Download Chapter 2.2 and 7 Notes

Section 2.2 and
Chapter 7
Electron Configurations and Waves
Waves
Waves
 Frequency: number of waves in a given
unit of time
Wavelength = speed/frequency
Waves
 Electromagnetic waves travel at the
speed of light regardless of wavelength
 c = 3.0*108 m/s
Wavelength = c/frequency
Electromagnetic Spectrum
 Visible Light is between 400 (violet) to
750 (red) nanometers
 1 nanometer = 1*10-9 meters
Quantum Theory
 Proposed by Max Planck in 1900
 The energy emitted or absorbed by an
object is restricted to “pieces” of
particular size

Each piece is called a quantum.
Quantum Theory
 The energy in a quantum is proportional
to the frequency of the radiation (wave).
E = hf
E = Energy (J)
h = Planck’s constant (6.626 x 10-34 J-s)
f = frequency
Quantum Theory
 Planck’s theory proposes that energy
absorbed or emitted from an atom is
quantized
 Why do we not notice this quantization of
energy?
Quantum Theory
 Planck’s constant is extremely small. In
the everyday world, the quantization of
energy is unnoticeable.
 On the atomic level, it is extremely
significant.
The photoelectric effect
 When light is shined on a metal,
electrons are ejected from the surface

Each metal has a minimum frequency of
light required to release electrons.
 Example: Sodium metal


Releases no electrons from red light
Easily releases electrons from violet light
 How can this be explained?
The photoelectric effect
 Albert Einstein proposed that light is
composed of quanta of energy that
behave like tiny particles


Photons
Each photon carries an amount of energy
determined by Planck’s equation
The photoelectric effect
 When photon strikes an atom, it
transfers its energy to an electron

“All or nothing”
 If the photon has a high enough
frequency, it will have the energy
required to eject the electron
Dual Nature of Light
 Light behaves as both waves and
particles (photons)

Acts as a wave because it has a frequency
and wavelength

Acts as particles because they carry
definite amounts of energy
The Bohr Model
http://phet.colorado.edu/sims/hydrogenatom/hydrogen-atom_en.jnlp
Line Spectra
 Each element will absorb energy and
release energy only in specific
wavelengths and frequencies
 Atomic fingerprint
 Sodium will produce yellow light
 Lithium produces red, potassium purple
So waves act like matter…can
matter act like waves?
 It turns out that matter does act like
waves in that it moves with wavelengths
and frequencies.
 Not noticed on macroscopic level, but it
is on the atomic level.
 Electrons move in wave-like forms
Heisenberg’s Uncertainty
Principle
 Proposed in 1927 by Werner Heisenberg
 The position and momentum of an
electron cannot simultaneously be
measured and known exactly

You can never know exactly where an
electron is or how it is moving
Electron Orbitals
 Electron orbitals do not show the exact
path of an electron.

Shows where an electron with a given
energy is likely to be located.
 Electrons are divided into energy levels

Each energy level is divided into sublevels
of different orbitals
 4 basic shapes (types) of orbitals

s, p, d, and f
s orbitals
 Lowest energy
 Holds up to 2
electrons
 1 orbital per energy
level
p orbitals
 3 per energy level
 Each can hold 2
electrons
 6 total electrons per
energy level
 First p orbitals are in
the second energy
level
d and f orbitals
 d orbitals:



5 orbitals, 2 electrons each
10 total electrons in each energy level
Do not show up until the 3rd energy level
 f orbitals



7 orbitals, 2 electrons each
14 total electrons in each energy level
Do not show up until the 4th energy level
Electron Configurations
 The distribution of electrons throughout
the orbitals of an atom
 Electrons will occupy the lowest energy
orbitals available
Electron Spin
 Two electrons in an orbital will have
opposite spins.


One may spin counterclockwise, the other
will spin clockwise
The direction is not important. What’s
important is that they spin opposite each
other.
Orbital Notation
 Same as electron configuration except it
shows the actual orbital and spin of each
electron.