Electrons and their
Excited State
Movement
Principal Quantum
Vocabulary
Number
Electromagnetic Radiation Principal Energy Level
Electromagnetic Spectrum Energy sublevel
Wavelength
Electron Configuration
Frequency
Aufbau Principle
Quantum
Pauli Exclusion Principle
Photon
Hund’s rule
Atomic emission spectrum Valence electrons
Ground State
Electron-dot Structure
Objectives:
1. Define a quantum of energy, and explain how it is related
to an energy change of matter.
2. Compare the Bohr and quantum mechanical models of
the atom.
3. Use the quantum numbers to describe the electrons in an
atom.
4. Apply the Pauli exclusion principle, the aufbau principle,
and Hund's rule to write electron configurations using
orbital diagrams and electron configuration notation.
5. Identify the relationships among a hydrogen atom's
energy levels, sublevels, and atomic orbitals.
6. Compare the wave and particle natures of light.
7. Explain the impact of de Broglie's wave article
duality and the Heisenberg uncertainty principle
on the current view of electrons in atoms.
Review:
What does a Bohr Model include?
How does this differ from Rutherford’s Model?
What do we know about the Quantum Mechanical
Model?
Rutherford
• Discovered the nucleus using the Gold Foil Exp.
• Rutherford’s model demonstrates his findings
– Mostly empty space containing e-’s
– At the center is a positive sphere (nucleus)
– NOTE: there aren’t e- orbitals or neutrons until
further models
Bohr Model
• An atomic physicist, Neils
Bohr developed a
planetary model based on
Rutherford’s findings
– Included a nucleus with
neutrons
– Electron’s moved in
definite orbits (paths)
around nucleus
• The e-’s have fixed
energies and do not
lose energy as they
orbit the nucleus
Quantum Mechanical Model
(our current model)
• The electrons no longer occur in orbits
– In this model, we have an idea of where the e-’s could
be, which is our e- cloud
– The electron cloud is made up of different orbitals of
different shapes
• The orbitals are not paths but rather probability densities
Energy Levels
Diagram: (Fig. 13.3)
• The energy levels in an atom are
sort of like rungs (steps) of a ladder.
• The more energy an electron has,
the farther away from the nucleus it
usually will be.
• The energy levels are not evenly
spaced. They get closer together as
you travel farther away.
• To move from one “rung” to
another requires a “quantum” of
energy.
Quantum Numbers
•
•
•
Describe the location of the e-’s around the nucleus.
Quantum #’s are sort of like a home address for the electron.
This information about the location of the e-’s in an atom can be
used to:
(1) determine chemical & physical properties for the
elements.
(2) show how the Periodic Table is organized.
(3) show how and why elements combine to form
compounds.
The four quantum numbers are abbreviated: n, l, ml, ms
The Principal Quantum Number: n
n, the principal quantum number
a) determines the overall energy of the atomic orbital
b)Tells shell number
c) May have any positive integral value from 1 to infinity
example: 1, 2, 3, 4…etc.
d) Tells the average distance to electron is from the nucleus
-as n increases, the distance between the outermost electron and
nucleus increases
e) The distances can be called
principal energy levels
1
2
3
nucleus
Angular momentum quantum number: l
(a.k.a orbital quantum number)
l: tells us which one of the orbitals in the n shell we
are looking at.
All of these orbital shapes are based on the probability
of finding the electron in the cloud.
a) l describes the shape of the electron’s momentum
around the nucleus with a letter: (s, p, d, & f) These
are sometimes called “ sublevels ”.
s= spherical cloud
p= ellipsoid
d & f orbital shapes are complex criscrossed ellipsoids, and
some d’s and f’s are an ellipsoid with a doughnut or two
around the middle.
s - orbital
p - orbitals
d - orbitals
f - orbitals
Magnetic Quantum Number: ml
Ml : tells us the orientation of the orbital in space
Each orbital has a specific number of orientations it can occur in:
s= 1 orientation
p= 3 orientations... (x, y, and z)
d= 5 orientations
f= 7 orientations
The orientations can be represented with a line or a box.
Examples: ___ This means a spherical orbital at a distance of
1s
“1” (the number 1 is the n) to the nucleus.
This orbital is centered about the x, y, and z
axis.
□□□
4p
This represents an ellipsoid orbital with its
3 possible orientations at a distance of “4”
from the nucleus.
Spin Quantum Number: ms
ms: spin quantum number
a) there are two possible orientations (spins) for electrons
b) the two electrons may occupy the same orbital if they have
different spins
ms describes how the electron in an orientation is spinning
around the nucleus. This spin can be thought of as “up” or
“down”. (Some like to imagine it spinning “clockwise” and
“counterclockwise”.) The spin is represented as an arrow in
the direction of the spin.
Example: ↑ This represents one electron in a spherical
2s
orbital with spin “up” at a distance of “2”
from the nucleus.
Figure 11.31: Orbitals being filled for elements in various parts of the periodic table.
What do we do with the quantum numbers?
We said that quantum numbers are like an address
for an electron but how do we write the address?
Electron configurations are used to describe the location of
electrons in an element. They incorporate the four quantum
numbers and are specific to each neutral element.
Electrons in these configurations follow four rules:
1) The Aufbau Principle
2) The Pauli exclusion principle (contains 2 rules)
3) Hund’s rule
The Pauli Exclusion Principle
Rule #1: Only 2 electrons can fit into each orbital
Example:
↑↓
↑
↑↓↑
___
___
not ____
1s 2s
1s
Rule #2: Electrons in the same orientation have opposite
spins.
Example:
↑___
↓
1s
not ↑↑
___
1s
The Aufbau Principle
Rule #3: Electrons fill the lower energy orbitals first.
-This means that energy level 1 will fill before 2.
Examples:
1s would be filled before 2s
3s would fill before 4s
Hund’s Rule
□
Rule #4 : “Bus seat rule”---> Every “ ” in an
orbital shape gets an electron before any
orientation gets a second e-.
Example:
□□□
↑
↑
↑
2p
not
□□□
↑↓ ↑
2p
How to write an Electron Configuration
1) Find the element on the periodic table. (Ne)
2) List ALL of the prior orbital sublevels, l (w/energy
level, n)
Ne: 1s 2s 2p
3) Write the amount of electrons in each orbital
-s: max of 2 e
-p: max of 6 e
-d: max of 10 e
Ne: 1s22s22p6
Electron Configurations
Practice Problems:
Write the electron configuration notation for each of the following
atoms:
H
Li
F
Fe
Br
Kr
The Shorthand
Shorthand Method:
ex: Mg
1) Find the previous Noble Gas (last column on
the periodic table)
2) Put the chemical symbol in brackets
3) Write the remaining e- configuration
Shorthand Method contd.
Practice Problems
Li
F
Fe
Br
Kr
Electron Configurations
(Energy Level Diagram)
Electron Configurations can be
drawn in an ENERGY LEVEL
DIAGRAM
-this shows energry levels rising
as the n (principle Q#) increases
-this diagram can aide in the
order in which the energy levels
and orbitals are filled
Silicon: 1s2, 2s2, 2p6, 3s2, 3p2
AFTER Si, go to your notes
and practice
filling/drawing these for
our previous examples.
The Exceptions
Like all other rules, there are exceptions to 2 of the rules. They occur
because of stability.
THE FOLLOWING ELEMENTS DO THIS: Cr, Mo, Cu, Ag, Au
(ANY OTHERS I WILL NOT HOLD YOU ACCOUNTABLE FOR)
-Half-filled or completely filled d & f sublevels have lower energies
and are more stable than partially filled d’s and f’s.
-This means that an atom can “borrow” one of its “s” electrons
from the previous orbital to become more stable.
Example:
becomes
___
___ ___ ___ ___ ___
5s
4d
___
___ ___ ___ ___ ___
5s
4d
Because the 4d sublevel is now full, the atom is at a lower energy state
and therefore more stable.
Electron Configurations & Properties
How do electron configurations relate to the chemical and physical
properties of an element?
-All elements w/ the same outer shell e- configurations have similar
properties.
-This means that elements in the same column (or group) have
similar properties.
Examples: (1) Li, Na, K, Rb, and Cs all have 1 lone “s” e- for
their last orbital... ( 1s1, 2s1, 3s1…etc.) They all react with water
to produce hydrogen gas.
(2) Ne, Ar, Kr, Xe, and Rn all have the outer energy level
completely filled with electrons...(2s22p6, 3s23p6, 4s24p6, etc.)
This makes all of them inert. (inert means it doesn’t react)
More Practice Problems
Bromine
(1) Which element has its last electron as a 4p5? ___________
F, Cl, I, At
(2) Which elements are similar in properties as Bromine? __________
(4) Which electron is added after 6s2? ________
4f1
(5) Which element would “borrow” a 5s electron to get a half-filled
“d” sublevel? ___________
Mo
(6) What is the shape of the last orbital filled in Calcium, (Ca)? sphere
_____
4
(7) How many electrons are in the last “p-orbital” of Sulfur, (S)? ____
What is Electromagnetic Radiation?
Any wave of energy traveling at a speed of light is called
electromagnetic radiation.
-Waves are made of photons
-equivalent to a quantum of light
-Electromagnetic radiation can be broken down by frequency
and wavelength into seven types of radiation
-frequency is the number of repeating periods over a specified time. A
period is the duration of one wavelength
-Wavelength is the distance between the crests of a wave
Electromagnetic Radiation
Below is the electromagnetic spectrum. It shows the types of radiation from the
longest to the shortest wavelength.
-This order also demonstrates the lowest to highest energies.
Electromagnetic Radiation
(1) Radio Waves – longest wavelength, lowest energy: used in
communications
(2) Microwaves-- broadcasts TV signals and used to cook food.
(3) Infrared (IR) -- we feel this as heat; Snakes & owls can “see” this.
infrared image of heating pipes under a floor
Infrared
Vision
Electromagnetic Radiation
(4) Visible Light -- the only radiation we can detect with our eyes. It
can be separated into the colors of the spectrum with a prism
ROYGBIV
(5) Ultraviolet (UV) -- gives you a sunburn; bees can “see” this; some
of this radiation from the sun gets blocked by the ozone layer
flower photo under normal light
flower photo under UV light
Electromagnetic Radiation
(6) X-Rays: used for medical imaging
Electromagnetic Radiation
(7) Gamma Rays– In the universe, gamma rays are produced by neutron
stars, pulsars, supernova explosions and is found near black holes.
-ON EARTH- gamma rays are generated by nuclear explosions,
lightning and radioactive decay
UNFORTUNATELY…
YOU WILL NOT BECOME HULK FROM EXCESSIVE EXPOSURE!
Credit: NASA/DOE/Fermi LAT Collaboration, CXC/SAO/JPLCaltech/Steward/O. Krause et al., and NRAO/AUI
COSMIC RAYS
The last type of radiation is sometimes grouped with gamma
rays…
(8) Cosmic Rays– not true radiation; they are comprised of
charged particles. Almost all of this radiation from the sun is
blocked by the ozone layer and our magnetic field. Most is
made up of matter, a little is made of anti-matter.
How is Light Produced?
Light is produced by the movement of
electrons in between orbitals within an atom
- e- movement is caused
by atoms getting hit with
energy (heat/electricity)
-e- absorb the energy
jump energy levels
- the electron falling back
to its original state cause
a release of energy as
light or other forms of
EMR
How Light is Produced
-Each photon emitted has a specific frequency
-The color of the light that is given off depends on how
far the electron falls
-The farther the fall, the greater energy
All the Photons Produced by Hydrogen
How hydrogen
produces the four
visible photons
How Light is Produced
•
•
Electrons are located in certain energy levels (and orbitals) around
the nucleus, only certain specific colors of light are emitted.
Scientists use a spectroscope to separate these colors into bands of
light. These bands of color look like a bar code of color which is
characteristic of that element. No two elements produce the same
spectrum of colors. This can be used to distinguish one element from
another contained in a sample. (See Fig. 13.11)
Chemical Spectra
Chemists use machines called spectrometers to analyze substances.
-All matter absorbs and reflects electromagnetic radiation differently.
Our eyes see the visible light that was reflected by objects.
-Chemists can compare information, called spectra, to previously
recorded information to identify substances and predict molecular
structure..
Emission Spectrum
Hydrogen
Spectrum
Neon
Spectrum
Chemical Spectra (IR spectra)
caffeine
NMR Spectra
NMR is similar to an MRI but it is used to
see molecular structure
-the spectra below is for benzoin