Transcript Chapter 5 – Early Atomic Theory & Structure 5.1 Early Thoughts 5.6 Subatomic Parts of the Atom 5.2 Dalton's Model of the Atom 5.7 The.

Chapter 5 – Early Atomic
Theory & Structure
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 General Arrangement of
Subatomic Particles
5.4 The Nature of Electric Charge 5.9 Atomic Numbers of the
Elements
5.5 Discovery of Ions
5.10 Isotopes of the Elements
5.11 Atomic Mass
1
Early Thoughts
• 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.
2
Early Thoughts
 Aristotle (384-322 B.C.) rejected the theory
of
Democritus
and
Empedoclean theory.
advanced
the
– Aristotle’s influence dominated the
thinking of scientists and philosophers
until the beginning of the 17th century.
3
Dalton’s Model of the Atom
2000 years after Aristotle, John Dalton an
English schoolmaster, proposed his
model of the atom–which was based on
experimentation.
4
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.
5
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.
6
Dalton’s atoms were individual particles.
Atoms of each element are alike in
mass and size.
7
5.1
Dalton’s atoms were individual particles.
Atoms of different elements are not alike
in mass and size.
8
5.1
H 2
=
O 1
H 1
=
O 1
Daltons atoms combine in specific ratios
to form compounds.
9
Composition of Compounds
The Law of Definite Composition
A compound always contains two
or more elements combined in a
definite proportion by mass.
10
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.
11
Composition of Compounds
The Law of Multiple Proportions
Atoms of two or more elements may
combine in different ratios to produce
more than one compound.
12
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.
13
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.
15
Composition of Compounds
16
The Nature of Electric Charge
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 less 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.
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Discovery of Ions
• 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).
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Discovery of Ions
• 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
+
Cl
19
Discovery of Ions
NaCl →
+
Na
+
Cl
 In the melt the positive Na+ ions moved
to the cathode (negative electrode).
Thus positive ions are called cations.
 In the melt the negative Cl- ions moved
to the anode (positive electrode). Thus
negative ions are called anions.
20
Subatomic Parts of the Atom
The diameter
an billionths
atom is 0.1ofto a
This
is 1 to 5often
0.5 nm.
meter.
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.
21
Subatomic Particle - Electron
• 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.
22
Subatomic Particle - Electron
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.
23
Subatomic Particle - Electron
This was the discovery of the
fundamental unit of charge
– the electron.
24
Subatomic Particle - Proton
• 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
indivisible.
held
that
atoms
were
25
Subatomic Particle - Neutron
• James Chadwick discovered the neutron
in 1932.
 Its actual mass is slightly greater than the
mass of a proton.
26
Subatomic Particles
27
Ions
• Positive ions were explained by
assuming that a neutral atom loses
electrons.
 Negative
ions were explained by
assuming that extra electrons can be
added to atoms.
28
When one or more electrons are lost
from an atom, a cation is formed.
5.4
29
When one or more electrons are added
to a neutral atom, an anion is formed.
5.4
30
The Nuclear Atom
• 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.
31
The Rutherford Experiment
• 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.
32
The Rutherford Experiment
5.5
Rutherford’s alpha particle scattering experiment.
33
The Rutherford Experiment
• 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.
34
The Rutherford Experiment
• 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.
35
The Rutherford Experiment
 Because alpha particles have relatively
high masses, the extent of the
deflections led Rutherford to conclude
that the nucleus was very heavy and
dense.
36
The Rutherford Experiment
Deflection
Scattering
37
Deflection and scattering of alpha particles by positive gold nuclei.
5.5
General Arrangement of
Subatomic Particles
• Rutherford’s experiment showed that an
atom had a dense, positively charged
nucleus.
 Chadwick’s work in 1932 demonstrated
the atom contains neutrons.
 Rutherford
also noted that light,
negatively charged electrons were
present in an atom and offset the positive
nuclear charge.
38
General Arrangement of Subatomic
Particles
• 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
39
General Arrangement of
Subatomic Particles
40
5.6
Atomic Numbers of the Elements
• 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.
41
Atomic Numbers of the Elements
Every atom with an atomic number of 1 is
a hydrogen atom.
Every hydrogen atom contains 1 proton in
its nucleus.
42
atomic
number
Every atom with an
atomic number of 1
is a hydrogen atom.
1 proton in the
nucleus
43
Atomic Numbers of the Elements
Every atom with an atomic
number of 6 is a carbon atom.
Every carbon atom contains 6
protons in its nucleus.
44
atomic
number
Every atom with an
atomic number of 6
is a carbon atom.
6 protons in the
nucleus
45
atomic
number
Every atom with an
atomic number of
92 is a uranium
atom.
92 protons
in the
nucleus
46
Atomic Numbers of the Elements
• Atoms of the same element can have
different masses.
 They always have the same number of
protons, but they can have different
numbers of neutrons in their nuclei.
 The difference in the number of neutrons
accounts for the difference in mass.
 These are isotopes of the same element.
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Atomic Numbers of the Elements
Isotopes of the Same Element
Have
Equal numbers of protons
Different numbers of
neutrons
48
Isotopic Notation
49
Isotopic Notation
6 protons + 6 neutrons
12
C
6
6 protons
50
Isotopic Notation
6 protons + 8 neutrons
14
C
6
6 protons
51
Isotopic Notation
8 protons + 8 neutrons
16
O
8
8 protons
52
Isotopic Notation
8 protons + 9 neutrons
17
O
8
8 protons
53
Isotopic Notation
8 protons + 10 neutrons
18
O
8
8 protons
54
Hydrogen has three isotopes
1 proton
1 proton
1 proton
0 neutrons
1 neutron
2 neutrons
55
Examples of Isotopes
Element Protons Electrons Neutrons Symbol
Hydrogen
Hydrogen
Hydrogen
Uranium
Uranium
Chlorine
Chlorine
1
1
1
92
92
17
17
1
1
1
92
92
17
17
0
1
2
143
146
18
20
1
1
H
2
1
H
3
1
H
235
192
236
192
U
35
17
37
17
U
Cl
Cl
56
Atomic Mass
• 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.
57
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
58
A typical reading from a mass spectrometer. The two
principal isotopes of copper are shown with the
abundance (%) given.
5.9
59
Atomic Mass
Using a mass spectrometer, the mass of one
hydrogen atom was determined to be 1.673
x 10-24 g.
To overcome this problem of such a
small mass relative atomic masses using
“atomic mass units” was devised to
express the masses of elements using
simple numbers.
60
Atomic Mass
The standard to which the masses of all other atoms
are compared to was chosen to be the most abundant
isotope of carbon – carbon 12.
A mass of exactly 12 atomic mass units (amu) was
assigned to carbon 12.
1
1 amu is defined as exactly equal to
the mass of
12
a carbon-12 atom
1 amu = 1.6606 x 10-24 g
12
6
C
61
Atomic Mass
H
K
U
Average atomic mass 1.00797 amu.
Average atomic mass 39.098 amu.
Average atomic mass 248.029 amu.
62
Average Relative Atomic Mass
• 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).
63
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)
(62.9298 amu) 0.6909 = 43.48 amu
(64.9278 amu) 0.3091 = 20.07 amu
63.55 amu
64
Concepts
Dalton’s Atomic Mode;
Law of Definite Composition
Law of Multiple Proportions
Three principle subatomic particles
Thomson Model of the Atom
Rutherford alpha-scattering experiment
Atomic Number, Mass number, number of
neutrons, number of Protons, Number of
Electrons
Three Isotopes of Hydrogen
Average Atomic Mass of an Element
66