divided.
 He

called these particles atoms
Aristotle – did not think there was a limit to the number of
times matter could be divided.
 Most
people believed Aristotle until the 1800s when scientists
had enough data to support Democritus
Aristotle thought that all substances were built up from
only four elements—earth, air, fire, and water. These
elements were a combination of four qualities—hot, cold,
dry, and wet. Fire was a combination of hot and dry. Water
was a combination of cold and wet.
DALTON’S ATOMIC THEORY

DALTON- born in England
in 1766.
 He
noticed no matter how
large or small a sample, the
ratio of masses of the
elements in a compound is
always the same.
 Dalton proposed a theory
that all matter is made up of
individual particles called
atoms, which cannot be
divided.
Magnesium reacts with oxygen to
form the compound magnesium
oxide. The ratio of magnesium to
oxygen, by mass, in magnesium
oxide is always about 3 : 2
MAIN POINTS OF DALTON’S ATOMIC THEORY
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All elements are composed of atoms.
All atoms of the same element have the same mass, and atoms of
different elements have different masses.
Compounds contain atoms of more than one element.
In a particular compound, atoms of different elements always
combine the same way.
In Dalton’s model, he thought elements were solid spheres. Each
type of atom is represented by a tiny, solid sphere with a different
mass.
Eventually, scientists discovered not all of Dalton’s theories were
correct.
THOMSON’S MODEL OF THE ATOM

J.J. Thomson- Joseph John
Thomson 1856-1940

Atoms have positive and negative
charges.
 Objects
with like charges repel, or push
apart.
 Objects with opposite charges attract or
pull together.
Some charged particles can flow
from one location to another (electric
current)
 Thomson used an electric current to
learn more about atoms.

CATHODODE RAY TUBE
If amber is rubbed with
wool, it becomes
charged and can attract
a feather.
CATHODE-RAY TUBE EXPERIMENT
A cathode-ray tube is a sealed tube with a metal
disk at each end. One is positive and one is
negative.
 A glowing beam appears between the two disks.

Thomson used a sealed
tube of gas in his
experiments. When the
current was on, the disks
became charged and a
glowing beam appeared in
the tube
CATHODE RAY TUBE EXPERIMENT
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Thomson discovered the beam was deflected when
additional charged plates were placed on the sides of the
tube.
Thomson concluded the beam must be negative charges.
He hypothesized the charges came from inside the atom.
Thomson’s experiments provided the first evidence that
atoms are made of even smaller particles.
The beam bent
toward a positively
charged plate placed
outside the tube
THOMSON’S MODEL

His model of the atom
looked like plum
pudding (or chocolate
chip ice cream). The
pudding had an overall
positive charge and the
negative charges were
randomly placed
throughout. Overall, the
atom is neutral.
RUTHERFORD’S HYPOTHESIS
1899 Ernest Rutherford discovered Uranium
emits fast-moving particles that have a positive
charge. (He called them alpha particles)
 1909 he asked his student, Ernest Marsden, to
see what happens when the alpha particles are
passed through a thin sheet of gold.

 He
hypothesized most particles would travel in a
straight path from their source.
 Some would be deflected slightly.
RUTHERFORD’S GOLD FOIL EXPERIMENT
GOLD FOIL EXPERIMENT
RUTHERFORD’S ATOMIC THEORY
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ACTUAL EXPERIMENTAL RESULTS
 More particles were deflected than he was
expecting. Some particles deflected as
much as 90º. Others bounced straight
back.
DISCOVERY OF THE NUCLEUS
 Nucleus- dense, positively charged mass
located in the center of the atom.
 Rutherford proposed a new model of the
atom.
 All of the atom’s positive charge is
concentrated in its nucleus. This
explains why alpha particles had a
greater deflection the closer they were
to the nucleus (both have positive
charges)
If the stadium were a
model for an atom, a
marble could represent
its nucleus.
4.2 STRUCTURE OF AN ATOM
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Properties of Subatomic Particles
Protons, electrons, and neutrons are all subatomic
particles
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PROTONS- Positive charge subatomic particle found in the
nucleus. They each have a charge of 1+. Each nucleus
contains at least one particle with a positive charge.
ELECTONS- Negatively charged subatomic particle that is
found in the space outside the nucleus. Each electron has a
charge of 1-.
NEUTRONS- Neutral subatomic particle that is found in the
nucleus of an atom. It’s mass is nearly equal to the mass of
a proton.

In 1932- James Chadwick designed an experiment to show neutrons
exist. It was similar to Rutherford’s gold foil experiment. The
neutrons showed no deflection.
COMPARING SUBATOMIC PARTICLES

Protons, electrons and neutrons can be distinguished by
mass, charge, and location in the atom.
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Protons and neutrons have equal mass.
Electrons are 1/2000 the mass of a proton.
Electrons have a charge that is equal in size to, but the opposite
of, the charge of a proton.
Neutrons have no charge.
Protons and neutrons are found in the nucleus, but electrons are
found in the space outside the nucleus.
ATOMIC NUMBER AND MASS NUMBER

Atomic number- Equal to the number of protons in an
atom of that element.
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Hydrogen (H) atoms are the only atoms with 1 proton.
Atoms of different elements have different numbers of
protons.
Each positive charge is balanced by a negative charge.
Each element has a
different atomic
number. A The atomic
number of sulfur (S) is
16. B The atomic
number of iron (Fe) is
26. C The atomic
number of silver (Ag) is
47.

Mass number- Sum of the protons and
neutrons in the nucleus of that atom.
#

of neutrons = mass # - atomic #
Isotopes
 Every
atom of a given element does have the same
number of protons and electrons. (electrons when in neutral state)
 But every atom of a given element does not have
the same number of neutrons.
 Isotopes of an element have the same atomic
number but a different mass number because they
have different numbers of neutrons.

With most elements, it’s hard to notice any differences
in the physical or chemical properties of their isotopes.

Hydrogen is the exception.
Normal hydrogen (H-1) has no neutrons (most of all H)
H-2 has 1 neutron –mass has doubled.
H-3 has 2 neutrons –mass has tripled.
Heavy water is made from H-2 atoms.
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4.3 MODERN ATOMIC THEORY
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BOHR’S MODEL OF THE ATOM
 Niels
Bohr’s model did something Rutherford’s
model did not do. It focused on the electrons.
 Electrons
move with constant speed in fixed orbits
around the nucleus (like planets around the sun)
 Each electron in an atom has a specific amount of
energy.
 ENERGY LEVELS- the possible energies that electrons in
an atom can have.
UNDERSTANDING ENERGY LEVELS

Picture energy levels as steps in a staircase.
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You can go up or down the steps, but only in whole-step
increments. You cannot stand between steps on a
staircase. Electrons cannot exist between energy
levels.
An electron in an atom can move from one energy
level to another when the atom gains or loses
energy.
The size of the jump determines the amount of energy
gained or lost.
 Energy released as the electron jumps back down to its
lower energy levels is often given off in the form of
visible light.
 Different elements emit different colors of light.

ELECTRON CLOUD MODEL
Bohr was incorrect in assuming electrons moved
like planets in a solar system. They are actually
less predictable.
 ELECTRON CLOUD- a visual model of the most
likely locations for electrons in an atom. The cloud
is denser at the locations where the probability of
finding an electron is higher.
 Scientists use the electron cloud model to
describe the possible locations of electrons
around the nucleus.
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ELECTRON CLOUD ANALOGY
When the propeller of an airplane is at rest, you can see the locations of the blades.
When the propeller is moving, you see only a blur that is similar to a drawing of an
electron cloud
ATOMIC ORBITALS
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Orbital- is a region of space around the nucleus where an
electron is likely to be found.
An electron cloud is a good approximation of how electrons
behave in their orbitals
The level in which an electron has the least energy—the
lowest energy level—has only one orbital. Higher energy
levels have more than one orbital
ELECTRON CONFIGURATIONS

electron configuration is the arrangement of
electrons in the orbitals of an atom.
The most stable electron configuration is the one in
which the electrons are in orbitals with the lowest
possible energies.
 When all the electrons in an atom have the lowest
possible energies, the atom is said to be in its ground
state
 If one electron can move to an orbital with a higher
energy it is referred to as an excited state.
 An excited state is less stable than the ground state.
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