How did we discover
electron arrangement in an
atom?
ELECTROMAGNETIC
RADIATION ! ! !
Waves
 Repeated disturbance through a medium (air, liquid)
from origin to distant points.
 Medium does not move
 Ex. Ocean waves, sound waves
Characteristics of Waves
 Wavelength
 Distance between 2 points within a wave cycle
 2 peaks
 Frequency
 # of wave cycles passing a point for a particular time
unit
 Usually seconds.
Wavelength and frequency are
inversely proportional.
c = νλ
c = speed of light, 3.0 x 10 8 m/s
 Constant
ν= frequency (s-1 or Hz)
λ= wavelength (m)
Example 1:
 Find the frequency of a green light that has a
wavelength of 545 nm.
Electromagnetic Waves
 Produced from electric charge movement
 Changes within electric and magnetic fields carried
over a distance
 No medium needed
Electromagnetic Spectrum
 Contains full range of wavelengths and frequencies
found with electromagnetic radiation
 Wavelength/frequency changes cause color changes
 Mostly invisible, visible range (390 nm -760 nm)
 Different materials absorb/transmit the spectrum
differently.
Types of Spectra
 What is a spectra?
 Spectrum– white light/radiation split into different
wavelengths and frequencies by a prism
 Continuous spectrum
 No breaks in spectrum
 Colors together
 Line spectrum
 Line pattern emitted by light from excited atoms of a
particular element
 Aided in determining atomic structure
Line Spectrum
 Pattern emitted by light from excited atoms of an
element
 Specific for each element
 Only certain wavelengths of visible spectrum present
 Used for element identification
Flame Tests
 Some atoms of elements produce visible light if heated
 Each element has a specific flame color
 Examples: Li, Na, Cs, Ca
A Bit of Quantum
Theory……
Max Planck
 1900
 Related energy and radiation
 E = hν
 h= 6.626 x 10 -34 Js (Planck’s constant)
 E = energy per photon (J)
 Quantum---smallest amount of energy
 Atoms can only absorb/emit specific quanta
Albert Einstein
 1905
 Added to Planck’s concept
 Photons—
 Bundles of light energy
 Same energy as quantum
 E = hν (energy of photon)
 Photons release energy and electrons gain energy
 Threshold frequency– minimum amount of energy needed
by photon to extract electron
THEREFORE ………
 Light is in the form of electromagnetic waves
 Photons can resemble particles
 Gave raise to the possibility of thinking about wave
AND particle qualities of subatomic particles
(electron)
Example 1
Calculate the energy found in a
photon of red light with a
wavelength of 700.0 nm
Example 2
How much energy (in joules) is found
in the radiation of the hydrogen atom
emission spectrum with a 656.3 nm
wavelength?
Example 3:
 A sodium atom emits yellow light with a wavelength
of 589 nm when it is excited. Find the energy per
photon of this light.
Coulomb’s Law
 Describes the attractive force between negative electrons and
positive nucleus.
 Force is directly related to the charge of electron and nucleus
 Force is inversely related to distance between particles
F = qe x qp
r2
(IE…an electron’s energy is dependent on distance from nucleus)
Early Models of the Atom
Bohr
 1913—hydrogen atom
structure
 Physics + quantum theory
 Electrons move in definite
orbits around the positively
charged nucleus—planetary
model
 Does not apply as atoms
increase in electron number
Bohr Model
 Electrons orbit nucleus in different energy levels
 Lower energy levels, closest to nucleus (n = 1)
 Higher energy levels increase electron’s distance from nucleus
 Electrons can “transition” or jump between energy levels through
photons
 Gain/absorb photon—higher energy level
 Lose/emit photon—lower energy level
Energy States in an Atom
 Atoms can gain or loss energy.
 Specific energy states within an atom.
 Can be counted
 Ground State = lowest energy state
 Excited State = higher energy level than
ground, gained energy
So, where does the Bohr
Model fit in?
 Electrons orbit around the nucleus at different energy
levels/orbits.
 Electron’s energy level = orbit level where electron is
located.
 Light absorption = electron moves from a state of
low energy to high energy. “becomes excited”
 Light Emitted = electron falls from an “excited” state
of energy to a lower energy level.
Ex. Li
Erwin Schrödinger
 Quantum mechanics
 1926---wave equation
 Electrons behave more like
waves than particles
Heisenberg’s Uncertainty
Principle
 Electron’s location and direction cannot be known
simultaneously
 Electron as cloud of negative charge
Modern Model of the Atom
The electron cloud
 Sometimes called the wave
model
 Electron as cloud of
negative charge
 Spherical cloud of varying
density
 Varying density shows
where an electron is more
or less likely to be
Quantum Theory
 Treats electron’s location as wave property
 Defined by quantum numbers
 Quantum numbers
 Provide information about size, shape, and orientation of
atomic orbitals
 Define atomic orbitals from general to specific
Principal Quantum
Number (n)
 Determines orbital size and electron energy
 Same as “n” value/orbital in Bohr model
 Positive whole number, NOT 0
 Shells – orbitals with same value
 n = 1, 2, 3, 4, etc.
Orbital Angular
Momentum Quantum
Number (l)
 Defines orbital shape for a particular region of atom
 Think as “subshell”
 l = n-1
 # of orbitals/subshells = principal quantum #
l
Orbital/Subshell
0
s
1
p
2
d
3
f
Magnetic Quantum
Number (ml)
 Describes orbital orientation within an atom
 Range from –l to +l, 0 is possible
 ml = 0, ± 1, ± 2, etc.
 ml = 2l + 1 (number of orientations)
How do you specify
orbitals?
2p
4f
Orbital Shapes
 s orbital
 1 possible orbital orientation, spherical shape
 n value determines size
 Charge cloud found near center, likely electron location
 p orbital
 3 possible orbital orientations, dumbbell shape
 pX, py, pz
What does atomic structure
REALLY look like?
P
Electron Spin
 Describes the motion of an electron, spinning
 As electron moves, magnetic field induced
 Electrons with opposite spins, cancel magnetic field of other
 Values: +1/2, -1/2
Homework
 Read lab procedure
 Read pp. 267-289 (for Friday)