Electrons and Quantum Mechanics
Unit 5
Electrons
• Rutherford described
the dense center of the
atom called the
nucleus.
• But the Electrons spin
around the outside of
that nucleus.
– Provide the chemical
properties of the atoms.
– Responsible for color
and reactivity.
Energy
• Energy is transmitted
from one place to
another.
– Light carries this energy.
– Converted into heat.
• Light is called
Electromagnetic
Radiation.
Electromagnetic Spectrum
• Radio
• Infrared
• Visible Light
– ROY G BIV
• Ultraviolet
• X Rays
• Gamma Rays
Light
• Light travels as a wave.
• Wave Properties
– Wavelength (λ) =
distance between two
waves (m)
– Frequency (f) = number
of peaks per second (Hz)
– Speed of Light (c) = how
fast light moves.
Light
• Light Equation
c= ƒλ
• Speed of light is a
constant = 3 x 108 m/s
• Nothing travel faster
than the speed of light!
– Maybe?!?!?!?!?!?!?!?!?
Light
• The Dual Nature of Light
– Light carries energy
through space like a
wave.
– Light also behaves like a
particle?!?
• A beam of light is made of
tiny packets of energy
called PHOTONS!
• Which travel in waves!?!
Light
• The Energy of a photon
depends on its frequency.
– So is the color of light!!!
E = hƒ
• ELECTRONS are like
photons!
– Act as waves and particles.
– Orbit the nucleus in a
wave-like motion.
Blackbody Radiation
• Rutherford could never
explain why objects
change colors when
they are heated.
• As the object heats, it
must give off electrons
of certain frequencies
and energies.
Photoelectric Effect
• Similarly, light on a
metal object can knock
off electrons.
– Shine different colors on
a metal.
– Measure the number of
electrons knocked off.
– Found that no electrons
were knocked off below
a certain frequency.
The Bohr Model
• Proposed the electrons
orbit the nucleus with
fixed energies.
– Called Energy Levels
– Much like the rungs of a
ladder.
• Quantum describes the
amount of energy
required to move an
electron from one level
to another.
The Bohr Model
• Ground State
– Lowest possible energy of
an electron.
– Normal location
• Excited State
– If electron absorbs energy,
it moves up an energy level
(absorption)
– If an electron gives off
energy, it moves down an
energy level (emission).
The Bohr Model
Atomic Spectra
• Hydrogen Atom Line
Emission Spectrum
– Expected continuous
spectrum of light
– But only specific
frequencies were given
off.
•
•
•
•
Red (656.6 nm)
Blue-green (486.1 nm)
Violet (434.1 nm)
Violet (419.2 nm)
Atomic Spectra
• Shine a light on an Atom
– When atoms absorb
energy, electrons move to
higher energy levels.
– When atoms release the
energy, electrons return to
the lower energy level.
• Atomic Spectra
– Frequencies of light
emitted by a certain
element.
– No two elements have the
same spectrum.
http://student.fizika.org/~nnctc/spectra.htm
Flame Tests
• Because no two atoms
produce the same
spectrum, elements can
be identified by the
colors they emit.
• Spectral Analysis uses
this properties to
identify elements.
Quantum Mechanics
• Max Planck (1900)
– Founder of Quantum
Mechanics
E = hf
• Albert Einstein (1905)
– Wave-Particle Duality
– Electrons are small
particles that move like
waves.
Quantum Mechanics
• Neils Bohr (1922)
– Electrons orbit in distinct
energy levels.
• Louis de Brogelie (1923)
– Wave Mechanics says
that ALL MATTER
behaves like waves.
mv/λ = h
Quantum Mechanics
• Werner Heisenberg (1927)
– Principle of Indeterminacy
– You can’t know both the
position and the velocity of
an electron.
• Erwin Schrödinger (1930)
– Used wave mechanics to
show the PROBABLE location
of an electron.
– Electrons exist in 3D clouds
of probability!!!
Quantum Mechanical Model
• Uses Schrodinger’s
equation to predict the
probable location of an
electron.
– Determines the energies
an electron is allowed to
have.
– Determines how likely it
is to find the electron in
various locations around
the nucleus.
Quantum Numbers
• Describes the location
and behavior of an
electron
– Like an electron’s
address
– No two electrons can
have the same quantum
numbers.
• Four Numbers
Quantum Numbers
• Principle (1st) Quantum
Number (n)
– The Energy Level
– Describes the size of the
cloud and the distance of
the cloud from the
nucleus.
– Shows the number of
electrons
n = 1 = 2 electrons
n = 2 = 8 en = 3 = 18 en = 4 = 32 e-
Quantum Numbers
• 2nd Quantum Number (l)
– Each energy level has
sublevels.
– The number of sublevels
equals n.
– Sublevels are called:
s = spherical
p = peanut-shaped
d = daisy-shaped
f = unknown?
Quantum Numbers
• 3rd Quantum Number (ml)
– Divides sublevels into orbitals.
– Tells the shape the electron
moves in.
– Number of orbitals = n2
– Examples
s = 1 orbital
p = 3 orbitals
d = 5 orbitals
f = 7 orbitals
Quantum Numbers
• 4th Quantum Number (ms)
– Describes the electron’s
spin.
– Only two electrons fit in an
orbital.
– Their charges repel causing
them to spin in opposite
directions (+½ or –½)
– Use up and down arrows.
Quantum Numbers
• Pauli Exclusion Principle
– No two electrons can
have the same set of 4
quantum numbers.
– The electrons repel each
other.
• Hund’s Rule
– Every orbital must get
one electron before
doubling up.
Quantum Numbers
• Diagonal Rule
– Electrons fill orbitals in
predictable patterns
– Some People Do Forget
– Electrons dill the lowest
energy level possible.
1s
2s
3s
4s
5s
2p
3p
4p
5p
3d
4d
5d
4f
5f
Orbital Notation
• Draw out the locations
of each electron in an
atom with arrows.
Electron Configuration
• Write out the
configurations of
electrons using
superscripts.
• Examples:
– H = 1s1
– He = 1s2
Electron Configurations
• Noble Gas Shorthand
– Write the Noble Gas just
before the element.
– Add the remainder of
the configuration.
Lewis Dot Diagrams
• A way to show the
number and position of
the valence electrons.
– Outermost energy level
– Look at the column
number to get this number.
• Use the chemical symbol
and number of valence
electrons.
– All four sides must have a
dot before you double up.
p orbitals
s orbital
p1
p2
X
p3
s