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Chapter 4
Atomic Structure
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Section 4.1 – Defining the Atom

Objectives

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Describe Democritus’s ideas about atoms
Explain Dalton’s atomic theory
Identify what instrument is used to
observe individual atoms
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Trust?

I am going to ask you to believe in
something that you can not see with
your unaided eye…

We can not see the tiny fundamental
particles that make up matter – yet all
matter is composed of such particles called
atoms
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Atom
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Atom
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The smallest particle of an element that retains
its identity in a chemical reaction
Think of a coin the size of a penny and
composed of pure copper (Cu)
Imagine grinding the coin into fine dust … each
speck still has the properties of Cu
 If you could continue to grind the dust into
smaller and smaller Cu dust particles you would
eventually come upon a particle of Cu that could
no longer be divided and still have the chemical
properties of Cu … this final particle is an atom

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Early Models of the Atom

Greek Philosopher Democritus (460 BC – 370
BC)
 Among the first to suggest the existence of
atoms

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He believed that atoms were
indivisible and indestructible
Although his ideas agree with
later scientific theory, they did
not explain chemical behavior
and lacked experimental support
– not based on the scientific
method – just philosophy
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Early Models of the Atom

Who was next?
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John Dalton (1766 – 1844)
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Experiment based – transformed Democritus’
ideas into scientific theory
Results of his experimentation and observations
are summarized in Dalton’s Atomic Theory
Dalton studied the ratios in which elements
combine in chemical reactions
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Early Models of the Atom
Dalton’s Atomic Theory

1.
2.
3.
4.
All elements are composed of tiny indivisible
particles called atoms
Atoms of the same element are identical. The
atoms of any other elements are different from
those of any other element
Atoms of different elements can physically mix
together or can chemically combine in simple
whole-number ratios to form compounds
In chemical reactions, atoms are combined,
separated or rearranged – but never changed
into atoms of another element
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Sizing up the atom
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Remember our copper (Cu) penny…We
ground it into smaller and smaller particles
until we came upon a Cu atom
How big (or small) would that Cu atom be?
Here’s a little perspective…

a pure Cu coin the size of a penny contains
2.4 x 1022 atoms!
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Sizing up the atom

Earths population is only 6 x 109 people!
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If you could line up 1 x 108 Cu atoms side by
side, they would produce a line only 1 cm long!
The radii of most atoms is
5 x 10-11 to 2 x 10-10 m
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Sizing up the atom
 Despite
their small size individual atoms are
observable with instruments such as scanning
tunneling microscope
 The image below shows an image of iron atoms
generated by a scanning tunneling microscope
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Section 4.2
– Structure of the Nuclear Atom

Objectives

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Identify three types of subatomic particles
Describe the structure of atoms according
to the Rutherford atomic model
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Subatomic Particles

One important change to Dalton’s
Atomic Theory…
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Atoms are known to be divisible
Atoms can be broken down into even
smaller, more fundamental particles called
subatomic particles

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Three kinds of subatomic particles are
electrons, protons and neutrons
How was this discovered?
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Subatomic Particles
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Electrons
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Discovered in 1897 by J.J. Thomson
Electrons are negatively charged subatomic
particles
Thomson performed experiments using a
cathode ray tube to deduce the presence
of negatively charged particles
Cathode ray tubes pass electricity through
a gas that is contained at a very low
pressure
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Thomson’s Experiment
Voltage source
-
+
Metal Disks
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Thomson’s Experiment
Voltage source

+
Passing an electric current makes a
beam appear to move from the
negative to the positive end
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Thomson’s Experiment
Voltage source
+

By adding an electric field he found
that the moving pieces were negative
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Thomson’s Experiment
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During experiment he used many different
metals
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By the amount the beam bent he could find the
ratio of charge to mass
This ratio was the same for every material
Thomson concluded it must be the same type of
piece in every kind of atom…the electron
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Mass of the Electron

In 1961 Robert Millikan determined the
mass of the electron to be 1/1840 the mass
of a hydrogen atom
 The mass of an electron is 9.11 x 10-28 g
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Conclusions from the study of the Electron
1.
Cathode rays have identical properties regardless
of the element used to produce them. All elements
must contain identically charged electrons.
2.
Atoms are neutral, so there must be positive
particles in the atom to balance the negative charge
of the electrons
3. Electrons have so little mass that atoms must contain
other particles that account for most of the mass
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Conclusions from the study of the Electron
4. Eugen Goldstein in 1886 observed what is now
called the “proton” - particles with a positive
charge, and a relative mass of 1 (or 1840 times
that of an electron)
5. 1932 – James Chadwick confirmed the existence
of the “neutron” – a particle with no charge, but a
mass nearly equal to a proton
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Subatomic particles
Name
Relative
Symbol Charge mass
Actual
mass (g)
Electron
e-
-1
1/1840 9.11 x 10-28
Proton
p+
+1
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
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The Atomic Nucleus
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Once subatomic particles were discovered
scientists wondered how they were put
together in as atom
We will look at two models:
Thomson’s Atomic Model (Plum Pudding Model)
 Rutherford’s Atomic Model
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The Atomic Nucleus
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He believed the
electrons were
struck into a lump of
positive charge
Thomson’s model turned
out to be short-lived
due to groundbreaking
work of Ernest
Rutherford a former
student of Thomson
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The Atomic Nucleus
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Ernest Rutherford’s Gold Foil Experiment
 Ernest Rutherford decided to test the
plum pudding model
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Wanted to determine how large the atoms were
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The Atomic Nucleus
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Used Alpha particles
 Helium atoms that have lost their two
electrons leaving a double positive
charge
Helium atom
Alpha particle
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The Atomic Nucleus
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Ernest Rutherford’s Gold Foil Experiment
 The alpha particles were fired at a thin
sheet of gold surrounded by a detecting
screen
 Particles that are hit on the detecting
screen are recorded
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The Atomic Nucleus
What Rutherford expected…
Because evenly distributed + charges of the plum pudding model
were not enough to stop the ++ of the alpha particles to pass
through without
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What he found …
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The Atomic Nucleus

What Rutherford’s results meant:
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The atom is mostly empty space.
The electrons are distributed around the nucleus,
and occupy most of the volume
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All the positive charge, and almost all the mass is
concentrated in a small area in the center.
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This is why most alpha particles passed through without
deflection (no + charge to deflect the ++ alpha particle)
This is why some alpha particles were deflected. The +
charge and mass of the center was enough to deflect the
++alpha particle.
The center is composed of protons and neutrons.
He called this a “nucleus”
His model was called the nuclear atom
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The Atomic Nucleus
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Section 4.3
– Distinguishing Among Atoms
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Objectives
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Explain what makes elements and isotopes
different from each other
Calculate the number of neutrons in an
atom
Calculate the atomic mass of an element
Explain why chemists use the periodic table
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Atomic Number
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All atoms are composed of protons, neutrons and
electrons
How then are atoms of one element different than
atoms of another element? Why is an atom of C
different from an atom of H?
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Elements are different because they contain different
numbers of protons
The atomic number of an element is the number of
protons in the nucleus of an atom of that element
# of protons in an atom = the number of electrons
 Atoms are electrically neutral
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Atomic Number
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Atomic number (Z) of an element is the number of
protons in the nucleus of each atom of that element
Element
# of Protons Atomic Number
Carbon
6
6
Phosphorus
15
15
Gold
79
79
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Mass Number
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The majority of mass of an atom is concentrated in the
nucleus
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The mass of an atom depends on the number of
protons and neutrons
Total number of protons and neutrons in an atom is
called the mass number
The number of neutrons in an atom is the difference
between the mass number and atomic number
Number of neutrons = mass number – atomic number
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Complete Symbols
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We began looking at symbols for some of
the elements, however:
Complete symbols contain the symbol of the
element, the mass number and the atomic
number
Mass
Superscript → Number
Atomic
Subscript →
Number
X
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Complete Symbols
Au
197
79
Find each of the following:
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Number of protons
79
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Number of neutrons
118
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Number of electrons
79
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Atomic number
79
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Mass number
197
Atoms can also be referred to by using the mass and the
name of the element – gold – 197
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Complete Symbols
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If an element has an atomic number of 34
and a mass number of 78, what is the:
Number of protons?
 Number of neutrons?
 Number of electrons?
 Name of element?
 Complete symbol?

34
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34
Selenium
78
Se
34
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Complete Symbols
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If an atom has 91 protons and 140 neutrons
what is the:
Atomic number?
 Mass number?
 Number of electrons?
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Complete symbol?
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Name of the element?
91
91+140 = 231
91
231
Pa
91
Protactinium
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Isotopes
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So, if two atoms have different number of
protons, they are different elements.
What if two atoms have different number
of neutrons?
Atoms with different number of neutrons are
isotopes of the same element
 They have different mass numbers
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Despite these differences, isotopes are
chemically alike because they have identical
numbers of protons and electrons

Protons and electrons are responsible for
chemical behavior
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Isotopes

Another way to say it: Isotopes are atoms of the same
element having different masses, due to varying
numbers of neutrons
 There are three known isotopes of hydrogen
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
Hydrogen-3
(tritium)
1
1
1
1
2
Nucleus
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Isotopes
•
Elements occur in
nature as mixtures
of isotopes
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Atomic Mass

How heavy is an atom of oxygen?
It depends, because there are different kinds
of oxygen atoms (different isotopes of Oxygen)
 We are more concerned with the average
atomic mass.
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This is based on the abundance
(percentage) of each isotope of that
element in nature.
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Determining Atomic Mass

We don’t use grams for this mass because the
numbers would be too small
Instead of grams, the unit we use is the Atomic
Mass Unit (amu)
It is defined as one-twelfth the mass of a carbon-12
atom
 Carbon-12 was chosen because of its isotope
purity.
Each isotope has its own atomic mass, thus we
determine the average from percent abundance
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Calculating Atomic Mass
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To calculate the average atomic mass of an
element we need to know 3 things:
What are the stable isotopes of the element
 The mass of each isotope of the element
 The natural percent abundance of each isotope of
the element
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Calculating Atomic Mass
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Sample Steps for An Element with 3 Naturally Occurring
Isotopes
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massisotope1 x natural abundanceisotope1 = avg. massisotope1
massisotope2 x natural abundanceisotope2 = avg. massisotope2
massisotope3 x natural abundanceisotope3 = avg. massisotope3
avg. massisotope1 + avg. massisotope2 + avg. massisotope3 = avg mass of
atom for given
element
The resulting sum is the weighted average mass of the
atoms of the element that occur in nature
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Atomic Mass
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Practice Problem
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The element copper has naturally occurring
isotopes with mass numbers of 63 and 65. The
relative abundance and atomic masses are 69.2%
for mass = 62.93 amu, and 30.8% for mass =
64.93 amu. Calculate the average atomic mass
of copper
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Periodic Table – A Preview

The periodic table is an arrangement of
elements in which the elements are
separated into groups based on a set of
repeating properties

The periodic table allows you to easily compare
the properties of one element to another
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Periodic Table – A Preview
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Each element is identified with its symbol
placed in a square
The elements are listed in order of
increasing atomic number from left to right
and top to bottom
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Periodic Table – A Preview
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Each horizontal row (there are 7 of them) is called
a period
 The properties of the elements vary as you
move across it from element to element
 The pattern then repeats as you move to the
next period
Each vertical column is called a group or family
 Elements in a group have similar chemical and
physical properties
 Identified with a number and either an “A” or a
“B”
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