Transcript d - Solon City Schools

Light and Energy
Electromagnetic Radiation is a form of energy that
emits wave-like behavior as it travels through space.
Examples: •Visible Light
•Microwaves
•X-rays
•Radio waves
Light and Energy
Several characteristics can be used to describe these waves.
2 sec
Amplitude (a) is the
height of a wave
from the origin to
the cresting point.
Frequency (n) is the
number of waves that
pass a given point per
second.
Time
Wavelength (l) is
the shortest distance
between equivalent
points on a continuous
wave.
Light and Energy
Which of the following waves has a higher frequency?
Light and Energy
Electromagnetic radiation can be displayed in the
electromagnetic spectrum. This spectrum places all
of the wave types in order based on wavelength (l)
and frequency (n).
Light and Energy
All electromagnetic waves travel at the same
speed in a vacuum
This speed (c) is 3.00 x 108 m/s in a vacuum.
c is more commonly referred to as the
speed of light.
c = ln
Sample Problem
• Green light has a wavelength of 545nm.
What is its frequency?
Light and Energy
In 1900 Max Planck helped us move
toward a better understanding of
electromagnetic radiation.
Matter can gain or lose energy only in small,
specific amounts called quanta.
A quantum is the minimum amount of energy
that can be gained or lost by an atom.
Light and Energy
Planck then demonstrated the relationship between that
quantum and the frequency of the emitted radiation.
E = hn
E = energy
n = frequency
h = Planck’s Constant
6.626 x 10-34 J s
Atomic Emission Spectrum
• Electrons gain energy and ‘jump’ to a
higher energy level.
• This is an unstable excited state.
• The electrons release energy when they
return to their ground state.
• We see this energy as light.
Atomic Emission Spectra
The Bohr Quantum Model
The electron in a hydrogen
atom moves around the
nucleus only in certain
allowed circular orbits
Ground State: The lowest
possible energy level.
Excited State: Higher energy
level.
Light and Energy
Light and energy can’t always be explained using waves.
In 1905 Albert Einstein proposed that electromagnetic
radiation has both wavelike and particle like natures.
While a beam of light has many wavelike
characteristics, the beam can also be thought
of as a stream of tiny particles or bundles of
energy called photons.
A photon is a particle of electromagnetic radiation
with no rest mass that carries a quantum of
energy.
Wave-Particle Duality Theory
Is energy a wave like light, or a particle?
 BOTH
Atomic Theory
Louis de Broglie (1892 – 1987)
•Matter has characteristics of both waves and
particles
•Electrons move in wavelike motion in the
circular orbits
Heisenberg Uncertainty
Principle
•There is a fundamental limitation
to just how precisely we can know
both the position and momentum of
a particle at a given time
•Watch this video!
Schrodinger
• Came up with an equation to describe the
probability of finding an electron within a
3D area of space.
• These equations are called wavefunctions.
The wavefunctions describe orbitals.
• Orbitals are the building block of electron
arrangement.
The Quantum Model!
• Atomic orbitals are used to describe the
possible position and energy of an
electron.
Orbitals
• The location of an electron is described
with 4 terms.
- Energy Level
- Sublevel
- Orbital
- Spin
Energy Level
• Describes the energy and distance from the
nucleus.
• Whole numbers, ranging from 1 to 7.
Orbital Shapes – s sublevel
- An orbital can contain up to 2 electrons.
- The s sublevel contains only 1 orbital.
P Sublevel
• 3 orbitals present in this sublevel.
• Each orbital can only have 2 electrons.
D Sublevel
• 5 orbitals present in this sublevel.
F Sublevel
• 7 orbitals present in this sublevel.
s sublevel
p sublevel
d sublevel
f sublevel
Summary
Energy
Level
1
2
3
4
Sublevels
Present
# of
Orbitals
Total # of
electrons
in Energy
Level
Arrangement of E• Electrons are arranged according to a few
rules.
– Aufbau Principle
– Pauli’s Exclusion Principle
– Hund’s Rule
Aufbau Principle
• As protons are added one by one to
the nucleus to build up the elements,
electrons are similarly added to
orbitals
• Electrons fill in low energy orbitals
before high energy orbitals
Increasing energy
Aufbau Principle
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
5f
4f
3d
He with 2
electrons
* orbital energy order found
on periodic table
s sublevel
p sublevel
d sublevel
f sublevel
Pauli’s Exclusion
• Each orbital can only hold 2 electrons and
they will have opposite spins.
– Example
Hund’s Rule
• When there are multiple orbitals, one
electron goes in each before pairing takes
place.
Practice
• Step 1: Start by drawing out the orbitals in
the correct order.
• Step 2: Determine the total number of
electrons.
Practice
• Step 3: Start arranging the electrons
• Step 4: Follow the rules!
Electron Configuration
• Electron configuration is a shortcut to the
orbital notation.
• It is done by writing only the energy level
and sublevel. The number of electrons in
each is represented by an exponent.
Noble Gas Shortcut
• Uses noble gases to represent the ‘inner’
electrons.
• Noble gases are the last column on the
right.
• Step 1: Find noble gas on the row above
the element.
• Step 2: Write remaining configuration.
Valence Electrons
• Are the electrons found on the highest
energy level. Remember, the energy level
is represented by the whole number.
• How do we find them?
Valence Electrons
• Practice
– Fe
– Cu
– Ag
• Do you see a pattern?
Lewis Dot Diagrams
• A Lewis Dot diagram shows the number of
valence electrons in an atom.
• You put the element symbol and dots
around it to represent the valence
electrons.