Transcript Atomic Structure Timeline Song

The History of the Atom
Atomic Theory
• Because we can not see atoms, we use
models to teach and learn about atoms.
• The atomic theory has changed over
time as new technologies have become
available.
– Remember: Scientific knowledge builds on past
research and experimentation.
Ancient Greece
• Aristotle:
– There are four elements:
• Earth, air, water, fire
– The four elements combine in various ways
to make all matter.
– Matter is continuous. (There is not a
“smallest” particle.)
Ancient Greece
• Democritus:
– Matter is made of tiny particles called
atoms.
– The atom is the smallest piece. It is
indivisible. (It can’t be divided any
further.)
– Atoms of an element have specific
properties (smooth, spiky, etc.) that give
the element its properties.
John Dalton-
Father of the Atomic Theory
• Matter is made of indivisible particles called atoms.
• Atoms of the same element are identical, but differ
from atoms of other elements.
• Atoms cannot be created or destroyed.
• Atoms of different elements can combine in simple,
whole number ratios to form compounds. (The Law
of Definite Proportions)
• Atoms of same element can combine in more than
one ratio to form two or more compounds. (The Law
of Multiple Proportions)
• The atom is the smallest unit of matter that can take
part in a chemical reaction.
John Dalton
• The “Billiard Ball” model:
– An atom is a solid, indivisible sphere
J.J. Thompson
• The “cathode ray tube” experiment
• Energizing matter in the tube removes charged
particles and makes a “cathode ray” (beam of
light).
• Magnets can deflect the cathode ray in a way
that shows it is made of negative particles.
• The amount of deflection tells us how massive
the particles are.
J.J. Thompson
• Thomson discovered that:
– The atom is NOT “indivisible”
– A small negatively-charged particle (the
electron) can be removed from the atom.
– He concluded that:
• The atom must have negative particles
(electrons) in it (it’s not like a billiard ball).
• The electrons are evenly distributed amongst
positive background “stuff” that makes up most
of the atom.
J.J. Thompson
This led to the “Plum Pudding” model
(The electrons are
like the plums in a
pudding– or the
chips in a chocolate
chip cookie.)
Plum pudding– plum-flavored pudding with plums distributed
throughout it– used to be a popular thing people ate. We
still call this the Plum Pudding Model even though we don’t
eat plum pudding.
Ernest Rutherford
• The Gold Foil Experiment
• Alpha particles (which
are positive) were
shot at a piece of
gold foil.
• They were expected
to go straight
through.
• Some particles were
(unexpectedly)
deflected.
• What deflected
them??
• (Positive) alpha
particles are
deflected by other
positives.
• (Like charges repel.)
• The nucleus must be
positive!
• The nucleus must
have most of the
mass, or it would get
pushed around by
the alpha particles.
• The nucleus must be
ridiculously small
because 10,000
alpha particles pass
straight through for
each one that is
deflected.
Ernest Rutherford
• The Gold Foil Experiment: Conclusions
– Atoms have a NUCLEUS
• It contains all the positive charge
• It contains almost all the mass
• It is TINY (1/10000 of the atom volume)
– Everything else is mostly empty space
Ernest Rutherford
• The “nuclear model” separates the
atom into two parts:
– Nucleus: positive & massive & tiny
– Electrons: negative, tiny, practically
weightless.
The Nuclear Model:
• This picture is not to
scale!
• The nucleus should be
WAY smaller than the rest.
Neils Bohr
• Since opposites attract, electrons are
attracted to the (positive) nucleus.
• Something must prevent electrons from
“falling into” the nucleus.
• Electrons are located in orbits
around the nucleus.
• Like planets orbiting the sun, they
are attracted, but if they stay at
the correct distance, they don’t
fall in.
• The distance from the nucleus
determines the energy.
Neils Bohr
Electrons
Or
are allowed here!
to be here.
• Bohr “quantized” the
atom.
• Only certain energies
are allowed, which
means only certain
orbits are allowed.
• All other places are
“forbidden” to the
electrons.
But not here. (This is forbidden because
no orbit for the electron to “land on.”)
Niels Bohr
• Moving from one energy level to
another requires the electrons to absorb
or emit energy
• Because only
certain orbits
are allowed…
• The energy
comes in
specific colors
of light for
each element.
Albert Einstein
• The photoelectric effect:
– Electrons can be ejected from metal by
light (energy) of a certain frequency.
Photoelectric Effect
• If the light isn’t high enough energy,
nothing happens.
• If it does have enough energy, electrons
are removed.
Albert Einstein
• The photoelectric effect proved:
– Different colors of light are worth different
amounts of energy.
– Electrons can be moved to different orbits
(or out of the atom altogether) by energy
in the form of light.
– Light has momentum– therefore it is a
particle (even though it has no mass).
– Light particles are called photons.
Thomas Young
• The double-slit experiment
• Electrons were sent through two small
openings and collected on a screen on
the other side.
• The electrons created
interference patterns.
• Interference patterns come
from waves overlapping
constructively and
Therefore electrons
destructively.
are a wave (?!)
Louis de Broglie
• Particle-Wave duality
– Einstein proved light could act like a
particle
– Young proved electrons could act like a
wave
– De Broglie concluded that “wave” and
“particle” aren’t mutually exclusive. The
electron can act as either one.
– Electrons have particle-wave duality.
Louis de Broglie
– Electrons-as-waves explains
the fact that only certain
orbits are “allowed” for each
element.
• The orbit’s circumference has to be a multiple
of the wavelength in order for it to be
“allowed.”
Heisenberg:
• The Uncertainty Principle
– Electrons can be moved by light.
– We see things because light bounced off of
them.
– So every time you “see” an electron, the
light that made it visible to you probably
also caused it to move.
– So it isn’t actually there anymore!
– It is impossible to know the location and
velocity of an electron at the same time.
The Uncertainty Principle
– Means that electrons cannot be traveling in
predictable circular orbits around the
nucleus.
– They move randomly and unpredictably.
– We can estimate the probability of finding
an electron somewhere, but we cannot say
where it is with certainty.
– (Let’s face it: if de Broglie is correct and the
electron is partly a wave, it doesn’t necessarily
“travel” as a particle anyway.)
Erwin Schrödinger
• Developed an equation to solve for the
locations the electron probability is
greatest.
– This basically describes a cloud-like region
where the electron is most likely to be
found.
– It cannot say with any certainty where the
electron actually is at any point in time, but
it describes where the electron could be.
Erwin Schrödinger
• The probable locations of
the electron predicted
by Schrödinger's equation
happen to coincide with
the locations specified in
Bohr's model.
• The difference is that
now everything is fuzzy
because of the lack of
certainty.
Modern Atom: the Quantum model
• Combining Heisenberg & Schrodinger
(with all the previous discoveries) gives us:
• Orbitals, not orbits
• Cloud-like regions (fuzzy
borders) where the
electron probability is
highest.
• Orbitals are 3-D (orbits were 2-D).
• Quantized energies
• Electrons can only exist at specific energies
allowed by the orbitals.
• Electrons can absorb or emit energy to move
to higher or lower energy orbitals.
The current (modern) atomic model: