Download Development of the Model of the Atom

Leading up to the Quantum Theory
DEVELOPMENT OF THE MODEL OF THE ATOM
WHAT WE THOUGHT WE KNEW ABOUT LIGHT



exhibits wavelike behavior
moves at a speed 3.8 × 108 m/s in a vacuum
there are measureable properties of light
(wavelength and frequency) and these
properties are related to each other
mathematically c = λν
THE PHOTOELECTRIC EFFECT

No electrons were emitted if the light’s
frequency was below a certain minimum,
regardless of the intensity of the light. If light is
thought of only in a wave context, then light of
any frequency should supply enough energy to
eject an electron. This was the conundrum
scientists were trying to explain.

Max Planck – suggested that energy comes in
small packets called quanta. He proposed the
relationship between a quantum of energy and
radiation to be E = hν
 E=the
energy of a quantum of radiation;
ν=frequency of the radiation emitted and;
h=constant (now a fundamental in physics called
Planck’s constant = 6.626×10-34 J•s)

Albert Einstein expanded on Planck’s Theory
and introduced an idea that all electromagnetic
radiation has a wave and particle nature to it.
Each particle of light carries a quantum of
energy. These particles are called photons.
Ephoton = hν

Thus Einstein explained the photoelectric effect
by proposing that radiation is absorbed by
matter in whole numbers of photons. In order
for an electron to be ejected, the electron must
be struck by a single photon possessing a
minimum amount of energy (which
corresponds to the frequency).
THE LINE EMISSION SPECTRUM OF HYDROGEN
Classical theory of energy predicted that
hydrogen atoms would be excited by whatever
amount of energy added to them – then you
would expect to see a continuous spectra of
frequencies.
What we really saw was four very specific
lines at 656 nm (red), 486 nm (green),
434 nm (blue) and 410 nm (purple).
Using the photon idea…These specific
frequencies indicated that the energy
emitted when an electron falls from an
excited state to its ground state or a lower
energy excited state. This also indicated that
an atom’s energy states were fixed.
BOHR MODEL OF THE ATOM



Proposed a model that linked the atom’s electron emission
to photon emission. Electrons can only circle the nucleus
only in allowed paths or orbits; the further the obit, the
higher the energy.
His model could explain emission spectrum observed for the
hydrogen atom. He calculated the allowed energy levels for
the hydrogen atom and related the possible energy level
changes to the lines in the emission line spectrum. His
calculated values agreed with experimentally observed
values.
What his model didn’t do: (1) Did not explain the spectra of
atoms with more than one electron. (2) Did not explain the
chemical behavior of atoms.
THE QUANTUM MODEL

de Broglie wondered that if light has a waveparticle duality, then maybe electrons may have
the same nature.
Scientists knew any wave confined to a space can have
only certain frequencies. De Broglie suggested that
electrons be considered waves confined to the space
around an atomic nucleus.
 Experiments demonstrated that electrons, like light
waves, can be bent or diffracted (bending as it passes
through a small opening). Also, some experiments
showed that electron beams interfere with each other.


http://youtu.be/O55XiriEaQI

Heisenburg worked on answering the question
of “where” electrons are in the atom, if they
(the electrons) are both wave and particle.
 Electrons
are detected by their interactions with
photons. Because photons have the same energy
as electrons, any attempt to locate them with a
photon knocks the electron off its course.
 The Heisenburg Uncertainty Principle states that it
is impossible to determining both the position and
velocity of an electron at the same time.

Schrödinger used the dual wave-particle nature
to create an equation that treated electrons as
waves. With this equation and Heisenburg’s
Uncertainty Principle laid the foundation for our
Modern Quantum Theory of the atom.
 Rather
than electrons travelling in distinct orbits
around the nucleus, electrons occupy threedimensional regions of high probability around the
nucleus.

http://www.youtube.com/watch?feature=player
_embedded&v=45KGS1Ro-sc