Chapter 5
• Light
• Electrons in atoms
Models of the atom
Rutherford’s model of the atom did not show or
explain chemical properties of elements
Needed some modifications
Along comes Bohr….
THE BOHR MODEL
A Danish physicist, Niels Bohr, proposed that an electron is
found only in specific circular paths, or orbits, around the
nucleus
 This is a combination of Rutherford’s model and
Einstein’s famous E=mc2
Each orbital has a fixed energy, known as energy levels
 These levels are similar to rungs on a ladder. The lowest
rung has the least amount of energy. The highest rung
has the most energy
A quantum of energy is the amount of energy
needed to move an electron from one orbital to
another
 When electrons jump from one orbital to another,
they release light of different frequencies
 These different frequencies result in different colors
of light
Light behaves two ways:
1. As a wave
2. As a particle
LIGHT: What Is It?
• Light Energy
•
Atoms
• As atoms absorb energy, electrons jump out
to a higher energy level.
• Electrons release photons (bundles of
energy released) when falling down to
the lower energy level.
Light Waves
ν
λ
crest
trough
• Light pieces are called = PHOTONS
• Light travels as a wave and has a QUANTUM of energy
Wavelength (λ)= distance from crest to crest, meters
Frequency (ν or f)= # of waves that pass a point in 1
sec., Hertz or Hz = 1/sec
Speed of Light (c)= 3.00 x 108 m/sec
c=λf
ELECTROMAGNETIC
SPECTRUM
Light as a Particle*
• photoelectric effect
•
•
Emission of electrons from a metal when
light is shined onto the metal.
This helped us understand it also acts as a
particle!
Light Energy
Quantum = minimum amount of Energy (in
joules) that can be lost or gained by an
atom
E=hf
Planck’s constant = h = 6.626 x 10-34 J sec
E=energy (J); f=frequency (1/sec, or Hz)
Light Spectra
• Ground State – lowest energy state of an atom
• Excited State – an atom in a higher energy state
• Line Emission Spectrum –is the spectrum of frequencies
of electromagnetic radiation emitted by atoms or
molecules when they are returned to their ground state.
•
Each element's emission spectrum is unique.
Line Emission Spectrum
Quantum Model Notes
• We don’t know exactly where the electrons
are, but we can have a MODEL showing
the most likely place for an electron.
Quantum Model- 3D model to
identify where electrons exist.
ENERGY LEVELS=
•
•
•
a possible location for electrons found
around the nucleus
There are seven (maybe more!)
energy levels.
• SUBLEVEL=
•
•
For each level, (EXCEPT LEVEL #1!)
there can be orbitals of different
shapes, called sublevels.
We name these sublevels s,p,d,f
Long Periodic Table
Orbitals
•
•
•
a single, allowed location for the
atomic electron.
Each orbital can hold at most 2
electrons
can be empty or only half-full
Sublevel
# Orbitals
s
1
Total # of e(# of
orbitals*2)
2
p
3
6
d
5
10
f
7
14
3. What does the quantum mechanical model determine about electrons
in atoms?
4. How do 2 sublevels of the same principal energy level differ from
each other?
5. How can electrons move from one energy level to another?
6. How many orbitals are in the following sublevels?
1. 3p
2. 3d
3. 2s
4. 4f
5. 4p
Aufbau Principle:
Electrons fill the lower energy levels
first
Example: Climbing a staircase. You
have to start at the bottom and
move your way up
Pauli Exclusion Principle:
Each orbital can hold only 2 electrons
of opposite spin
Example: Shoes in shoe box
QUESTION FOR DISCUSSION
Modeling the Pauli Exclusion Principle:
How are 2 bar magnets similar to electrons’
spins in the same orbitals?
Rules for electron
Configurations
• Hund’s Rule:
• Electrons fill orbitals of equal energy
before sharing orbitals.
• Ex: In a house, usually all kids
bedrooms are filled with one kid
before their parents make them
share a room.
The Energy Levels
of an atom
n=4
n=3
Energy
n=2
n=1
nucleus
4s
4p
4d
3s
3p
3d
2s
2p
4f
1s
Sublevels = s,p,d,f
orbitals =