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Unit 3, Part 1: Atomic
Structure & The Periodic Table
3.1 Early Models of the Atom
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What are Atoms?
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4th century B.C. Greek philosopher
Democritus stated the universe
was made of invisible units called
atoms (atom “unable to be divided”).
Believed movements in atoms
caused changes observed in
matter.
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Late 1700’s

French Chemist, Antoine
Lavoisier established the Law of
Conservation of Matter.
Matter can not be created or
destroyed… only change forms

Joseph Proust, later established
the Law of Constant Composition:
the compound always contains the same
elements in the same proportions by
mass.
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In 1803 John Dalton proposed
an atomic theory:
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Each element is composed of
extremely small particles called
atoms .
Atoms of the same element are
exactly alike .
Every compound always has the
same ratio and kinds of atoms .
A chemical rxn is a
rearrangement of atoms; they are
not created or destroyed .
There are some exceptions to Dalton’s theory;
however they still are the basis for understanding
Chemistry.
3.2 Discovering Atomic
Structure
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Dalton and contemporaries
thought the atom was like a
marble… small, hard and round…
but couldn’t explain why atoms
from other element behaved
differently.
The English Chemist, Michael
Faraday, in 1839, showed atoms
contain electrical charge.
Electricity

American, Benjamin Franklin,
experimented with electricity (think
kite and key experiment). Franklin was
able to determine the 2 charges an
object has and it was he who
coined the names “positive” and
“negative” charge. Opposite
charges attract and like charges
repel.
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But where do these charges come
from? What are they?
Cathode Rays and Electrons

Cathode Ray Tube (CRT) - A battery is
connected to a tubing of partially evacuated
glass. The glass is lined with fluorescent
material, current flows to the ends of the tube.
The end connected to the (-) terminal of the
battery is called the cathode and the other is the
anode (+). A stream of radiation flows from the
cathode to the anode when the battery is turned
on.
At the end of the 19th Century
we knew:
The cathode ray could spin a small
paddle wheel = actual a stream of
particles.
Magnet deflects ray as expected for a (-)
charge = ray made of (-) charged
particles.
HOW do we prove it’s made of particles?
Electrons
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English physicist, J.J. Thompson (1856-1940),
determined the ray WAS made (-) particles by
allowing the ray to pass through a hole in the
anode and then through a magnetic field. He
determined that the negative particles emanate
from the cathode, they had structure and he
named them electrons.
American physicist, Robert Millikan (18681953), was able to measure the charge of an
electron. He sprayed oil and used X-rays to
give the oil a negative charged. Then measured
how different magnetic charges changed the
rate the oil fell. He calculated the mass of the
e- to be 9.11 X 10-19 grams.
Radioactivity
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French physicist, Henri Becquerel (18521908) , accidentally placed uranium (U) on
unexposed film. The U had produced an
image suggesting it was emitting some
type of radiation (the spontaneous
emission of particles) and that elements
could be radioactive.
His associates, Marie Sklodowska Curie
and her husband Pierre, discovered two
more radioactive elements, radium and
polonium.
Radioactivity was key to understanding
the atom, as a whole.
3 Types of Radiation

New Zealander, Ernest Rutherford passed
the cathode ray of a radioactive substance
between two charged plates. The ray splitpart of the beam was deflected towards the (-)
plate (alpha radiation), another was deflected
towards the (+) plate (beta radiation) and a third
passed straight through undisturbed (gamma
radiation).
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Alpha particles 2+ charge
Beta particles 1- charge
Gamma rays have no charge
This experiment demonstrated that the atom was
much more complex than previously thought.
The Nuclear Atom

Thompson showed us that atoms had electrons, but that
doesn’t explain why atoms are electrically neutral (they
don’t have a charge). If they have electrons they must
have some type of (+) charges, too.

1909 The Alpha-scattering Experiment by Rutherford- A
beam of high-speed alpha particles bombarded a thin
sheet of Au foil. Most of the alpha particles went straight
through the foil but a small portion DID deflect. And they
would scatter in every direction possible. Why??
Atomic Model Battle!
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Thompson had suggested,
earlier, that the atom was
like plum pudding. The e-s
were spaced evenly
throughout the atoms (+)
interior, the way the plums
were distributed through
the pudding.
(You can also think of chocolate chip cookie
dough- the dough is the positive interior while
the chocolate chips are the e-s).
Atomic Model Battle! Cont.
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Rutherford’s data
suggested this was not true.
He determined that there
must be an (+) core in the
atom. He named this the
nucleus. The particles
which went straight through
suggested the atom was
mostly empty space.
3.3 Modern Atomic Theory
The Structure of the Atom
 We now know that atoms can be divided into many
different subatomic particles. For example:
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Nucleus- the center of the atom, very dense, contains
protons and neutrons, has an overall positive (+) charge.
Protons- positively (+) charged subatomic particle.
Neutrons- neutral (not charged) subatomic particle.
Electrons- negatively (-) charged subatomic particle that
has very little mass. Creates a “cloud” that encircles the
nucleus.
Rutherford thought of it as like a mini solar system:
the nucleus in the middle encircled by electrons.
Atomic Numbers

Englishman, Henry Mosley (18871915) discovered that each element
had a unique amount of positive
charge. This concept leads to the
understanding of why atoms of
different elements are unique. The
identity of the atom comes from
the number of protons contained in
the nucleus.
Atomic #
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Atomic # (Z)- the # of protons in the nucleus of an atom.
(Always a whole #)
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Now, individual atoms are electrically neutral… meaning they
have an equal # of protons as electrons.
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Ex. O-8 = 8 protons in the nucleus
C-6 = 6 protons in the nucleus
Pb-82= 82 protons in the nucleus
O has 8 protons (we know this by its atomic number) and 8
electrons.
C has 6 protons and 6 electrons
Pb has 82 protons and 82 electrons.
WHAT HAPPENS WHEN THEY ARE NOT EQUAL??
Ions
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The atom can gain or lose electrons. This is an ion.
(NOT PROTONS- if it change the # of protons you change the element/atom.)
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Charge on Ion = # of protons - # of electrons
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If you gain an electron you get a negative ion.
If you lose an electron you have a positive ion.
Magnesium (Mg) can loose 2 electrons making it a positive ionwritten as:Mg2+
O will gain 2 electrons making it a negative ion. It is written as
such: O2-
Atoms that lose electrons, like Mg, are called cations.
Atoms that gain electrons, like O, are called anions.
Atomic Mass #
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Atomic mass # (A) - the # of protons + neutrons in an atom.
Because the bulk of the atom’s mass is provided by the
protons and neutrons, we only consider their masses when
calculating the atomic mass #.
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Ex: O has 8 protons and 8 neutrons, so A= 16.
Dalton said that every atom of the same element is exactly
alike… NOT SO. They do have the same # of protons;
however, they do not necessarily have the same # of
neutrons!
Isotopes
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This means that different atoms of the same
element may/will have different masses.
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isotopes (atoms having the same # of protons but different # of
neutrons).
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Some isotopes are more common than others.
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3 isotopes of H (why A = 1.00794).
 Protium (only has one proton in the nucleus) A=1
 Deuterium (1 proton + 1 neutron) A=2
 Tritium (1 proton + 2 neutrons) A=3
Determining The # of Neutrons
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Calculating the # of neutrons in an
atom:
General Formula:
A – Z= # of neutrons.
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Ex: calculate the # of neutrons in the isotopes of Cl (Cl 35
and Cl 37)
The Mass of an Atom
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The mass of a single element is extremely small
(in the one trillionth of a billionth range) and is very difficult to
work with. So instead we express the mass of
atoms in atomic mass units (amu).
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One amu is equal to 1/12th of the mass of a carbon-12
atom. This isotope has exactly 6 protons and 6
neutrons so the mass of each has to be about 1.0 amu.
1amu = 1/12 (mass of 126C atom) = 1.66x10-24 g
Average Atomic Mass

The atomic mass of an element is often
listed as the average atomic mass as found
in nature. This is a weighted average of the
isotopes for that particular element. The
more commonly found isotopes have a
greater effect on the averages mass than
the more rare isotopes.
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Ex: Cl 24% Cl-37 and 76% Cl-35, thus the
average atomic mass (35.45 amu) is much
closer to 35 than 37.