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Chapter 4
The Structure of the Atom
Section 4.1—Early Theories of
Matter
Science as we know it did not exist
several thousand years ago. Curiosity
sparked the investigations of most
scholarly thinkers and they based their
thoughts on life experiences.
 The prefix “phil” means loving or dear
& is seen in many science & nonscience words. “Philosophy” refers to a
love of knowledge.

Democritus
Section 4.1—Early Theories of
Matter

Democritus was a Greek philosopher and the
1st person to propose the idea that matter
was not infinitely divisible. He also thought
that:

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


Matter is composed of empty space through which atoms move
Atoms are solid, homogeneous, indestructible, & indivisible
Different kinds of atoms have different sizes & shapes
Apparent changes in matter result from changes in the atoms
themselves
The differing properties of matter are due to the size, shape, &
movement of the atoms
Section 4.1—Early Theories of
Matter

Democritus idea of the atomic “theory” was
rejected because he could not answer the
question of what held the atoms together.
Also, the most influential Greek philosopher
Aristotle rejected Democritus’ theory just
because it did not agree with his own ideas
of nature.
Section 4.1—Early Theories of
Matter

In 1808, another atomic theory was
published by John Dalton, a
schoolteacher in England.
Dalton’s theory marks
the beginning of the
development of modern
atomic theory.
Section 4.1—Early Theories of
Matter

Dalton revised Democritus’s ideas based
upon the results of scientific research he
conducted. Dalton thought that:

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All matter is composed of extremely small particles called atoms
All atoms of a given element are identical, having the same size,
mass, & chemical properties. Atoms of a specific element are
different from those of any other element
Atoms cannot be created, divided into smaller particles, or
destroyed
Different atoms combine in simple whole-number ratios to form
compounds
In a chemical reaction, atoms are separated, combined, or
rearranged
Section 4.1—Early Theories of
Matter
Dalton’s theory easily explained the law
of conservation of mass
 Dalton’s theory has received 2 revisions
over time

atoms are divisible into several subatomic
particles
 atoms of an element may have slightly
different masses

Section 4.1—Early Theories of
Matter
Atom—smallest particle of an element that
retains the properties of the element
 Scanning tunneling microscope allows us to
see atoms and to actually move atoms to
form shapes, patterns, & simple machines.
This has lead to the discovery of
nanotechnology—the atom-by-atom
building of machines the size of molecules.

Scanning tunneling microscope
Section 4.1—Early Theories of
Matter

Molecule—is a group of atoms that are
bonded together & act as a unit
Section 4.1—Early Theories of
Matter
**Draw a double bubble map of Dalton’s
& Democritus’s atomic theory.
Section 4.2—Subatomic Particles
& the Nuclear Atom
In the 1800s, scientists were looking for a
relationship between matter & electric
charge.
 Cathode ray tube—glass tube of which most
of the air and matter had been sealed out of
with electrodes connected to the ends of the
tube and then to a battery discovered by
William Crookes

negative terminal end—cathode
 positive terminal end—anode

Section 4.2—Subatomic Particles
& the Nuclear Atom

As a result of continued research, Crookes
was convinced that
cathode rays were actually a stream of charged
particles
 the particles carried a negative charge

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These negatively charged particles were
found to be in all matter and were called
electrons.
Section 4.2—Subatomic Particles
& the Nuclear Atom

As a result of William Crookes cathode ray tube,
one of the most important technological & social
developments was discovered—the television.
Section 4.2—Subatomic Particles
& the Nuclear Atom

J.J. Thomson, an English physicist, began to
run cathode ray experiments to measure the
ratio of the electron’s charge to its mass.
After obtaining the ratio, he compared it to
other known ratios, for instance the hydrogen
atom. Thomson found that the mass of the
electron was lighter than that of hydrogen,
which disproved part of Dalton’s atomic
theory (atoms can’t be divided into smaller
particles).
Section 4.2—Subatomic Particles
& the Nuclear Atom

In 1909, American physicist Robert Millikan
determined the actual charge of an electron,
-1. Knowing the mass to charge ratio and
the charge of an electron, Millikan was able
to determine the mass of a single electron—
9.1x10-28 g or 1/1840 mass of a hydrogen
atom.
Section 4.2—Subatomic Particles
& the Nuclear Atom

It was known that matter is neutral because
you don’t go around getting shocked when
you touch any object. So it was proposed
that electrons were evenly spaced within a
uniformly distributed positive charged. This
was known as the “plum pudding” model.
AKA chocolate-chip cookie dough model.
Section 4.2—Subatomic Particles
& the Nuclear Atom

In 1911, Ernest Rutherford conducted an
experiment which led to the rejection of the “plum
pudding” model. Using gold and an alpha particle
emitting source, Rutherford calculated that an
atom consisted mostly of empty space through
which the electrons move. He also coined the
term nucleus--the tiny, dense region in the center
of the atom that contained most of the atom’s
positive charge & virtually all of its mass.
Rutherford’s new theory was called the nuclear
atomic theory.
Gold foil experiment
Section 4.2—Subatomic Particles
& the Nuclear Atom
In 1919, Rutherford fine-tuned his concept of the
nucleus, stating that it contained positively
charged subatomic particles equal to but opposite
that of an electron (+1 charge)—proton.
 In 1932, Rutherford & his co-worker, James
Chadwick, concluded that the nucleus contained
yet another subatomic particle, a neutral particle
called the neutron, having a mass nearly equal to
that of a proton.

Section 4.2—Subatomic Particles
& the Nuclear Atom
The mass of the nucleus (protons &
neutrons) is about 99.7% of the atom’s
total mass.
**Draw a double bubble map about
Thomson’s “plum pudding” atomic
model and Rutherford’s nuclear atomic
model.

Section 4.3—How Atoms Differ
Not long after Rutherford’s gold foil
experiment, Henry Moseley discovered that
atoms of each element contain a unique
positive charge in their nuclei. The number
of protons in an atom is referred to as the
element’s atomic number.
 The periodic table is arranged left-to-right,
top-to-bottom by increasing atomic number.

Section 4.3—How Atoms Differ
Remember that all atoms are neutral—SO, if
Atomic number = number of protons =
number of electrons
 How many protons are in each type of
atom?
Gold (Au) _______
Silver (Ag) _________
Potassium (K) ________

example
A
B
C
element Atomic Protons electrons
#
Pb
82
8
30
Practice problems
1.
How many protons & electrons are in each
of the following atoms?
 a.
Boron
c. Platinum
 b. Radon
d. Magnesium

2. An atom of an element contains 66
electrons. What element is it?
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3. An atom of an element contains 14
protons. What element is it?
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**worksheet

Determine the element, symbol, # of p, n,
e, atomic #, or atomic mass where
needed.
Section 4.3—How Atoms Differ
Isotopes have the same number of protons &
electrons but different numbers of neutrons.
 Most elements in nature are a mixture of
isotopes. Isotopes with more neutrons have
higher mass numbers—sum of the number of
protons & the number of neutrons in the nucleus.
To make it easier to identify each of the various
isotopes of the elements, chemists add the mass
number after the name. Chemists also use
abbreviations to represent isotopes also.
Fig 4-15, p 100

Section 4.3—How Atoms Differ

Mass number – atomic number = number
of neutrons
107
47
Ar
109
47
Ar
Example 4-2 & practice b-f

Determine the number of protons,
electrons, & neutrons in the isotope of
neon. Also, name the isotope given.
element
Atomic #
Mass
number
a. neon
10
22
b. calcium
20
46
c. oxygen
8
17
d. iron
26
57
e. zinc
30
64
f. mercury
80
204
Section 4.3—How Atoms Differ

You can calculate the atomic mass of
any element if you know its number of
isotopes, their masses, & their %
abundances.
Example 4-3

Calculate the atomic mass of unknown
element X. Then identify the unknown
element, which is used medically to
treat some mental disorders.
Isotope
Mass
(amu)
Percent
abundance
6X
6.015
7.5%
7X
7.016
92.5%
Practice
1.
2.
3.
Boron has 2 naturally occurring isotopes: Boron-10 (abundance=
19.8%, mass= 10.013 amu), boron-11 (abundance= 80.2%, mass=
11.009 amu). Calculate the atomic mass of boron
Helium has 2 naturally occurring isotopes, helium-3 & helium-4.
The atomic mass of helium is 4.003 amu. Which isotope is more
abundant in nature?
Calculate the atomic mass of magnesium. The 3 magnesium
isotopes have atomic masses & relative abundances of 23.985 amu
(78.99%), 24.986 amu (10.00%), & 25.982 amu (11.01%).
Section 4.4—Unstable Nuclei &
Radioactive Decay
In a chemical reaction, only the electrons are
involved in the reaction—NOT the particles of the
nucleus. This is the reason that an atoms identity
does not change during a chemical reaction.
 However, there are some reactions that DO
involve the changing of an atoms nucleus—
nuclear reactions. In the late 1890s, scientists
noticed that some substances spontaneously
emitted radiation in a process called radioactivity.
The actual rays & particles emitted were called
radiation.

Section 4.4—Unstable Nuclei &
Radioactive Decay
Radioactive atoms emit
radiation because their
nuclei are unstable. To
gain stability, they emit
radiation so that they
can lose energy, trying
to gain stability—a
process called
radioactive decay.
 There are 3 types of
radiation:

Radiation
Type
Symbol
Mass
(amu)
Charge
Alpha
α
4
2+
Beta
β
1/1840
1-
Gamma
γ
0
0
Section 4.4—Unstable Nuclei &
Radioactive Decay
A nuclear equation shows the atomic number &
mass number of the particles involved. It is
important to note that in a nuclear equation, both
mass number & atomic number are conserved.
 The primary reason that atoms are not stable is
the neutron-to-proton ratio. Eventually,
radioactive atoms undergo enough radioactive
decay to form stable, non-reactive atoms—this
explains their rare existence in nature.
