Chapter 4
The Structure of the Atom
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Early Theories of Matter
• The concept of the atom developed over
a few thousand years.
• Consider that there were no controlled
experiments and few tools for scientific
exploration.
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• The power of the mind and intellectual
thought were the primary ways to truth
about the universe.
• Curiosity drove discovery
• Philosophers, scholarly thinkers,
speculated about the nature of matter.
• Ideas came from their own life
experiences.
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Early Greek philosophers thought that matter
was made of Earth, Wind, Water, and Fire.
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• Commonly accepted that matter could
be divided endlessly into smaller and
smaller pieces…..
– The Continuous Theory of Matter
• No methods to test the validity of these
ideas.
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Democritus
• a Greek Philosopher
• the first to propose
the idea that matter
was not infinitely
divisible.
discontinuous theory
• He proposed the idea
that matter was
made up of tiny
individual particles
called atomos.
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• The English word for atom comes
from this Greek word atomos.
• He believed that atoms could not
be created, destroyed, or further
divided (discontinuous theory).
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• His belief in atoms was amazingly
accurate.
• However, his ideas were met with
much criticism from other
philosophers.
• The harshest criticism was from the
influential Greek philosopher,
Aristotle.
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Aristotle
• Aristotle rejected the
theory of atoms
because it did not
agree with own ideas
on nature.
• He did not believe
that atoms could
move through empty
space.
• He did not believe
that the
“nothingness” of
empty space could
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exist.
Democritus Rejected
• Because Democritus had no way to
answer the challenges to his ideas, his
theory was eventually rejected.
• It is important to realize that these ideas
were just that—ideas and not science.
• Incredibly, the ideas of Aristotle were so
great and the science so primitive that his
denial of the existence of the atom went
unchallenged for two thousand years
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John Dalton
• In the 19th century John Dalton
– an English School teacher
– revived and revised the
work of Democritus.
– This time it was based on
scientific research
conducted by Dalton.
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Dalton’s Atomic Theory
– 1808
– marks the beginning of modern atomic
theory.
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Dalton was wrong about:
1.the atom being divisible (atoms are
divisible into subatomic particles)
2.all atoms of an element having
identical properties. (atoms of an
element can have slightly different
masses.)
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Dalton’s Model of the Atom
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Defining the Atom
• An atom is the smallest part of an element that
retains the characteristics of the element.
– How small is a typical atom?
o In the year 2000, the world population was about
6 billion people.
o In comparison, a copper penny contain about 5
billion times as many atoms of copper.
•
•
•
•
World population
6 000 000 000
Atoms in a Penny 29 000 000 000 000 000 000 000
Diameter of a copper atom 1.28 x 10-10 m
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Next Set of Questions
1. What is an atom like?
2. How are atoms shaped?
3. Is the atom composed of other
particles?
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The Discovery of the Electron
• British scientist named J. J.Thompson
• Used a glass tube, called a
cathode-ray tube, connected to
two electrodes.
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Cathode Ray Tube
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• One electrode was attached to one
end of the tube.
• It was negative in charge and was
thus called the cathode.
• At the other end was attached a
positive electrode, the anode.
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Thomson’s Experiment
• Thomson placed a paddle wheel inside
the tube, between the cathode and the
anode.
• He passed an electric current through a
variety of gases in the tube. The gases
within the tube were changed with each
experiment.
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Thomson made the following
observations
1. The current passing through the tube creates
different colors in different gases.
2. The anode (positive end) glows where as the
cathode (negative end) does not.
3. An object placed within the cathode-ray casts
a shadow.
4. The paddle-wheel spins in the direction of the
anode.
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Conclusions:
1. The different colors that were created by using
different gases showed that atoms of different
elements possessed different energies.
2. The cast shadow was thought to be due to the
beam of light created by the cathode-ray.
However, the experiment made with the
spinning paddle-wheel showed that the
cathode-ray was composed of particles.
3. The fact that the anode glowed and the
cathode did not shows that the cathode-ray
travels from a negative potential to a positive
potential. This was one of the most important
findings of J. J. Thomson's experiment and was
repeated using another apparatus.
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• To test the polarity of the cathode ray,
Thompson replaced the paddle-wheel with
another pair of electrodes.
• When the current was passed through this new
apparatus, Thompson found that the cathoderay was "bent" towards the positive electrode
and repelled by the negative electrode.
• Thompson attributed the "bending" of the
cathode-ray to charge-charge repulsion of the
second cathode and the negatively charged
particles in the cathode-ray
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Electrons
• He named these negatively
charged particles electrons.
• Since Thompson could get these
results regardless of what element
he used within the tube, he
concluded that all atoms have
electrons.
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Thomson’s Model of the Atom
• The Plum Pudding Model
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Oil Drop Experiment
• In 1909, American scientist, Robert Millikan,
determined the charge of an electron in his
famous Oil Drop Experiment.
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Discovery of the Nucleus
• 1911 Ernst Rutherford
• Rutherford’s Gold Foil Experiment
– determined that nearly all of the mass
of the atom consisted of a positively
charged nucleus
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Rutherford’s Gold Foil Experiment
• He did this by measuring the deflection of helium
nuclei by gold atoms.
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Predicted Results
31
•
While most atoms were not deflected
much at all, a few were deflected by 180
degrees.
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How was this possible?
• Rutherford calculated that the only
way this was possible was if the gold
atoms consisted of a cloud of
electrons surrounding the very
dense, positively charged nucleus,
for only then could the gold atom
transfer enough momentum to the
ions to turn them around.
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Rutherford’s Conclusions
• Atoms are mostly empty space and
the nucleus is incredibly dense.
– If the nucleus were a marble, the atom
would be the size of a football field
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Rutherford ‘s Nuclear Model
• By 1919 Rutherford
refined his concept of
the nucleus and
concluded that it
contained protons.
• Protons have a
charge of +1.
• This model could not
account for the mass
of the atom.
• 20 more years before
this mystery was
solved.
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Discovery of the Neutron
• 1932
• James Chadwick, an English Physicist
• He showed that the nucleus also
contained another particle, the neutron
• The neutron is neutral, meaning it has no
charge.
• The neutron has a mass equal to the
proton.
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Composition of the Nucleus
• Protons
• Mass of the proton = 1.673 X 10-24grams
• Positive Charge
• Neutrons
• Mass = 1.675 X 10-24grams
• No charge
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Properties of Subatomic Particles
Properties of Subatomic Particles
Relative
charge
mass
Particle Symbol Electrical

Electron
e
1
1/1840
Actual
mass
(g)
9.111028
Proton
p
1
1
1.6731024
Neutron
n0
0
1
1.6751024
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Nuclear Model of the Atom
• An atom is an electrically neutral particle
composed of protons, neutrons, and
electrons.
• Atoms are spherical in shape with a tiny,
dense nucleus of positive charge
surrounded by one or more negatively
charged electrons
• Nucleus contains 99.7 % of the mass of an
atom.
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Subatomic Particles
• Most of the atom consists of fast-moving
electrons traveling through the space
around the nucleus.
• The number of protons is equal to the
number of electrons since atoms are
neutral in charge.
• The number of Protons in the nucleus
determines the identity of an atom.
(Atomic Number)
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How Atoms Differ
• Henry Moseley discovered that atoms of
each element contain a unique positive
charge in their nucleus.
• Therefore, the proton number of an atom
identifies the atom.
• The number of protons in an atom is
known as atomic number.
• Example: All Hydrogen atoms have one
proton
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Atomic Number(Z): the number of protons
in the nucleus of each atom of the element
• Atomic Number = Protons
• Elements are arranged in the
periodic table from left to right in
order of increasing atomic number.
• Atomic number identifies the
element.
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Atoms are neutral, therefore:
# of protons = # of electrons
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Isotopes and Mass Number
• Naturally occurring elements are a
mixture of atoms that have different
numbers of neutrons.
• Atoms with the same number of
protons, but different number of
neutrons are called Isotopes.
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Hydrogen Has Three
Isotopes:
• Protium
• Deuterium
• Tritium
1 proton
1 proton
1 proton
1 electron
1 electron
1 electron
0 neutrons
1 neutron
2 neutrons
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Isotopes and Mass
Number
• Isotopes: atoms of the same element that
have different masses.
• Different number of neutrons so mass is
different.
• Atoms of different isotopes have different
masses so identity is given by name and
mass.
• Mass Number: the total number of
protons and neutrons in the nucleus of an
isotope
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Isotope (Nuclear) Symbols:
Consists of three parts
1. the symbol of the element
2. the atomic number of the element
3. the mass number of the specific
isotope.
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Nuclear Symbol
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Isotope Names
Element Name-Mass Number
Examples:
Helium-4
Potassium-39
Hydrogen-3
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Neutrons = Mass Number – Atomic Number
Atomic Number: _________
Mass Number: __________
# of Protons: _________
# of Electrons: _______
# of Neutrons: ________
Name of Isotope: ____________
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Example
Atomic Number: _________
Mass Number: __________
# of Protons: _________
# of Electrons: _______
# of Neutrons: ________
Name of Isotope: ____________
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Example
Atomic Number: _________
Mass Number: __________
# of Protons: _________
# of Electrons: _______
# of Neutrons: ________
Isotope Name: _________
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Draw the isotope symbol for
Calcium with 21 neutrons.
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Comparing Potassium Isotopes
Potassium-41
• protons = ________
• electrons = ______
• neutrons = _______
Potassium-40
• protons = ________
• electrons = ______
• neutrons = _______
Potassium-39
• Protons = __________
• Electrons = _________
• Neutrons = __________
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Relative Atomic Masses
• Atomic masses measured in grams
are very small.
• Example: An atom of Oxygen-16
has a mass of 2.657 x 19-23grams.
• It is most convenient to measure
mass in relative atomic masses.
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To set up a relative scale of
atomic masses:
1. One atom is chosen and assigned a
relative mass value,
2. All the other masses are expressed in
relation to this defined standard.
• Carbon-12 is the chosen standard.
• A single atom of C-12 is assigned a mass
of exactly 12 atomic mass units(u).
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An AMU
• One amu is exactly 1/12th of the
mass of a carbon-12 atom
– 1.660 5402 x 10-24grams.
• The atomic mass of a carbon-12
atom is exactly 12 u.
• Atomic Mass: The mass of an atom
expressed in atomic mass units.
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Average Atomic Mass
• Average Atomic Mass
– given on the periodic table
– weighted average for the naturally occurring mixtures
of isotopes in each element
• Average Atomic Mass depends on
– mass and relative abundance of the isotopes.
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Calculating average atomic mass:
Step 1:
Isotope 1: Atomic mass x relative abundance = mass contribution 1
Isotope 2: Atomic mass x relative abundance = mass contribution 2
Step 2: Add the mass contributions results.
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Example Problem:
Naturally occurring copper consists of
69.17% Cu-63, mass of 62.939 598u and
30.83% Cu-65, mass of 64.927 793u
Cu-63
Cu-65
0.6917 x 62.939 598u = 43.535u
0.3083 x 64.927 793u = 20.017
63.552u
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Example Problem 2
Gallium occurs in nature as a mixture of two
isotopes. They are Ga-69 with a 60.108% abundance
and a mass of 68.926 amu and Ga-71 with a
39.892% abundance and an atomic mass of 70.925.
Calculate the atomic mass of gallium.
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Example 3
The element chlorine occurs in nature as a mixture of two
isotopes. Chlorine-35 has an atomic mass of 34.969 amu and
makes up 75.77% of chlorine atoms. Chlorine-37 atoms make
up the remaining 24.23% of all chlorine. Use the average
atomic mass of chlorine from the periodic table to calculate
the atomic mass of Cl-37 atoms.
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Example 4
The atomic mass of bromine given in the periodic
table is 79.904 amu, which is very close to 80 amu.
Use a reference book to find the percent of Br-80 in
naturally occurring bromine. Explain the value of the
atomic mass of bromine from the data you find.
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