The Development of Atomic Theory

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The Development of Atomic Theory
Larry Scheffler Lincoln High School Portland OR
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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

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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
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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
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Daltons Atomic Theory

• All elements are composed of indivisible and
  indestructible particles called atoms.
• Atoms of the same element are exactly alike, They
  have the same masses.

• Atoms of different elements have different
  masses.
• Atoms combine to form compounds in small whole
  number ratios..

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Objections to DaltonsAtomic 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

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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.
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The Electron

• The electron was the first subatomic particle to
  be identified.
• In 1897 J.J Thomson used a cathode ray tube to
  establish the presence of a charged particle
  known as the electron
• Thomson established the charge to mass ratio

E/m 1.76 x 108 coulombs/gram
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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.

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Thomsons Charge to Mass Ratio
E/m 1.76 x 108 coulombs/gram
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Thomsens 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
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Millikans Experiment

• By varying the charge on the plates, Millikan
  found that he could suspend the oil drops or make
  them levitate.

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Millikans Experiment
Millikan used his data to measure the charge of
an electron and then to calculate the mass of
the electron from Thomsons 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
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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
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Rutherfords Experiment
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. They found that although most of them
passed through. About 1 in 10,000 hit and were
deflected
1910
Ernest Rutherford
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Rutherfords Experiment
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Rutherfords Experiment
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Rutherfords Experiment

• By studying this pattern, Rutherford concluded
  that atoms have a very dense nucleus, but there
  are mostly empty space.

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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 Neutron
s
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Neutrons
Discovered by James Chadwick in 1932 Slightly
heavier than a proton Mass of a neutron 1.675
x 10-24 grams
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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.
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Aspects of the Bohr Model

• Bohr put together Balmers and Planks
  discoveries to form a new atomic model
• In Bohrs model
• Electrons can orbit only at certain allowed
  distances from the nucleus.
• Electrons that are further away from the nucleus
  have higher energy levels (explaining the faults
  with Rutherfords model).

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The Electromagnetic Spectrum
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Wave Characteristics
Energy of a wave E hn Frequency n number
of peaks per unit of time Speed of light c nl
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Emission Spectra

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Flame Tests
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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 high-energy orbit.
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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.

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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.

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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
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Modern View of the Atom
Albert Einstein demonstrated the equivalence of
matter and energy. Hence matter and energy in
Einsteins 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
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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
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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
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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.
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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.
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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.
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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.

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Wave Mechanical Model

• The location of the electron in a hydrogen atom
  is a probability distribution.

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Progression of Atomic Models
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• Our view of the atom has changed over time

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ATOMIC STRUCTURE
Particle
Charge
Mass
proton
charge
1
neutron
No charge
1
electron
- charge
0
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ATOMIC NUMBER AND MASS NUMBER
Mass Number
He
4
the number of protons and neutrons in an atom
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Atomic Number
the number of protons in an atom
Number of electrons Number of protons in a
neutral atom
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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
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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
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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.
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or Na-24
or Na-23
Isotope one of two or more atoms having the same
number of protons but different numbers of
neutrons
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Measuring Atomic Mass--the Mass Spectrometer

• The mass spectrometer can be used to determine
  the atomic mass of isotopes.

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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.

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Mass Spectrum of Germanium

• The mass spectrum of germanium shows 5 peaks at
  relative atomic masses of 70, 72,73,74, and 75

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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 Abundance 70
20.55 72 27.37
73 7.67 74 36.74
75 7.67
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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 Abundance 70
20.55 72 27.37
73 7.67 74 36.74
75 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
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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
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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

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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.

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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 Avogadros Number and it defines the
  size of a quantity we call a mole.

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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.

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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.

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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.

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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.

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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.