Transcript Chapter 6 review

Chapter 6
Review and Breathe
The Wave Nature of Light
• Electromagnetic
radiation is one way
energy travels
through space.
• Wavelength is
inversely proportional
to frequency
• =c/
The Particle Nature of Light
• Energy is gained or
lost in whole number
multiples.
• ΔE=nh
• This energy is
Quantized.
• Electromagnetic
radiation is quantized
as photons.
Dual Nature of Light
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Photons have energy of Ephoton = hc/
Energy is also E=mc2 (c=speed of light)
Thus: m=E/c2
Substitute E=hc/ in the above equation
m=hc/c2 
And finally: m=h/c 
Light acts as a wave and has mass
Do particles have wave
characteristics?
• If a photon has mass m=h/c  while it is
moving…
• Then a particle moving at a velocity v has
a wavelength using the equation m=h/v 
• Solve for  and =h/mv
• This is de Broglie’s equation.
Atomic spectrum of Hydrogen
• When light is passed
through a prism one
gets an emission
spectrum.
• When all wavelengths
are possible one gets
a continuous
spectrum.
• When energy is
quantized the
spectrum is a line
spectrum, or discrete.
How do electrons emit light?
The Bohr model
• Bohr proposed that in
a hydrogen atom the
electron orbits the
nucleus in allowed
circular orbits.
• Each orbit has an
energy associated
with it.
• E=-2.178x10-18J
(Z2/n2)
Quantum Mechanical Model
• Bohr’s model doesn’t
work for atoms larger
than hydrogen.
• Electrons are not
behaving in a manner
that agrees with the
circular orbit model.
Wave function
• Electrons in the
higher energy levels
are acting more like
standing waves than
like particles.
• Schrödinger looked at
the wave function of
the electron. This
wave function
describes the
electron’s orbital.
But where is the electron?
• Heisenberg’s
uncertainty principle
states that we cannot
know with certainty
both how fast an
electron is moving
and where it is.
• Δx * Δ(mv)> h/4
• Probability distribution
Quantum numbers
• each energy level is designated by the value n
which is an integer from 1 (lowest energy or
"ground state") on up
• the number of types of orbitals possible on an
energy level is also equal to n
• the maximum number of actual orbitals on an
energy level is equal to n2
• the maximum number of electrons in an orbital is
equal to 2
• the maximum number of electrons on an energy
level is equal to 2n2
Quantum numbers
• Principle quantum number, n=integral
values 1, 2, 3 … Represents the energy
level.
• Angular quantum number, l = 0, 1, 2 etc
for n-1. It represents the shape of the
orbital.
• Magnetic quantum number m= -l to l. It is
related to the orientation of the orbital in
relationship to other orbitals in the atom.
Pauli Exclusion Principle
• In given atom, no two electrons may have
the same set of quantum numbers. This is
known as the Pauli Exclusion Principle.
• Since each orbital may hold up to two
elections. Each election is assigned a
separate spin.
• +½ and -½
Summary Table
Angular quantum numbers
Take note of the nodes!
• For l= 0,
1,
2,
and bottom row
3.
Aufbau Principle
• As protons are added
to a nucleus to build
up elements,
electrons are added
too.
• These electrons are
added into the lower
energy orbitals first.
Hund’s Rule
• The lowest energy
configuration In an
orbital is one having
the maximum number
of unpaired electrons
allowed by the Pauli
Principle in a set of
degenerative orbitals.
Electron Configuration and the
Periodic Table
• Elements’ reactivity is
based on its valence
electrons.
• The periodic table
demonstrates the
valence electrons of
each group.