Transcript Electrons

Light and Electrons
Electromagnetic Radiation
Light is electromagnetic radiation:
combined electric and magnetic
waves Source
Electric vector
Magnetic vector
direction of propagation
Electromagnetic Radiation
Light is more than what we can see…
Electromagnetic Radiation
Subatomic particles (electron, photon,
proton, etc) exhibit both PARTICLE
and WAVE properties. This is known as
Wave-Particle Duality.
Diffraction: wave-like
Photoelectric Effect: particle-like
Electromagnetic Radiation
Wave Properties of Light:
1. It’s fast! …c = 3.0 x 108 m/s
2. It relfects, refracts, diffracts
(Transverse wave)
3.
Electromagnetic Radiation
All light waves have
frequency
wavelength
symbol:
l (Greek “lambda”)
units:
f
“cycles per sec” = Hertz
“distance” (m, nm)
c=l f
where c = velocity of light
= 3.00 x 108 m/sec
Increasing frequency
Electromagnetic Radiation
Example: Red light has l = 700 nm.
Calculate the frequency, f.
C
f=
l
=
3.00 x 10
8
7.00 x 10
m/s
-7
m
= 4.29 x 10
14
Hz
Electromagnetic Radiation
Particle Properties of Light:
1. A particle of light is called a photon
2. Energy of a photon is calculated by
E=h·f
where E = energy (Joules, J)
f = frequency (Hertz, Hz, 1/sec)
h = Planck’s constant 6.63 x 10
-34
J·s
Electromagnetic Radiation
Albert Einstein postulates the
Photoelectric Effect to explain
two observations:
1. No electrons are observed until
a minimum energy is applied.
2. Number of electrons ejected
depends upon light intensity –
not light frequency!
Light is created by the
Photoelectric Effect
Electromagnetic Radiation
The photoelectric effect and the idea of
discrete, quantized energies neatly explain the
observation of emission spectra.
Electromagnetic Radiation
Example: Red light has l = 700 nm.
Calculate the energy per photon.
E = hf and c = lf
So f = c/l and E = hc/l
-34
8
E = (6.63 x 10 Js)(3.0 x 10 m/s)
-9
700 x 10 m
E = 2.84 x 10
-19
J
Electron Orbitals
While thinking about the emission spectrum of hydrogen, Neils
Bohr came up with the planetary model of the atom. In this
model, electrons can only orbit the nucleus at discrete distances
and particular orbital shape.
Orbital model of Na
Sharp-line spectrum of H
Neils Bohr
Electron Orbitals (n)
n = energy level or shell (n = 1, 2, 3, 4, 5, 6, 7)
1. Energy levels are whole numbers
2. The maximum number of electrons in each
energy level equals 2n2.
3. The rows of the periodic table correspond to
energy levels.
1. Whole number energy levels – like a standing wave
Electron Orbitals (n)
3. The rows (periods) of the periodic table correspond
to energy levels.
Electron Orbitals (l)
l = subshell (s, p, d, f, g, h, i, j…)
1. s, p, d, and f are named after the four lines in
the hydrogen emission spectrum…Sharp,
Principle, Diffuse, Fundamental.
2. Each subshell has a different shape
3. The number of
subshells in an
energy level is
equal to the
number of the
level.
Energy
Level
1
Number of
Sublevels
1
Name of
sublevels
s
2
2
s, p
3
3
s, p ,d
energy
4
4
s, p, d, f
Electron Orbitals (l)
1. Sharp, Principle, Diffuse, and Fundamental refer to
the way the spectral lines look. It was thought that
electrons traveling between certain energy
sublevels
produced
those certain
lines. This
was not
correct, but
the names
stuck.
Electron Orbitals (l)
2. Each subshell has a different shape
s-orbital
1. Has a
spherical
shape
2. Can hold up to
2 electrons
3. Lowest energy
subshell
Electron Orbitals (l)
p-orbitals
1.Said to have a “dumbbell shape”
2.Can hold up to 6 electrons
Electron Orbitals (l)
d-orbitals
1. Said to
have a
“clover
leaf”
shape
2. Can hold
up to 10
electrons
d-orbitals
combined
orbitals
Electron Orbitals (l)
f-orbitals
1. Can hold up to 14 electrons
f-orbitals
combined
orbitals
Electron Orbitals
To write a ground-state electron configuration:
1. Determine how many electrons are present.
2. Follow the Aufbau Diagram (Diagonal Rule)
Aufbau Diagram
Electron Orbitals
Example: Write the ground-state electron
configuration for nitrogen.
1. Nitrogen has 7 electrons
2. Follow the Aufbau Diagram
3. N: 1s22s22p3
Electron Orbitals
So why does it work like this?
1. Pauli Exclusion Principle – states that “no two
electrons in an atom can have the same set of
four quantum numbers.” In other words, no
atomic orbital can contain more than two
electrons.
2. Hund’s Rule – The most stable arrangement of
electrons around an atom is one with the
maximum number of unpaired electrons. This
minimizes electron-electron repulsion.
Electron Orbitals
So why does it work like this? (cont.)
3. Aufbau Principle – Electrons occupy the
lowest energy state possible.
4. Heisenberg Uncertainty Principle – The
orbitals are probabilities – not shapes in
space like planetary orbits. The uncertainty
principle states that you cannot know the
location and velocity of an electron
simultaneously.
s-orbitals in Zinc
p-orbitals in Zinc
Electron Configuration Shortcut…
Electron Orbitals
Electron orbital notation goes one step further
than electron configuration. It describes,
specifically, each electron.
Compare them
Electron Configuration of Oxygen: 1s22s22p4
Electron Orbital Notation of Oxygen:
.
1s
.
2s
.
.
2p
.
Electron Orbitals
Orbital
s
Notation
.
or
1s
2s
.
p
.
3s
.
.
.
.
.
.
.
.
2p
.
d
f
. or
.
3d
.
.
.
4f
.
Electron Orbitals
Example: What is the electron orbital notation
for sulfur?
.
.
1s
2s
.
.
.
.
2p
.
3s
.
.
3p
Example: What is the non-core electron orbital
notation for gold?
[Xe]
.
6s
.
.
.
5d
.
.
Electron Orbitals
Example: What is the non-core electron orbital
notation for gold?
[Xe]
.
.
.
6s
.
.
5d
Electrons are
more stable
in full or halffull orbitals.
…or more likely,
[Xe]
.
6s
.
.
.
5d
.
.
.
Electron Orbitals
Octet Rule: Atoms will gain or lose electrons to
achieve a full valence shell (usually this means
8 electrons).
Oxidation State: The value of the charge on an
ion (positive or negative), after the atom has
achieved a full valence shell.
- metals tend to lose electrons, forming
positive (+) ions (cations).
- non-metals tend to gain electrons,
forming negative (-) ions (anions).
Electron
Orbitals States
Periodic Table
of Oxidation