Light
• Light is a kind of electromagnetic
radiation, which is a from of
energy that exhibits wavelike
behavior as it travels through
space.
• Other forms of electromagnetic
radiation include X-rays, UV rays,
infrared rays, microwaves and
radio waves.
• All forms move at a constant
speed of 3.0 x 108 m/s
Wavelength and frequency
• Wavelength is the
distance between
corresponding points
on adjacent waves.
• Frequency is the
number of waves that
pass a given point in a
specific time (usually
1 second).
The photoelectric effect
• The photoelectric effect refers to the
emission of electrons from a metal
when light shines on the metal.
• If the light’s frequency was below a
certain minimum no electrons were
emitted, no matter how long it was
shone.
• Planck suggested that the light was
emitted as quanta of energy rather than
waves.
• E = hv
Photons
• Einstein –
electromagnetic
radiation has a dual
wave-particle nature.
• Each particle of light
carries a quantum of
energy, called a
photon.
• The lowest energy state of
an atom is its ground state.
A state in which an atom
has a higher potential
energy than its ground
state is called an excited
state.
• When an excited atom
returns to its ground state,
it gives off energy it
gained in the form of
electromagnetic radiation.
Emission Spectra
• Emission spectra –
produced by passing
electric current through a
vacuum tube containing
gas at low pressures, and
then shone through a
prism.
• Not continuous. Hydrogen
emits only at specific
frequency – energy
differences between the
atoms’ energy states are
fixed.
Bohr’s Model
• Neils Bohr proposed a
model that linked the
atom’s electron with
photon emission.
• He said that electrons
could only orbit in
fixed energy states.
• His model explained
only the hydrogen
atom.
Quantum model of the atom
• De Broglie – electrons could also be considered
as waves. Electron waves could exist only at
specific frequencies.
• Heisenberg uncertainty principle – it is
impossible to determine at the same time, both
the position and velocity of an electron or any
other particle
• Quantum theory describes mathematically the
wave properties of electrons and other very small
particles.
Orbitals
• Wave functions give only the probability of
finding an electron at a given place around
the nucleus
• Do not travel in neat orbits, but exist in
certain regions called orbitals.
• An orbital is the three dimensional region
around the nucleus that indicates the
probable location of an electron.
Quantum numbers
• The principal quantum
number (n) indicates
the main energy level
occupied by an
electron
• The angular momentum quantum number
(l) indicates the shape of the orbital.
• s is spherical
• p is dumb-bell shape
• d is complex
• Magnetic quantum number
indicates the orientation of an
orbital around the nucleus.
• Spin quantum number – only two
possible values
Electron Configuration
• Aufbau’s principle
Bohr Model – energy levels
An electron can be
in any of these
energy levels, but
cannot be found
in between.
• In the ground state –
atom is unexcited –
electrons are in the
lowest possible energy
level.
• If an atom is excited
the electron may go
up to the next level.
• When they fall back
they emit light.
• This is an element’s
emission spectrum
Emission spectra
• Later complications – De Broglie –
electrons act both as particles and waves
• Heisenberg – you can never know
exactly the position and velocity of an
electron (or any other particle)
• Schrodinger – wave equations –
(quantum theory)
• An orbital is the region around the
nucleus where an electron can be found.
Why do we need to know
where electrons are found in
the atom?
• Because the chemical behavior
(reactivity) of an element
depends on the energy levels of
its electrons
The Principal quantum number,
n, is the main energy level
occupied by an electron
Sublevels – the shape of the
orbital
d-orbitals
Spin quantum number
A single orbital can
hold a maximum
of two electrons
with opposite
spins
Filling electron orbitals
Filling electron orbitals
Filling energy levels
(Aufbau’s rule)