A
Brief Introduction to Quantum Chemistry - Part 2
Wave-Matter
Duality of Light
In the early
1900’s, the fundamental nature of light was still a mystery. Light
exhibited many properties that were consistent with both matter and
waves/energy. For example, when light enters a prism, it divides into
many colours, which are visual cues of electromagnetic radiation.
Electromagnetic radiation is a form of energy, like UV rays and
microwaves. However, light acts like matter when it is drawn into black
holes. Attractive gravitational forces require that objects have mass,
as matter does. As such, light remains a confusing amalgam of both being
a wave and matter, but truly being neither.
In 1924, Louis
deBroglie suggested an interesting claim. If light can show the
characteristics of both matter and wave, why couldn’t matter show
wave-like characteristics? deBroglie claimed that the wavelength of a
moving object was Planck’s constant divided by p, its momentum. Later,
this result was “proven” when both X-rays (a wave) and electrons
(matter) were fired into a thin sheet of aluminum. The pattern that was
observed from both was very similar, and wave-like in nature.
Heisenberg
Uncertainty Principle
From deBroglie’s
results, it could be seen that any moving object traveled in a wave-like
path. While the result it generally inconsequential for large objects,
the problem remained for smaller particles, such as electrons.
In order us to
measure the location of a moving electron, at a single moment in time,
light must be used to determine its location. Just as one sees by
receiving light that has bounced off objects, one can “see” an electron
by measuring light that has bounced off an electron. The light
needs to interact with the electron in someway, or there would be no
detectable change, as it would be the same as light going through a
vacuum. However, once the light hits the electron, it gives it an
instantaneous boost of kinetic energy, which increases it momentum.
Conversely, it we know an electron’s exact position, there is no way we
can learn its momentum without colliding with light. At that moment, the
electron becomes “displaced” from its original position. As such, it is
impossible to know both the position and momentum of an object exactly
with perfect accuracy.
The Heisenberg
Uncertainty Principle is a key result used in Quantum Mechanics because
it provides a boundary for the accuracy and limitations of measurement.
In order for a measurement to be made, it requires interactions between
the various objects. It is this fundamental interaction that causes this
variability. Given the Uncertainty Principle, many results are given
probabilistic interpretations where surfaces and ranges are often best
representations of electron behaviour.
Introduction
Introduction to Quantum Chemistry
Periodic Table
MO Theory
References