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 = hĥ 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 electronsn = 2 = 8 e-
n = 3 = 18 e-
n = 4 = 32 e-
Quantum Numbers
• 2nd Quantum Number (l)– Each energy level has
sublevels.– The number of sublevels
equals n.– Sublevels are called:
s = sphericalp = peanut-shapedd = 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 orbitalp = 3 orbitalsd = 5 orbitalsf = 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.
1s2s 2p3s 3p
3d4s 4p4d 4f5s 5p5d 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.
X
p1
p3
p2 s
p orbitals s orbital