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The Development of
Atomic Theory
Larry Scheffler
Lincoln High School
Portland OR
1
The Atom
• The term atom is derived from the
Greek word atomos (atomos)
meaning invisible
• Democritius (470-370 BC )
suggested that all matter was made
up of invisible particles called atoms
2
Law of Constant
Composition
A compound always contains atoms of two or
More elements combined in definite proportions
by mass
Example:
Water H2O always contains 8
grams of oxygen to 1 gram of
hydrogen
3
Law of Multiple
Proportions
Atoms of two or more elements may
combine in different ratios to produce
more than one compound.
Examples:
NO
NO2
N2O
N2O5
4
Dalton’s Atomic Theory
1. All elements are
composed of indivisible
and indestructible
particles called atoms.
2. Atoms of the same
element are exactly
alike, They have the
3.
Atoms
of
different
same masses.
elements have different
masses.
4. Atoms combine to form
compounds in small
whole number ratios.. 5
Objections to Dalton’s
Atomic Theory
• Atoms are not indivisible. They are
composed of subatomic particles.
• Not all atoms of a particular element
have exactly the same mass.
• Some nuclear transformations
alter (destroy) atoms
6
Crookes Experiment
Crookes found that passing an electrical current
through a gas at very low pressure caused the gas to
glow. Putting a magnet next to the beam caused it to
be deflected.
7
The Electron
1. The electron was the first subatomic
particle to be identified.
2. In 1897 J.J Thomson used a cathode ray
tube to establish the presence of a
charged particle known as the electron
3. Thomson established the charge to
mass ratio
E/m = 1.76 x 108 coulombs/gram
8
A Cathode Ray Tube
Thomson found that an electrical field would
also deflect an electron beam. He surmised
that the ratio of charge to mass is constant.
Thomson’s Charge to Mass
Ratio
E/m = 1.76 x 108 coulombs/gram
Thomsen’s Plum Pudding
Model
Thompson proposed that
an atom was made up of
electrons scattered
unevenly through out an
elastic sphere. These
charges were surrounded
by a sea of positive
charge to balance the
electron's charge like
plums surrounded by
pudding.
This early model of the atom
was called The Plum
Pudding Model. A more
contemporary American
label might be the “chocolate
chip cookie” model
11
Millikan’s Experiment

By varying the charge on the plates, Millikan found
that he could suspend the oil drops or make them
levitate.
12
Millikan’s Experiment
Millikan used his data to
measure the charge of an
electron and then to
calculate the mass of the
electron from Thomson’s
charge to mass ratio.
Given the charge =
1.60 x 10-19 coulomb and
the ratio of E/m = 1.76 x 108
coulombs/gram it is
possible to calculate the
mass
Mass
= 9.11 x 10-28 gram
13
Protons
First observed by E. Goldstein in 1896
J.J. Thomson established the presence
of positive charges.
The mass of the proton is
1.673 x 10-24 grams
14
Rutherford’s Experiment
1910
Rutherford oversaw Geiger and
Marsden carrying out his famous
experiment.
They fired high speed alpha
particles (Helium nuclei) at a
piece of gold foil which was only
a few atoms thick.
Ernest Rutherford
They found that although most of
them passed through. About 1 in
10,000 hit and were deflected
15
Rutherford’s Experiment
16
Rutherford’s Experiment
17
Rutherford’s Experiment
By studying this
pattern, Rutherford
concluded that
atoms have a very
dense nucleus, but
there are mostly
empty space.
18
Subatomic Particles
The diameter of a single atom ranges
From 0.1 to 0.5 nm. (1 nm = 10-9 m).
Within the atom are smaller particles:
Electrons
Protons
Neutrons
19
Neutrons
Discovered by James Chadwick in 1932
Slightly heavier than a proton
Mass of a neutron = 1.675 x 10-24 grams
20
The Bohr Model
Niels Bohr proposed the
Planetary Model in 1913.
Electrons move in
definite orbits around the
nucleus like planets
moving around the
nucleus. Bohr proposed
that each electron moves
in a specific energy level.
21
Aspects of the Bohr Model
Bohr put together Balmer’s and
Plank’s discoveries to form a new
atomic model
In Bohr’s model:
1. Electrons can orbit only at certain
allowed distances from the
nucleus.
2. Electrons that are further away
from the nucleus have higher
energy levels (explaining the
faults with Rutherford’s model).
22
The Electromagnetic Spectrum
Wave Characteristics
Energy of a
wave
E = hn
Frequency = n
= number of
peaks per unit
of time
Speed of light
c = nl
Emission Spectra
25
Flame Tests
According to Bohr
Atoms radiate energy
whenever an electron jumps
from a higher-energy orbit to a
lower-energy orbit. Also, an
atom absorbs energy when an
electron gets boosted from a
low-energy orbit to a highenergy orbit.
27
Problems with the Bohr
Model

The Bohr model provided a model that gave
precise results for simple atoms like hydrogen.

Using the Bohr model precise energies could
be calculated for energy level transitions in
hydrogen.

Unfortunately these calculations did not work
for atoms with more than 1 electron.
28
Weakness of the Bohr
Model
• According to the Bohr model electrons could
be found in orbitals with distinct energies.
• When the data for energies measured using
spectral methods where compared to the
values predicted by the Rydberg equation, they
were accurate only for hydrogen.
• By the 1920s, further experiments showed that
Bohr's model of the atom had some difficulties.
Bohr's atom seemed too simple to describe the
heavier elements.
29
Modern View of the Atom
The wave mechanical model for the
atom was developed to answer some of
the objections that were raised about
the Bohr model. It is based on the
work of a number of scientists and
evolved over a period of time
The quantum theorists such as Maxwell
Planck suggested that energy
consists of small particles known as
photons. These photons can have only
discreet energies
Maxwell Planck
30
Modern View of the Atom
Albert Einstein demonstrated the equivalence
of matter and energy. Hence matter and energy in
Einstein’s theory were not different entities but
different expressions of the same thing
Einstein then proposed the
equivalence of Matter and
Energy given by his famous
equation
E = mc2
31
Modern View of the Atom
Louis de Broglie suggested
that if energy could be
thought of as having particle
properties, perhaps matter
could be thought of as having
wave like characteristics
Louis de Broglie
32
Modern View of the Atom
Louis de Broglie proposed
that an electron is not just a
particle but it also has wave
characteristics.
E = mc2 = hn
33
Modern View of the Atom
The more precisely the position is determined, the less precisely
the momentum is known in this instant, and vice versa.
--Heisenberg, Uncertainty paper, 1927
Heisenberg proposed that it was
impossible to know the location
and the momentum of a high
speed particle such as an
electron.
34
Modern View of the Atom
The more precisely the position is determined, the
less precisely the momentum is known in this
instant, and vice versa.
--Werner Heisenberg,
Uncertainty paper, 1927
The atom cannot be defined as a
solar system with discreet orbits
for the electrons. The best that
we could do was define the
probability of finding an electron
in a particular location.
35
Modern View of the Atom
Edwin Schroedinger proposed
that the electron is really a
wave. It only exists when we
identify its location. Therefore
the electrons are best thought
of probability distributions
rather than discreet particles.
36
Modern View of the Atom
The modern view of the atom suggests
that the atom is more like a cloud.
Atomic orbitals around the nucleus
define the places where electrons are
most likely to be found.
37
Wave Mechanical Model
The location of the
electron in a hydrogen
atom is a probability
distribution.
38
Progression of Atomic Models
Our view of the atom has changed over time
39
ATOMIC STRUCTURE
Particle
Charge
Mass
proton
+ charge
1
neutron
No charge
1
electron
- charge
0
40
ATOMIC NUMBER AND MASS NUMBER
4
2
He
Mass Number
the number of protons and
neutrons in an atom
Atomic Number
the number of protons in an
atom
Number of electrons = Number of protons
in a neutral atom
41
Atomic Mass
The atomic mass of an atom is a relative
number that is used to compare the
mass of atoms.
An atomic mass unit is defined as 1/12
of the mass of an atom of carbon 12.
The atomic masses of all other atoms
are a ratio to carbon 12
42
Isotopes
Many elements have atoms that have multiple
forms
Different forms of the same element having
different numbers of neutrons are called isotopes.
For example: Carbon exists as both Carbon 12
and Carbon 14
Carbon 12
Carbon 14
6 electrons
6 electrons
6 protons
6 protons
6 neutrons
8 neutrons
43
Isotopes and Atomic Mass
Many elements have atoms that have
multiple isotopes.
Isotopes vary in abundance. Some are
quite common while others are very rare.
The atomic mass that appears in the
periodic table is a weighted average taking
into account the relative abundance of each
isotope.
44
or Na-23
or Na-24
Isotope: one of two or more atoms having the
same number of protons but different
numbers of neutrons
Measuring Atomic Mass
--the Mass Spectrometer
The mass
spectrometer
can be used
to determine
the atomic
mass of
isotopes.
Mass Spectrum of Neon

The mass spectrum neon shows three isotopes with the
isotope at atomic mass = 20 accounting for more than
90% of neon.
Mass Spectrum of Germanium

The mass spectrum of germanium shows 5 peaks at
relative atomic masses of 70, 72,73,74, and 75
Calculating the average relative
atomic mass

The average atomic mass that is shown in the
periodic table is really the weighted average of
the atomic masses of each of the elements
isotopes. Germanium has 5 isotopes whose
relative atomic masses are shown in the table
Mass Number
70
72
73
74
75
% Abundance
20.55
27.37
7.67
36.74
7.67
Calculating the Average
Relative Atomic Mass

To calculate the average atomic mass multiply the
atomic mass of each isotope by its abundance
(expressed as a decimal fraction)
Mass Number
70
72
73
74
75
% Abundance
20.55
27.37
7.67
36.74
7.67
Average atomic mass
= (0.2055)(70) + (0.2737)(72) + (0.0767)(73) + (0.3674)(74)+ (0.0767)(75)
= 72.36
Note: atomic masses are ratios so they do not have real units
although they are sometimes called atomic mass units or amu
Problem

The mass spectrum of an element, A, contained
4 lines at mass/charge ratios of 54, 56, 57 and
58 with relative intensities of 5.84, 91.68, 2.17
and 0.31 respectively. Calculate the relative
atomic mass of element A.
Average atomic mass
= (0.0584)(54) + (0.9168)(56) + (0.0217)(57) + (0.0031)(58)
= 56.02
The Nucleus

The nucleus is very small — of the order
of 10-14 meter whereas the atom is of the
order of 10-9 meters. By analogy, the
nucleus occupies as much of the total
volume of the atom as a fly in a cathedral
Protons and Neutrons


Protons and neutrons have nearly equal masses, and
their combined number, the mass number, is
approximately equal to the atomic mass of an atom.
The combined mass of the electrons is very small in
comparison to the mass of the nucleus, since protons
and neutrons weigh roughly 2000 times more than
electrons.
Charge
Mass
Electron
-1
9.109383 x 10-28 g
Relative
Mass
1/1837
Proton
+1
1.6726217 x 10-24 g
1
0
1.6749273 x 10-24 g
1
Neutron
Atomic Mass Units




An atomic mass unit (amu) is equal to exactly 1/12 of
the mass of an atom of Carbon 12.
One atomic mass unit is equal to 1.66054 x 10-24 grams.
Note that this is slightly less than the mass of a proton
or a neutron.
An atomic mass unit is sometimes called a Dalton (D).
1.00 g = 6.02214 x 1023 amu. This number is also known
as Avogadro’s Number and it defines the size of a
quantity we call a mole.
Radioactive Nuclei



The presence of neutrons in the nucleus tends
to buffer the repulsions of multiple protons in the
nucleus.
There appears to be an optimal number of
neutrons for the number of protons in a given
atom in a stable atom.
In a radioactive element the nucleus may
disintegrate releasing either an alpha particle or
a beta particle as well as some high energy
gamma radiation.
Alpha Particles



An alpha particle consists of two protons and two neutrons. This
makes it equivalent to a helium nucleus.
When a radioactive element undergoes alpha decay its nucleus is
decreased in mass by 2 protons and 2 neutrons.
Since the number of protons changes, it has a new atomic number
and hence it is a different element. The mass number decreases
by 4.
238
4
234
U
92
146 neutrons
92 protons
Ratio = 1.52
-->
Th +
90
144 neutrons
90 protons
Ratio = 1.60
He
2
Alpha decay raises the neutron
to proton ratio. It occurs in
radioactive nuclei where the ratio
is too low
Beta Particles


A beta particle consists of a high-speed electron
released from the nucleus.
When a radioactive element undergoes beta decay,
the number of protons increases by one and the
number of neutrons decreases by one. The mass
number remains the same.
14
C
-->
14
6
8 neutrons
6 protons
Ratio = 1.333
7
7 neutrons
7 protons
Ratio = 1.00
N +
0
e-
-1
Beta decay lowers the
neutron to proton ratio. It
occurs in radioactive nuclei
where the ratio is too high
Beta Particles


A beta particle consists of a high-speed electron
released from the nucleus.
When a radioactive element undergoes beta decay,
the number of protons increases by one and the
number of neutrons decreases by one. The mass
number remains the same.
14
C
-->
14
6
8 neutrons
6 protons
Ratio = 1.333
7
7 neutrons
7 protons
Ratio = 1.00
N +
0
e-
-1
Beta decay lowers the
neutron to proton ratio. It
occurs in radioactive nuclei
where the ratio is too high
The Half-Life
The rates at which various radioactive elements undergo
decay vary considerably. The half-life of a radioactive
element if the time required for half of the nuclei in a
sample of radioactive nuclei to disintegrate.
Isotope
Type
Half-life
Uranium 238
Carbon 14
Iodine 131
Radon-222
Alpha
Beta
Beta
Alpha
4.51 x 109
5730 years
8 days
3.825 days
Cesium -137
Polonium 210
Beta
Alpha
30 years
138 days