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Atom and Ev
Atoms,
Energy,
and
the Heisenberg
Uncertainty Principle
By Lee Wignall
The idea that matter is made of small,
indivisible “pieces” has been around for
thousands of years.
The word “atom” is derived from the Greek
word “atomos”, meaning “that which cannot
be divided”.
Leucippus: Greek philosopher/scientist
Democritus: student of Leucippus and considered “Father of Modern Science”
Fast forward over 1000 years…
In 1808, John Dalton came up with four
postulates of his “atomic theory.”
(there are actually five, but only the first four are wellknown)
John Dalton
Fast forward 100 years…
In 1906, Thomson created a current of
negatively charged particles in a
cathode-ray tube that were much smaller
than an atom.
J.J. Thomson
He called these tiny particles
“corpuscles” and created the
new model of the atom.
Negatively charged
“corpuscles” were spread
around a positively charged
sphere, like raisins in plum
pudding.
Plum pudding atomic model
2 years later…
In 1908, Ernest Rutherford fired alpha particles
(protons) at gold atoms and discovered that the
center of the atom is densely packed and
positively charged.
Gold foil experiment
He called the center of the atom the
nucleus and proposed that the
negatively charged electrons orbit
around the nucleus like planets.
5 years later…
In 1913, Neils Bohr introduced quantum
mechanics into the atomic model.
He proposed the idea that electrons were only
allowed to orbit the nucleus at certain energy
levels. Instead of an infinite number of possible
orbits, only certain orbits were allowed.
(“quantizing the atom”)
Electrons could only jump to a higher
energy orbit if they absorbed a photon
with enough energy.
Electrons jumping down energy levels
would emit a photon with the exact
energy difference between orbits.
The Electron Volt (Ev)
-13.6 Ev
-3.4 Ev
The energy of subatomic particles is
measured in electron volts (Ev).
An electron absorbs a
photon with enough
energy to jump to the
next highest energy
level.
The Electron Volt (Ev)
An electron jumps down
to a lower energy level
and emits a photon with
the energy difference.
-13.6 Ev
-3.4 Ev
The lowest “natural” energy state is the
ground state.
The energy of subatomic particles is
measured in electron volts (Ev).
Glow in the Dark
Electrons in the paint absorb energy
from ambient light and jump to higher
energy orbits.
However, this energetic state is unstable
and as soon as the energy source (light) is
removed, the electrons jump back down to
their ground-state, emitting photons with
energy equal to the difference between
orbits.
Louis de Broglie
de Broglie proposed in 1924 that if light can act as
both a particle and a wave, then shouldn’t all
moving particles act as waves as well?
Matter waves: the more
massive an object, the smaller
the associated wavelength.
Therefore, large objects have
wavelengths that are WAY too
short to be noticeable on any
level.
13 years later…
In 1926, Erwin Schrodinger introduced
probabilty into the structure of the atom.
Borrowing de Broglie’s idea that all
matter exhibited wave-like properties,
Schrodinger created an equation that
describes the way a wave evolves.
The Schrodinger Wave Equation
?
Don’t worry, I’ll break it
down for you.
The wavefunction: “psi”
Called
the how
“del-squared
number”:
it’s
operator”,
the
square
this
root
ofover
minus
one.
Describes
Describes
allthe
thewavefunction
forces
acting
changes
on
thequantity
particle.
time.
This“imaginary
is “h-bar”,
Planck’s
constant
(6.6
*10^-34
J*s)
The
mass
ofhow
thepi.
particle
being described.
describes
the
wavefunction,
, changes from
divided
by 2 times
one place to another.
Applying the Schrodinger Equation to an Atom
When you apply the Schrodinger Equation to
an electron in an atom, you end up with a
wavefunction that gives the probability of
finding the electron in any given location.
All the possible locations “smear” together to
create an electron cloud.
This computer analysis shows the
results of the Schrodinger Equation
under certain circumstances. You can
clearly see the denser “cloud” where
the probability of finding the electron is
higher.
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?
?
nucleus
?
?
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?
?
?
?
Electron cloud
Without direct observation, the position of the electron is
given by the wavefunction--a probability wave that
describes the chances of finding it at any given location.
Observed location of electron
?
?
?
?
?
?
?
nucleus
?
?
?
?
?
?
?
Electron cloud
As
soon
as observation
we make andoes
observation
aboutbegin
the precise
Only
upon
the
electron
to so is
Without
But
electrons
direct
observation,
are
both
a
wave
the
position
and
a
particle,
of
the
electron
and
location
of
the electron,
wewith
reduce
the probability of
act
as
a
particle.
Just
like
the
double-slit
given
byitthe
wavefunction--a
probability
wave by
thatdirect
we should
be
able
to location
determine
it’s position
finding
in
any
other
to
zero,
thereby
experiment!
describes
observation
thethe
(with
chances
high-tech
of finding
equipment).
it at any given location.
“collapsing
wavefunction.”
The Uncertainty Principle
An interesting feature of the Schrodinger wave
equation is the more precise you calculate the
position of the particle, the less you can determine
the momentum (and vice versa).
Werner Heisenberg
This is called the Heisenberg Uncertainty Principle.
(formulated in 1927)
The
combined
uncertainty
in boththe
theless
position
So, the
more
you
know
about
position,
you
The
uncertainty
ininthe
particle’s
momentum.
The
uncertainty
the
particle’s
position.
momentum
has to be greater than (or
know and
about
momentum!
equal to) h-bar divided by two.
Click below to begin youtube video on
Heisenberg’s Uncertainty Principle.
http://www.youtube.com/watch?v=KT7xJ0tjB4A
I know it’s 1932 and I’m tardy to the party, but I
just discovered the neutron.
I’m kind of a big deal.
James Chadwick
Heisenberg and Bohr, two
heavyweights in atomic
theory for their quantum
contributions, immediately
adopted the theory and used
the neutron to explain small
discrepancies in their
experiments.