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Early Atomic Theory
and Structure
Chapter 5
Hein and Arena
Version 2.0
12th Edition
Eugene Passer
Chemistry Department
Bronx Community
1 College
© John Wiley and Sons, Inc
Chapter Outline
5.1 Early Thoughts
5.6 Subatomic Parts of the Atom
5.2 Dalton's Model of the Atom
5.7 The Nuclear Atom
5.3 Composition of Compounds
5.8 Isotopes of the Elements
5.4 The Nature of Electric Charge 5.9 Atomic Mass
5.5 Discovery of Ions
2
5.1
Early Thoughts
3
• The earliest models of the atom were
developed by the ancient Greek
philosophers.
– Empedocles stated that matter was made
of 4 elements: earth, air, fire, and water.
– Democritus (about 470-370 B.C.) thought
that all forms of matter were divisible into
tiny indivisible particles. He called them
“atoms” from the Greek “atomos”,
indivisible.
4
• Aristotle (384-322 B.C.) rejected the
theory of Democritus and advanced the
Empedoclean theory.
– Aristotle’s influence dominated the
thinking of scientists and philosophers
until the beginning of the 17th century.
5
5.2
Dalton’s Model
of the Atom
6
2000 years after Aristotle, John Dalton,
an English schoolmaster, proposed his
model of the atom–which was based on
experimentation.
7
Dalton’s Atomic Theory
1. Elements are composed of minute
indivisible particles called atoms.
2. Atoms of the same element are alike in
mass and size.
3. Atoms of different elements have
different masses and sizes.
Modern research has demonstrated that
Atoms
under special
circumstances
can
4.
Chemical
compounds
are
formed
by
atoms are composed of subatomic
be decomposed.
the union of two or atoms of different
particles.
elements.
8
Dalton’s Atomic Theory
5. Atoms combine to form compounds in
simple numerical ratios, such as one to
one , two to two, two to three, and so on.
6. Atoms of two elements may combine in
different ratios to form more than one
compound.
9
Dalton’s atoms were individual particles.
Atoms of each element are alike in
mass and size.
10
5.1
Dalton’s atoms were individual particles.
Atoms of different elements are not alike
in mass and size.
11
5.1
H 2
=
O 1
H 1
=
O 1
Dalton’s atoms combine in specific
ratios to form compounds.
12
5.3
Composition of
Compounds
13
The Law of Definite Composition
A compound always contains two or
more elements chemically combined
in a definite proportion by mass.
14
Composition of Water
• Water always contains the same two
elements: hydrogen and oxygen.
• The percent by mass of hydrogen in
water is 11.2%.
• The percent by mass of oxygen in
water is 88.8%.
• Water always has these percentages. If
the percentages were different, the
compound would not be water.
15
Composition of Hydrogen Peroxide
• Hydrogen peroxide always contains the same
two elements: hydrogen and oxygen.
• The percent by mass of hydrogen in hydrogen
peroxide is 5.9%.
• The percent by mass of oxygen in hydrogen
peroxide is 94.1%.
• Hydrogen peroxide always has these
percentages.
If the percentages were
different, the compound would not be
hydrogen peroxide.
16
The Law of Multiple Proportions
Atoms of two or more elements may
combine in different ratios to produce
more than one compound.
17
Combining Masses of Hydrogen and Oxygen
Mass
Hydrogen(g)
Mass
Oxygen(g)
Water
1.0
8.0
Hydrogen
Peroxide
1.0
16.0
Hydrogen
peroxide
has peroxide
twice as much
mass
of oxygen
in hydrogen
16g 2
=
=18 ¹
oxygen
water.
mass(by
of mass)
oxygenasindoes
water
8g 1
Combining Ratios of Hydrogen and Oxygen
• Hydrogen peroxide has twice as many
oxygens per hydrogen atom as does
water.
• The formula for water is H2O.
• The formula for hydrogen peroxide is
H2O2.
19
20
5.4
The Nature of
Electric Charge
21
Properties of Electric Charge
• Charge may be of two types: positive and negative.
• Unlike charges attract (positive attracts negative),
and like charges repel (negative repels negative and
positive repels positive).
• Charge may be transferred from one object to
another, by contact or induction.
• The smaller the distance between two charges, the
greater the force of attraction between unlike
charges (or repulsion between identical charges).
kq1q 2
F=
2
r
q1 and q2 are charges, r is the
distance between charges, and
k is a constant.
22
5.5
Discovery of Ions
23
• Michael Faraday discovered that
certain substances, when dissolved in
water, conducted an electric current.
• He found that atoms of some elements
moved to the cathode (negative
electrode) and some moved to the
anode (positive electrode).
• He concluded they were electrically
charged and called them ions (Greek
wanderer).
24
• Svante Arrhenius reasoned that an ion
is an atom (or a group of atoms)
carrying a positive or negative electric
charge.
• Arrhenius accounted for the electrical
conduction of molten sodium chloride
(NaCl) by proposing that melted NaCl
dissociated into the charged ions Na+
and Cl-.
Δ
NaCl → Na+ + Cl25
NaCl → Na+ + Cl• When melted, the positive Na+ ions moved
to the cathode (negative electrode). Thus
positive ions are called cations.
• When melted, the negative Cl- ions moved
to the anode (positive electrode). Thus
negative ions are called anions.
26
5.6
Subatomic Parts
of the Atom
27
An atom is very small
28
This
The diameter
is 1 to 5often
anbillionths
atom is 0.1oftoa
meter.
0.5 nm.
If the diameter of this dot is 1
Even smaller particles than atoms
mm, then 10 million hydrogen
exist. These are called subatomic
atoms would form a line across
particles.
the dot.
29
Subatomic Particles
30
Electron
31
• In 1875 Sir William Crookes invented
the Crookes tube.
• Crookes tubes experiments led the way
to an understanding of the subatomic
structure of the atom.
• Crookes tube emissions are called
cathode rays.
32
In 1897 Sir Joseph Thompson demonstrated
that cathode rays:
• travel in straight lines.
• are negative in charge.
• are deflected by electric and magnetic
fields.
• produce sharp shadows
• are capable of moving a small paddle
wheel.
33
This was the discovery of the
fundamental unit of charge
– the electron.
34
Proton
35
• Eugen Goldstein, a German physicist,
first observed protons in 1886:
• Thompson determined
characteristics.
the
proton’s
• Thompson showed that atoms contained
both positive and negative charges.
• This disproved the Dalton model of the
atom which held that atoms were
indivisible.
36
Neutron
37
• James Chadwick discovered the neutron
in 1932.
• Its actual mass is slightly greater than
the mass of a proton.
38
39
Ions
40
• Positive ions were explained by
assuming that a neutral atom loses
electrons.
• Negative ions were explained by
assuming that additional electrons can
be added to atoms.
41
When one or more electrons are lost
from an atom, a cation is formed.
5.4
42
When one or more electrons are added
to a neutral atom, an anion is formed.
5.4
43
5.7
The Nuclear Atom
44
• Radioactivity was discovered by Becquerel
in 1896.
• Radioactive elements spontaneously emit
alpha particles, beta particles and gamma
rays from their nuclei.
• By 1907 Rutherford found that alpha
particles emitted by certain radioactive
elements were helium nuclei.
45
The Rutherford Experiment
46
• Rutherford in 1911 performed experiments
that shot a stream of alpha particles at a
gold foil.
• Most of the alpha particles passed through
the foil with little or no deflection.
• He found that a few were deflected at large
angles and some alpha particles even
bounced back.
47
Rutherford’s alpha particle scattering experiment.
48
5.5
• An electron with a mass of 1/1837 amu
could not have deflected an alpha
particle with a mass of 4 amu.
• Rutherford knew that like charges
repel.
• Rutherford concluded that each gold
atom contained a positively charged
mass that occupied a tiny volume. He
called this mass the nucleus.
49
• If a positive alpha particle approached
close enough to the positive mass it
was deflected.
• Most of the alpha particles passed
through the gold foil.
This led
Rutherford to conclude that a gold
atom was mostly empty space.
50
• Because alpha particles have relatively
high masses, the extent of the
deflections led Rutherford to conclude
that the nucleus was very heavy and
dense.
51
Deflection
Scattering
Deflection and scattering of alpha particles by positive gold nuclei.
52
5.5
General Arrangement of
Subatomic Particles
53
• Rutherford’s experiment showed that an
atom had a dense, positively charged
nucleus.
• Chadwick’s work in 1932 demonstrated
that the atom contains neutrons.
• Rutherford also noted that light,
negatively charged electrons were
present in an atom and offset the positive
nuclear charge.
54
• Rutherford put forward a model of the
atom in which a dense, positively
charged nucleus is located at the
atom’s center.
• The negative electrons surround the
nucleus.
• The nucleus contains protons and
neutrons.
55
56
5.6
Atomic Numbers
of the Elements
57
• The atomic number of an element is
equal to the number of protons in the
nucleus of that element.
• The atomic number of an atom
determines which element the atom is.
58
Every atom with an atomic
number of 1 is a hydrogen atom.
Every hydrogen atom contains 1
proton in its nucleus.
59
atomic
number
Every atom with an
atomic number of 1
is a hydrogen atom.
1 proton in the
nucleus
60
Every atom with an atomic
number of 6 is a carbon atom.
Every carbon atom contains 6
protons in its nucleus.
61
atomic
number
Every atom with an
atomic number of 6
is a carbon atom.
6 protons in the
nucleus
62
atomic
number
Every atom with an
atomic number of
92 is a uranium
atom.
92 protons
in the
nucleus
63
5.8
Isotopes
of the Elements
64
• Atoms of the same element have the
same number of protons.
• Atoms of the same element can have
different masses, because they can have
different numbers of neutrons.
• These are isotopes of the same element.
65
Isotopes of the Same
Element Have
Equal numbers of protons
Different numbers of
neutrons
66
Isotopic Notation
67
Isotopic Notation
6 protons + 6 neutrons
12
C
6
6 protons
68
Isotopic Notation
6 protons + 8 neutrons
14
C
6
6 protons
69
Isotopic Notation
8 protons + 8 neutrons
16
O
8
8 protons
70
Isotopic Notation
8 protons + 9 neutrons
17
O
8
8 protons
71
Isotopic Notation
8 protons + 10 neutrons
18
O
8
8 protons
72
Hydrogen has three isotopes
1 proton
1 proton
1 proton
0 neutrons
1 neutron
2 neutrons
73
Examples of Isotopes
Element Protons
Electrons Neutrons Symbol
Hydrogen
Hydrogen
Hydrogen
1
1
1
1
1
1
0
1
2
Uranium
Uranium
92
92
92
92
143
146
Chlorine
Chlorine
17
17
17
17
18
20
11H
11
22
H
11 H
33
H
11 H
235
235
U
92 92
238
238
92 92 U
35
35
Cl
17 17
37
37
74
17 17Cl
5.9
Atomic Mass
75
• The mass of a single atom is too small to
measure on a balance.
• Using a mass spectrometer, the mass of
the hydrogen atom was determined.
76
A Modern Mass Spectrometer
Positive ions
formed from
sample.
Electrical field
From the intensity and positions
at slits
A mass
of the lines
on the
Deflection
of mass
accelerates
spectrogram
spectrogram,
theions different
positive
ions. positive
is recorded.
isotopes occurs
and their
relative
at
amounts can
be determined.
magnetic
field.
5.8
77
A typical reading from a mass spectrometer. The two
principal isotopes of copper are shown with the
78
abundance (%) given.
5.9
Using a mass spectrometer, the mass of one
hydrogen atom was determined to be 1.673
x 10-24 g.
79
This number is very small.
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
small
80
Numbers
of this
size
are too small
for of
To overcome
this
problem,
a system
practical
use.
relative atomic
masses using “atomic
mass units” was devised to express the
masses of elements using simple
numbers.
81
The standard to which the masses of all
other atoms are compared to was chosen
to be the most abundant isotope of
carbon.
12
6
C
82
A mass of exactly 12 atomic mass units
(amu) was assigned to
12
6
C
83
1
1 amu is defined as exactly equal to
12
the mass of a carbon-12 atom
1 amu = 1.6606 x 10-24 g
12
6
C
84
Average atomic mass 1.00797 amu.
H
85
Average atomic mass 39.098 amu.
K
86
Average atomic mass 248.029 amu.
U
87
Average Relative
Atomic Mass
88
• Most elements occur as mixtures of
isotopes.
• Isotopes of the same element have
different masses.
• The listed atomic mass of an element is
the average relative mass of the isotopes
of that element compared to the mass of
carbon-12 (exactly 12.0000…amu).
89
To calculate the atomic mass, multiply the
atomic mass of each isotope by its percent
abundance and add the results.
Isotope
Isotopic mass
(amu)
Abundance
(%)
63
29
Cu
62.9298
69.09
65
29
Cu
64.9278
30.91
Average
atomic mass
(amu)
63.55
(62.9998 amu) 0.6909 = 43.48 amu
(64.9278 amu) 0.3091 = 20.07 amu
63.55 amu
90
Relationship Between Mass
Number and Atomic Number
91
The mass number minus the atomic
number equals the number of neutrons in
the nucleus.
mass
number
atomic
number
109
47
Ag
atomic
mass number number
109
47
=
=
number of
neutrons
62
92
93