Atomos: Not to Be Cut
The History of Atomic Theory
Atomic Models
• At right is the image most people
have of what an atom looks like. It
shows electrons in various energy
levels orbiting a nucleus.
• Models use familiar ideas to explain
unfamiliar facts observed in nature.
Models can be changed as new
information is collected.
The Six Atomic Models are:
• So, is this the most accurate model
of the atom? Or is it far from reality?
1. The Greek Model
• This presentation discusses six
major models of what we call
“atoms”. Each new model was
developed based on new discoveries
at the time.
3. Thomson’s Model
2. Dalton’s Model
4. Rutherford’s Model
5. Bohr’s Model
6. Wave-Mechanical Model
1. The Greek Model
• Democritus was a Greek
philosopher who began the
search for a description of
matter more than 2400 years
ago (around 400 BC).
• He asked: could matter be
divided into smaller and
smaller pieces forever, or was
there a limit to the number of
times a piece of matter could
be divided?
It’s only
logical!
• Democritus knew this question could not
be easily answered directly. Let’s say you
were breaking up a piece of rock -- how
would you know when the smallest piece
of rock was actually reached?
• Democritus logically concluded that a
piece of matter could not be divided into
smaller pieces forever. Eventually the
smallest possible piece would be obtained.
• Since this piece is the smallest possible it
would be indivisible, meaning it could not
be further divided.
• He named the smallest possible piece of matter “atomos,”
which means “not to be cut.”
• To Democritus, atoms were small, hard particles that were
all made of the same material but were different shapes
and sizes. They were infinite in number, always moving,
and capable of joining together.
• However, Democritus’ theory was ignored and forgotten for
more than 2000 years! Why?
• The Greeks did not experiment, but tried to win arguments
through logic and debate.
• The famous philosopher Aristotle believed
differently, that matter was composed of
four elements: earth, fire, air and water.
These four elements were supposedly
blended in different proportions to make
up all the various types substances in the
world.
• Aristotle had such a great reputation as a
thinker that many of his ideas were simply
accepted as true without any experimental
evidence.
• Poor Democritus! He had the right idea
but was ignored. His idea of the nature of
atoms was not popular again until the time
of John Dalton in the19th century.
2: Dalton’s Model
• John Dalton proposed his “modern” atomic
theory in 1803, about 2000 years after
Democritus.
• Prior to Dalton matter was considered to
be continuous in nature, not composed of
indivisible bits.
John Dalton
• The Law of Multiple Proportions is the third
postulate of Dalton's atomic theory. It
states that the masses of one element will
combine with a fixed mass of a second
element in ratios of whole numbers.
1766-1824
See Sec 4.3 of
text.
Dalton’s Atomic Theory
• All elements are composed of atoms.
• Atoms of the same element are identical.
• Atoms of a given element are different
from those of another element.
• Compounds are formed by the joining of
atoms of two or more elements. A given
compound always has the same relative
number and types of atoms.
• Atoms are indivisible and may not be
created or destroyed in chemical
reactions. Chemical reactions simple
rearrange how atoms are grouped
together.
• Do the points of
Dalton’s theory
sound familiar?
• They should!
Dalton’s atomic
theory became one
of the foundations
of modern
chemistry.
3: Thomson’s Plum Pudding Model
• See section 4.5 in text.
• In 1897, the English
scientist J.J. Thomson
proposed that atoms
themselves are made of
even smaller, subatomic, particles.
• Thomson experimented with cathode ray tubes. When such a
tube is evacuated of air and an electric current is pass through
it, a glow (cathode rays) was observed.
• Cathode rays are so named because they are emitted at the
negative (cathode) end and travel to the anode (positive) end
of the tube.
• Cathode rays are deflected towards a positively charged plate
which showed they are composed of negative particles.
• Regardless of the type of metal used at the cathodes and
anodes of the vacuum tubes, the cathode rays were seen and
had the same properties.
• From these observations Thomson concluded that the
negatively charged particles came from within the atom.
• A particle smaller than an atom had to exist. Therefore,
the atom was divisible!
• Thomson called the negatively charged particles
“corpuscles”. We, of course, know them as today as
electrons.
•
Since atoms were neutral he reasoned that there must
also be positively charged particles in the atom, but he
could never identify them.
•
http://www.youtube.com/watch?v=IdTxGJjA4Jw&list=PLD6A43875B9DEC5
7E&index=24
• Thompson proposed a
model of the atom that
became known as the
“Plum Pudding” model.
• Atoms were made from a
positively charged
substance with negatively
charged electrons
scattered about, like
raisins in a pudding.
• The “Plum Pudding”
model was also proposed
around 1900 by William
Thompson (better known
as Lord Kelvin), who was
no relation to J. J.
Thompson.
4: Rutherford’s Model
• See sections 4.5, 10.1
in text
• In 1910, New Zealand
physicist Ernst
Rutherford performed a
series of experiments
to study the structure of
atoms.
• Rutherford’s experiment (see next slide) involved firing a
stream of positively (+) charged “bullets” – called alpha
particles (actually, helium nuclei) -- at a thin sheet of
gold foil about 2000 atoms thick.
• Most of the positively charged alpha particle “bullets”
passed nearly straight through the gold atoms in the
sheet of foil.
• However, some of the alpha particles bounced away
from the gold sheet as if they had hit something solid.
• How could something like this be explained given the
current model of what atoms were?
• Rutherford concluded that the gold atoms in the sheet
were mostly open space, and not like a solid “pudding.”
• He further concluded that both the mass and the positive
charge of an atom are concentrated within a tiny fraction
of the atom’s volume, which he called the nucleus.
• The few alpha particle “bullets” that were deflected were
bouncing off the tiny, dense nucleus at the center of the
mostly empty atom. Those not deflected were passing
through empty space!
What Rutherford expected: alpha particle paths virtually
uninterrupted as they went through the atom because at
the time it was believed that atoms had the same
consistence throughout.
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What he observed: scattering of some alpha particles as they
bounced off the dense nucleus at the center of the atom.
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• These observations certainly changed
the idea of how atoms could be
pictured.
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• Rutherford’s atomic model contained
positive charges in a very tiny, very
dense nucleus.
• The negatively charged electrons filled
the remaining much larger volume of
the atom and orbited outside the
nucleus.
• By 1920 Rutherford established that the nucleus
of an atom consisted of positively charged
particles, called protons, and, for stability, must
contain neutral particles with the same mass as
protons.
• Hydrogen is the exception with a single proton in
its nucleus.
• The neutral particle, called the neutron, was
eventually discovered in 1932 by James
Chadwick, a student of Rutherford’s.
• Another example of the Rutherford model of
the atom in which electrons can orbit the
atom in infinite number of paths.
5: Bohr Model
• See section 10.5 in text.
• Rutherford’s model, however, could
not explain how electrons could
maintain stable orbits around the
nucleus. They should, over time,
fall into the nucleus. But this is not
observed. So what was keeping
electrons in stable orbits – and
atoms in existence?
• In 1913, the Danish scientist Niels
Bohr proposed that electrons are
restricted to certain fixed
(quantized) orbits around the
nucleus.
• These orbits, or energy levels, are
only located at certain distances
from the nucleus.
• If electrons gain energy (become
energized) they can jump to a higher
fixed energy level.
• When electrons lose energy they
drop back down to a lower energy
level. The energy lost is released
(emitted) as a photon of light of a
specific frequency (color).
• A quantum (fixed amount) of energy is required
to move electrons to the next highest level.
Likewise, the same fixed amount of energy is
emitted when an electron drops to a lower
energy.
• Bohr’s model explained the observed emission
spectrum data for hydrogen (see fig 10.11, pg
286).
• In Bohr’s model, the position of electrons is
analogous to the steps of a stair (or rungs of a
ladder).
• Electrons cannot exist between energy levels,
just like you can’t stand between steps on a
stairs (or rungs on ladder).
6: The Wave Mechanical Model
• See sections 10.6, 10.7, 10.8 in text.
• Also called the quantum mechanical model, or cloud
model.
• By mid-1920s the Bohr model was shown to be incorrect
because it only worked well for the hydrogen atom.
• A new model suggested that electrons exhibit wave-like
as well as particle-like behavior (just as photons do).
• Electrons do not “orbit” the nucleus as in the Bohr model,
but are located within regions of space around the
nucleus.
• Scientists responsible for the model:
– Louis de Broglie was the first to support the idea
that electrons exhibit wave characteristics.
– Erwin Schrodinger developed a mathematical
wave formula that described electrons as waves.
– Werner Heisenberg showed that it is impossible
to know both the exact location and the exact
speed of an electron (the Uncertainty Principle).
• Summary of wave-mechanical model:
– The nucleus remains as defined by Rutherford.
– Electron states are described as orbitals, not orbits.
– Orbitals are regions around a nucleus where
electrons have a probability of being found. The
precise location and speed of electrons within
orbitals cannot be accurately determined (see firefly
example on pg 289).
– Schrodinger’s wave equations describe orbitals of
various shapes – such as spherical or dumbbell
(see pg 291 in text).
– Orbitals are also described as electron cloud regions.
– Electron clouds are visual models that map the
possible location of electrons in an atom.
Electrons can only be
present in certain
regions around the
nucleus, not just
anywhere.
not here
here
– The edge of an orbital or cloud is “fuzzy,” meaning it
does not have an exact size.
– Electron clouds are denser closer to the nucleus
where electrons are more likely to be found.
– Location of an electron depends upon how much
energy the electron has.
– Electrons closer to the nucleus have less energy than
those further away.
– Higher energy states of electrons correspond to
orbitals with different shapes.
Summary of Models
• Greeks
– Atoms are indivisible, infinite, moving, join together
• Dalton
– Atoms are indivisible, different for each element
– Form compounds by joining together in fixed ratios
• Thompson
– Atoms are divisible. Consist of negatively-charged electrons and
a particles with a positive charge.
– Plum-Pudding model
•
Rutherford
– Atoms consist of positively-charged particles called protons located in a
tiny, very dense nucleus. Neutral particles called neutrons are also located
within the nucleus for stability.
– Negatively-charged electrons are arranged in orbits around the nucleus.
– Atoms are mostly empty space.
•
Bohr
– Retains Rutherford’s basic model except that electrons are restricted to
certain orbits around nucleus. (Like planets around the sun.)
– Orbits of electrons are quantized energy levels.
– At best this model could only be applied to hydrogen.
•
Wave-Mechanical
– Electrons occupy orbitals which are regions of space around the nucleus
where they are likely to be found. They do not orbit the nucleus like
planets around the sun.
– This model worked for atoms of any element.
Some links to videos on the web
• Atomic timeline video:
http://www.youtube.com/watch?v=NSAgLvKOPLQ
• Lots of interesting videos! Pick the ones that interest
you.
http://www.youtube.com/watch?v=bw5TE5o7JtE&list=PLD6A43875B9DEC5
7E&index=1