Download Radiant Energy Electromagnetic Wave Electromagnetic Wave

Radiant Energy
• The light form that
travels through space
• what makes you feel
hot in bright sunlight
• has properties of both
waves and particles
• A form of
electromagnetic
radiation
Electromagnetic Wave
Electromagnetic Wave
• Electric and magnetic fields oscillating
at right angles to each other and to
the direction of the motion of the
wave
Electric field
Magnetic field
Properties of a Wave
• Amplitude - height of a wave from
the origin to the crest (peak);
determines brightness or intensity of
light
amplitude
Direction of
wave travel
Properties of a Wave
• Wavelength - distance between two
successive crests or troughs; distance
the wave travels as it completes one
full cycle of up & down movement
wavelength
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Properties of a Wave
• Visible light wavelengths range from
400-700 nm
Properties of a Wave
• Frequency - tells how fast a wave is
oscillating up and down (cycles per
second, 1/s, s-1, Hz)
• Speed of light is always 3.00 x 108 m/s
• THE SHORTER THE WAVELENGTH THE
GREATER THE FREQUENCY!!!!!!!!!!!!!!!!
Lambda λ = wavelength
Nu ν = frequency
C = speed of light
Electromagnetic Spectrum
Electromagnetic Spectrum
Electromagnetic Spectrum
Quantum Theory
• Visible spectrum - continuous; each color
fades into the next. Each color has its own
wavelength and frequency. A very small
part of the spectrum but the only part we
can see!
• Infrared - radiant heat (ex: heat lamp)
• Microwaves - transfer energy to the
moisture in food
• 1900 - Max Planck predicted how the
spectrum of radiation
emitted from an object
changes with
temperature
• There is a restriction on
the amounts of energy
an object emits or
absorbs and that each
piece of energy is a
quantum
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Quantum Theory
• E = amount of energy
E = hν
• h = 6.6262 x 10-34 J·s
(Planck’s constant)
KNOW!!
• ν = frequency
• Can use to determine the
temperatures of planets, stars,
etc.
Photoelectric Effect
• 1905 - Einstein - used
Planck’s constant to
explain the
photoelectric effect
• Electrons are ejected
from the surface of a
metal when light shines
on the metal
Photoelectric Effect
• When a photon strikes
the metal, the energy
is transferred to the
metal’s electrons
• The greater the
frequency, the more
energy
• If the photon has enough energy, the electron
receives the energy and is ejected from the metal’s
surface
Quantum Energy
or
• Quantized energy is like stair steps
• We don’t feel this because the
“steps” are so small!! (equal to h)
• A small change on our scale = HUGE
change on atomic scale!
Photoelectric Effect
• Einstein proposed that light consists of
photons
• Photon - a quanta of energy that
behaves like a tiny particle
• Each photon carries an amount of
energy given by E=hν
• Photons can transfer energy to a
substance when they strike the
surface
Photoelectric Effect
• Photons behave like particles traveling
at the speed of light, with frequencies
AND wavelengths!
• X-rays = photons with high frequency,
high energy; can cause damage to
cells!
**Thus, light has a dual nature - it is
like a particle AND a wave (waveparticle duality)**
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Line Spectrum
Line Spectrum
• Spectrum with only certain colors or
wavelengths
• All elements emit light when heated
or when electricity is passed through
them when they are in their gas state
• When elements absorb energy, they
re-emit that energy as light with only
certain wavelengths
• Gives each element a unique atomic
emission spectrum (“fingerprint”)
Line Spectrum
Line Spectrum
• 1911 - Niels Bohr said that to get
spectral lines, energy
of electrons must be
quantized
• Electrons are only
allowed certain orbits
corresponding to
different amounts of
energy
Quantum electrons!
• Bohr gave each energy level (orbit) a
quantum number
• Lowest energy level = ground state,
closest to nucleus, n = 1
• Electron plus right amount of energy =
EXCITED STATE!
• Excited states have quantum numbers
n=2, n=3, n=4…..larger orbits are
farther from the nucleus
Quantum electrons!
• When energy
(radiation) is absorbed
- electron jumps from
ground state to
excited state!
• Radiation is emitted
when electron falls
back down, resulting in
visible emissions of
light!
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More Quantum Things
• Light travels through space
like a wave, but behaves like
a stream of particles when
it hits matter
• 1924 - Louis de Broglie said wavelike behavior of
particles = matter waves;
predicted that all moving
objects have a wavelike
behavior!
More Quantum Things
• 1927 - Heisenberg said that the position
and momentum of a
moving object cannot be
simultaneously
measured
• You can’t pin down an
electron, but you have
a good idea of where it
will probably be!
A New Approach to the Atom
Quantum-Mechanical Model
• Known:
• Explains the
properties by treating
electrons as a wave
with quantized energy
• Electron density density of an electron
cloud; high probability
= high density, low
probability = low
density
– Electrons can only have certain amounts
of energy because energy is quantized
(only certain orbits correspond to certain
amounts of energy)
– Electrons have wave-like behavior
– The exact position and momentum of an
electron is unknown
Atomic Orbitals
Atomic Orbitals
• Atomic orbital- a region around the
nucleus of an atom where an electron
with a given energy is likely to be
found
• Each has a specific shape, size, and
energy. The amount of energy in an
electron determines the orbital in
which it is located
• S-orbital: spherical
• P-orbital: dumbbell
• D- and f- orbitals: complex!
• Go here to play with orbitals!
http://www.d.umn.edu/~pkiprof/ChemW
ebV2/AOs/index.html
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Principal Energy Levels
• Main energy levels given the principal
quantum numbers n=1, n=2, etc
• Energy increases as n increases
• Each energy level has one or more sublevels
(s, p, d, f)
• Number before sublevel = quantum number
(ex: 4s)
• Superscript = number of electrons in
sublevel (ex: 4s2)
Sublevels
•
•
•
•
S sublevel = 1 s-orbital
P sublevel = 3 p-orbitals
D sublevel = 5 d-orbitals
F sublevel = 7 f-orbitals
• Pauli Exclusion Principle each orbital of an atom
can hold only 2 electrons
with opposite spins, thus
determining how many
electrons are allowed in
each principle quantum
energy level and sublevel
• Paired = 2 electrons,
unpaired = 1 electron
• So principal energy levels are made
out of sublevels, which are made out
of orbitals
• The higher the number and the
orbital, the more energy it has
• The 2p sublevel has 3 “p-orbitals” of
equal energy (2px, 2py, 2pz)
Electron Spin
Increasing
energy!
• EACH ORBITAL HOLDS 2 ELECTRONS!
Electron Spin
Principal Energy Levels
• Electrons spin! (but not really!)
• Clockwise (north pole points up)
• Counterclockwise (north pole points
down)
• Creates a magnetic field; a pair with
opposite spins cancel each other; a
pair with parallel spins = magnetic
effect!
Electron Configuration
• The distribution of electrons among
orbitals describing where electrons
are found and what energies they
have
• Determined by distributing among
energy levels, sublevels, and orbitals
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Electron Configuration
• Aufbau principle - electrons are
added one at a time to the LOWEST
ENERGY orbitals until all electrons of
the atom have been accounted for
• Aufbau - german for “construction”,
theory came from Bohr
• Exceptions - Cr, Cu, inner transition
elements
Electron Configuration
• Hund’s Rule - electrons
occupy equal energy
orbitals so that a
maximum number of
UNpaired electrons
results
Orbitals fill up in
this order!
Electron Configuration
• Orbital diagrams - combination of all
information about electron placement!
• Know how to write these!
Which is the correct
orbital diagram for
oxygen?
Must follow Aufbau,
Pauli exclusion, and
Hund’s rule!
Electron Configuration
• Compact form - does not show orbitals and
electron spins
• Principal quantum energy number, sublevels,
electrons (superscripts)
ex: oxygen : 1s22s22p4
• Sum of superscripts = atomic number of
element!
Summary
• Energy levels - n=1, n=2, n=3, etc
• Each energy level has SUBLEVELS (s,
p, d, f)
• S sublevel - 1 orbital - 2 electrons
• P sublevel - 3 orbitals - 6 electrons
• D sublevel - 5 orbitals - 10 electrons
• F sublevel - 7 orbitals - 14 electrons
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