Download Unit 5 Power Point

Electronic Structure and Light
UNIT 5
Electromagnetic
Radiation
• Energy that moves through space in the form
of a wave.
• Also known as Radiant energy or Light!
• Types include: Radio waves, microwaves,
infrared rays, visible light, ultraviolet rays, xrays and gamma rays
Measurable Properties of Waves
• Speed = distance traveled
Time
• SI unit = m/s
• Speed in a vacuum is 3.00 x 108 m/s
• Wavelength = distance between identical
points on back to back waves.
• SI unit = meter
other units = nm, km, or μm
Electromagnetic
Spectrum
1. Which type of
electromagnetic
radiation has
the shortest
wavelength?
The longest?
R
O
Y
G
B
V
3. Which color has the longest wavelength?
4. Which color has the shortest wavelength?
2. What is the
wavelength of
visible light?
Measurable Properties of Waves
• Frequency = number of complete
wavelengths that pass a given point in 1
second
• SI unit = s1- or 1/s. Hertz (Hz) is a common unit
• 1 Hz = 1/s
Which wave has
the highest
frequency?
Electromagnetic Spectrum
Light
• c = speed of light = 3.00 x 108
• Wavelength = 
• Frequency = ⱱ
Related by
m/s
c = ⱱ
• Frequency and wavelength are _________ proportional.
Homework
• Complete 6.1-6.6 on the handout. Write the
answers on a separate sheet of paper.
A revolution is born
• Physcists wanted to understand the relationship
between the temperature and the intensity and
wavelengths of the emitted radiation.
• Prevailing laws could not explain three
phenomena
• Black body radiation
• Photoelectric effect
• Emission Spectra.
Max Planck (1858-1947)
• Explained how radiation is emitted by hot
objects (black-body radiation)
• Assumed energy can be released (or absorbed)
by atoms or molecules only in discrete quantities
• Quantum (plural quanta) – discrete chunk of
energy, a fixed amount.
• E = hf
• Planck’s constant, h = 6.63 x 10-34 J s
• Energy and frequency are _______ proportional
Albert Einstein (18791955)
• Extended Planck’s theory, Light is composed of particles
or photons
• Photons - packets of energy that behave like particles;
particles of light
• Energy of photon = E = hf (radiant energy itself is
quantized)
• Explained photoelectric effect (When
electrons are ejected off of the surface of a
metal when light of a certain minimum
frequency shines on it.
• If photons have enough energy when they
strike the metal, they will pass that energy
to the electrons causing the electrons to
fly off!
The Wave Nature of
Light
• electrons, electrons, electrons
• Knowing the arrangement, number of
electrons, and energy of the electrons
(electronic structure) in an atom is the key to
understanding the physical and chemical
properties of an element.
• What does the study of light have to do with
understanding electronic structure?
Is light composed of
waves or particles?
• What do you think?
Spectrum
• Continuous Spectrum - Continuous range of colors
• Line (Emission) Spectrum – A spectrum containing
radiation of specific wavelengths emitted by a
substance.
• Absorption Spectrum – A spectrum containing
radiation of all wavelengths except those absorbed
by a substance.
• Who was the first to shine light through a prism and
study it?
Hydrogen Spectra
• Absorption spectrum
• Line spectrum
Emission Spectrum of
Elements
Bohr’s Model
Three postulates:
• Assumed that electrons move in circular orbits around the
nucleus. These orbits correspond to certain definite
amounts of energy.
• An electron in a permitted orbit has a specific energy and is
in an allowed energy state. It will not spiral into the nucleus.
• Energy is emitted or absorbed by an electron as it changes
from one energy state to another . This energy exists as a
photon.
• Ground State: Electrons are as close to the nucleus as they can
be; they are in the lowest energy level
• Excited State: Electrons are not as close to the nucleus as they
can be; they are in a higher energy level
Louis de Broglie
 Dual nature of the electron - suggests that if
light can behave like a stream of particles
then electrons may possess wave properties.
 Soon after this was published it was
experimentally demonstrated.
Werner Heisenberg
 A wave extends into space, therefore its exact
location can not be found.
 Also because photons are used to detect
electrons and the energy of each of these
particles is similar, then any attempt to locate
an electron with a photon will knock the
elctron off its course.
Heisenberg Uncertainty
Principle
 It is impossible to know the momentum and
position of a particle with certainty.
 Because of this we know the electron does not
orbit the nucleus in a well defined path (as Bohr
thought)
Quantum Mechanical Model
• The Schrodinger Equation incorporates both
the wave and particle behavior of electrons.
• The location of an electron cannot be described so
simply.
• Launched a field of physics called quantum
mechanics.
Quantum Mechanical Model
 Quantum mechanics mathematically defines
the region where the electron might be at a
given time.
 Electron Density – the probability that an
electron will be found in a particular region of
an atom.
Quantum Mechanical Model v. Bohr Model
Quantum Model
Bohr Model
Treats electron as a wave
Treats electron as a particle
Energy and location described in terms
of probability (called orbitals)
Energy and location is described in
terms of a definite orbit.
Electron Density
Orbitals
 Electrons do not travel around the nucleus,
instead they exist in regions called orbitals.
 Region in the atom where there is a high
probability of finding an electron.
 Each has a characteristic shape and energy.
Quantum Numbers
 Principal quantum number (n): 1,2,3, etc.
 Determines the size of the orbital, distance from
the nucleus, and the energy level.
 Azimuthal quantum number (l ): can have
values from 0 to n-1.
 Determines shape of the orbital.
n
Value for l
Type of
sublevel
l
1
0
2
0, 1
s
0
3
0, 1, 2
p
1
4
0, 1, 2, 3
d
2
5
0, 1, 2, 3, 4
f
3
Quantum Numbers
 Magnetic quantum number (ml): values of l
to -l
 Orientation of the orbital (x, y, or z axis)
 Subshell in orbital
l
Possible values for ml
0
0
1
-1, 0, 1
2
-2, -1, 0, 1, 2
3
-3, -2, -1, 0, 1, 2, 3
Quantum Numbers
 Spin magnetic quantum number (ms): +1/2 or
-1/2
 Electrons move in opposite directions within
subshell
Electron Configuration
 H: 1s1
1 – gives energy level
s – gives orbital shape
1s – gives subshell
Aufbau Principle
 As protons are added one by one to the
nucleus to build up the elements, electrons
are similarly added to the atomic orbitals.
Lower energy subshells are filled first.
Electron Configuration
 Pauli Exclusion Principle – No two electrons in
an atom can have the same four quantum
numbers. They will have different spins.
Electron Configuration
 Hund’s Rule – the most stable arrangement of
electrons in subshells is the one with the
greatest number of parallel spins. Place
electrons in each orbital one at a time, then go
back and place two in them if neccesary.