Download Atomic Theory Notes- Chapters 5 and 13

9/17/14
By convention there is color,
By convention sweetness,
By convention bitterness,
But in reality there are atoms and space.
-Democritus (c. 400 BCE)
Development of Early Atomic
Theory
Democritus
Through
Rutherford
Democritus (460-370 BC)
•  Greek Philosopher
•  Pondered the
composition of matter
•  Matter is made up of
tiny indivisible
particles called
________________
Law of Conservation of Mass
•  Mass is neither created or destroyed in chemical
reactions. The atoms are _________________.
Hg(NO3)2 + 2KI → HgI2 + 2KNO3
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Which diagram best represents the
LOCM?
Heat
Heat
Law of Definite Proportions
•  Formulated by Joseph Proust
•  Different samples of a pure
substance always contain the
________________ proportion
of elements by ____________
•  Water is 11% H and 89% O
regardless of its source.
•  Elements combine in
___________________
proportions, not in random
proportions.
Dalton’s Atomic Theory
•  Elements are made of
______________________________________.
•  Atoms of the same element are the ______________
and atoms of different elements are
______________________.
•  Atoms combine in simple
____________________________________________.
•  Atoms are not
_____________________________________________
in a chemical reaction. They are just rearranged.
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Radioactivity
•  Becquerel, 1896
•  Alpha-α, ________________, heaviest, lowest
energy,positive charge,
_________________________________________
•  Beta-β, ___________________, light weight, higher
energy, negative charge, _____________________
•  Gamma-γ, ___________________, highest energy,
_________________________________________
http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/
animations_center.html#
J.J. Thomson and Cathode Ray
Tubes-1897
• 
Objects cast shadows.
• 
Rays can be formed with different metals.
• 
The rays exert a ___________________.
• 
Rays are deflected by ___________________
and magnetic fields.
• 
J.J. Thomson was the first person to deflect
the rays.
• 
Called the particles corpuscles.
• 
The term ______________________ coined
in 1891 by G. Johnstone Stoney.
http://highered.mcgraw-hill.com/sites/
0072512644/student_view0/chapter2/
animations_center.html#
Paddle
Wheel
Maltese
Cross
Simulation
Plum Pudding Model
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The Oil Drop Experiment
•  ____________________
(1868-1953)
•  Suspended a charged oil
drop between two plates
•  The _________________
of the electron was
determined to be 1.61
x10-19 C.
•  The mass of the electronhttp://highered.mcgraw-hill.com/sites/0072512644/
student_view0/chapter2/animations_center.html#
explanation
simulation
Rutherford Gold Foil Experiment
http://cameron4gfs.glogster.com/Gold-foil-experiment/
Summary
•  Experiment performed by
Ernest
____________________
•  Most of the particles passed
straight through as expected—
atoms are mostly
________________________!
•  However, some were
_________________ at angles
and some bounced back—
atom contains small, dense,
positively charged
___________________
•  Called the planetary or
________________ model
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The Atom
Three Major Subatomic Particles
Nuclear Symbols
Isotopes
Average Atomic Mass
Three Major Subatomic Particles
Particle
Symbol
Relative Charge
Mass
Number
Rel.
Mass
Act.
Mass
Electron
0.0005u
9.11•10-28 g
Proton
1.007 u
1.637•10-24 g
Neutron
1.009 u
1.675•10-24 g
1 u = 1.66•10-24 g
Carbon has 6 p+ and 6no in its nucleus.
By international agreement, 1 amu is 1/12 the mass of a carbon atom.
Symbols on the Periodic Table
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Isotopes and Mass Number
• 
Isotopes
- 
Atoms of the _________ ___________that have
different masses.
- 
Mass difference is due to the number of
______________ in the nucleus.
- 
To identify an isotope, you must know the
______________ of the element and the
_________ of the isotope
- A _________ is an individual isotope.
Mass Number, A
•  The total number of protons and neutrons in the
nucleus of an atom.
Mass Number, A = _______________+_______________
•  Uranium that has 92 protons and 143 neutrons.
•  Uranium that has 92 protons and 146 neutrons.
Representing Isotopes
•  Individual isotopes can be represented using
_____________ ____________ or the name of
the element followed by the mass number of the
isotope.
14
6
C
Carbon-14
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Nuclear Symbols
•  A =
•  Z =
A
Z
X Charge
•  Charge =
•  Positive
•  Neutral
•  Negative
Examples # 1 – 4
•  Write the nuclear
symbol for an atom
that has 15 protons,
16 neutrons and 15
electrons.
•  Write the nuclear
symbol for an element
that has 30 protons,
28 electrons and a
mass number of 65.
•  Calculate the number
of protons, neutrons,
and electrons
represented by the
nuclear symbol on the
27
3+
right.
13
•  An element has an
atomic number of 12,
13 neutrons in the
nucleus, and 10
electrons. Write the
nuclear symbol for
this ion.
Al
Isotopes of Hydrogen
MCMURRAY, J., & Fay, R. C. (2001). Isotopes of Hydrogen. In Chemistry (3rd ed.,
p. 46). Upper Saddle River, NJ: Prentice Hall.
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Average Atomic Mass
•  The average atomic mass is a
____________average of the masses of
all of the isotopes for a given element.
•  Average atomic mass is determined using
the following formula:
A.A.M. = (Mass I #1)(Fractional Abundance #1)+(Mass I#2)(Fractional
Abundance#2)+…
Isotope Problem #1
•  Calculate the average atomic mass for Ne.
Isotope
Atomic Mass
Percent
Abundance
Ne-20
20.00 amu
90.92
Ne-21
21.00 amu
0.26
Ne-22
22.00 amu
8.82
Fractional
Abundance
Isotope Problem #2
•  Natural chlorine is a mixture of isotopes. Determine its
atomic weight if 75.53% of the naturally occurring
element is chlorine 35, and 24.47% is chlorine 37.
Fogiel, M., Dr. (Ed.). (1996). Problem 708. In REA's Problem Solvers Chemistry
(p. 672). Piscataway, NJ: Research and Education Association.
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Isotope Problem #3
•  Only two isotopes of copper occur
naturally, copper-63 and copper-65. Their
masses are 62.9298 amu and 64.9278
amu respectively. If the average atomic
mass for copper is 63.546 amu calculate
the relative abundance of each isotope.
The Periodic Table
Periods / Series
•  ______ on the periodic table
• 
Represent the shell that the valence electrons occupy.
•  __________Electrons
–  ____________ electrons in an atom.
–  The electrons that are ________, _________, or
________ in chemical reactions.
–  The _____________ of valence electrons in an
element's outer shell determines the
______________ properties of that element.
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Groups/Families
•  ____________
•  All elements in the same group have the
same number of _________ electrons, and
therefore have similar
__________properties.
•  Some groups have special names
Special Groups
• 
• 
• 
• 
Group 1 Group 2 Group 7 Group 8 -
• 
___________
Elements
• 
In addition some
regions of the
periodic table have
special names.
Metals
•  Located on the _____
side of the periodic
table.
•  ______ valence
electrons in chemical
reactions.
•  ________,________,
________,________
•  Transition metals.
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Nonmetals
•  Located on the ______
side of the periodic table.
•  _______,__________
•  ______ electrons in
chemical reactions.
•  __________ ______:
Every atom wants ___
_________electrons.
Metalloids
•  Located on the
_________-_____
line.
•  Have ______ and
_______properties.
Electronic Structure of Atoms
II
Chapter 13
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Wave nature of light
•  ___________________________________—
carries energy through space
–  Examples: visible light, x-rays, radio waves
–  Speed of light in a vacuum c = 3.00x108 m/s
•  ____________________ (λ)—distance between
two adjacent crests or troughs
•  ____________________ (ν)—number of
complete wavelength cycles that pass a given
point each second (units: Hz or s-1)
Wave nature of light
•  Wavelength and frequency vary inversely
–  νλ = c
•  high frequency, short wavelength
•  low frequency, longer wavelength
What is the
amplitude of a
wave?
http://www.qrg.northwestern.edu/projects/vss/docs/Communications/1-what-is-wavelength.html
Electromagnetic spectrum
Unit/symbol
Length (m)
Type of radiation
Angstrom (Å)
10-10
X-ray
nanometer (nm)
10-9
Ultraviolet; visible
micrometer (µm)
10-6
Infrared
Other:
mm—infrared
cm—microwave
m—TV/radio
http://www.antonine-education.co.uk/physics_gcse/Unit_1/Topic_5/topic_5_what_are_the_uses_and_ha.htm
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http://micro.magnet.fsu.edu/primer/java/wavebasics/index.html
Example problems
•  Yellow light from a sodium vapor lamp has a
wavelength of 589 nm. What is the frequency of
this light?
–  A:
•  An FM radio station broadcasts electromagnetic
radiation at a frequency of 103.4 MHz. Calculate
the wavelength of this radiation
–  A:
Relationship between Energy and
Frequency
•  E = hν
•  h = Planck’s Constant, ____________ J*s
•  Energy and frequency are
______________________ related.
•  Energy and wavelength are
______________________ related.
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Example problems
•  Calculate the energy electromagnetic radiation
that has a frequency of 2.44 x 1015 Hz.
–  A:
•  Calculate the frequency of electromagnetic
radiation that has an energy of 2.99x10-19J.
–  A:
•  Calculate the energy of electromagnetic
radiation that has a wavelength of 750 nm.
–  A:
So light is a wave!
•  Light Diffracts
•  Light Refracts
•  Light has a
wavelength and
a frequency
http://lukewest.edublogs.org/
The wave model of light can’t
explain everything...
I. Blackbody radiation—emission of light from hot
objects)
II. Photoelectric effect—electron emission from
metals when excited by light
III. Emission spectra of light for excited gas atoms
•  Let’s look at these more closely...
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I. Blackbody radiation
•  If you heat a solid, it
emits radiation
–  Stove burner; light bulb
•  When metals are heated
they change color from
red to yellow and from
yellow to white.
•  As the wavelength
decreases, the intensity
increases.
I. Blackbody radiation
I. Blackbody radiation
•  In 1900, German physicist Max Planck
determined that energy was only released/
absorbed in discrete chunks
–  He called the smallest quantity of energy a quantum
•  Energy equals a constant times the frequency
–  E=hν
h = 6.626x10-34 J*s
–  Can only have E in multiples of hν—2hν, 3hν, 4hν, etc
–  Energy is _________________________—can only
have certain values!
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Quantized Energy
•  ___________________the smallest amount of
energy that can be
emitted or absorbed as
electromagnetic radiation.
Plural-______________.
•  An elevator
•  A ramp vs. stairs
•  So why does energy
seem continuous to us?
II. Photoelectric effect
•  1905—Einstein realized that when light struck a metal,
the metal emitted electrons
•  __________________ of the light was important
–  minimum needed to get emission
–  Increasing beyond this minimal affect
•  Einstein theorized
that light was behaving
more like an energy
“packet” than a wave
•  Called this “packet”
a _______________,
which behaves
like a tiny particle
http://www.cem.msu.edu/~harrison/cem483/
Example problem
•  A laser emits light with a frequency of 4.69x1014
s-1. What is the energy of one photon?
–  A:
•  If the laser pulse has 5.0x1017 photons, what is
the total energy?
–  A:
•  If the laser emits 1.3x10-2 J of energy during a
pulse, how many photons are emitted during the
pulse?
–  A:
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Behavior of light
•  Light behaves like a wave (c=νλ)
•  Energy of light depends on frequency
(E=hν)
•  Light behaves like a particle (photoelectric
effect)
•  So is light a particle or a wave?
–  Both!!! Wave-particle duality
III. Spectra
•  _______________ light—light with
only one wavelength
•  ________________—light separated
into different wavelength components
–  ____________________ spectrum—
contains light of all wavelengths
–  ________________ spectrum—contains
light of a few distinct wavelengths
http://www.upscale.utoronto.ca/IYearLab/Intros/Spectra/Spectra.html
Emission Spectrum
•  Atoms ______________
radiation when they are
excited.
•  Light given off consists of
discrete wavelengths
called the line or emission
spectrum.
•  What is special about an
element’s line spectrum?
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Atomic Absorption Spectrum
•  Atoms will also ________
discrete wavelengths of
light.
•  The absorptions are
observed as
_________________ in a
continuous spectrum.
•  How does an element’s
emission spectrum relate
to its absorption
spectrum?
Hydrogen spectrum
•  Mid-1800’s, only 4 distinct lines seen in the
hydrogen gas spectrum
http://faculty.sdmiramar.edu/fgarces/LabMatters/Instruments/AA/AAS_Theory/AtomicLineOrigins.htm
http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.htm
•  1885 Johann Balmer developed an equation for
the wavelengths of these 4 lines
Bohr’s model: three postulates
•  Only orbits of certain radii, corresponding
to certain _____________, are permitted.
http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm
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Bohr’s model: three postulates
•  Electron in permitted orbit has specific E,
and is “allowed”. It does not radiate E,
and won’t spiral into the nucleus.
http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm
Bohr’s model: three postulates
•  E is emitted or absorbed by electron only if
electron changes from one “allowed” orbit
to another. E is emitted or absorbed as a
photon.
Lowest energy, n=1
is called the
________________
When electron is in a
higher energy orbit (n=2
or higher), it is in an
___________________
http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm
Principal quantum number
•  Bohr calculated energy for each orbit:
E = (−hcRH )(
1
1
) = (−2.18 x10 −18 J )( 2 )
n2
n
•  h = Planck’s constant, c = speed of light,
RH= Rydberg constant
•  The integer n is called the
_________________________________
What happens as n becomes infinitely large?
State when electron is removed is called the _____________________
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Changing energy
•  Electrons jumping from n to n change E
•  Calculate the change in energy of an
electron
ΔE = E f − Ei = E photon = hν
•  Recall that c=νλ and the Rydberg
equation.....
1
1
ΔE = (−2.18 x10 −18 J )(
n 2f
−
ni2
)
This explains line spectra...
If nf is smaller than ni,
the electron moves
__________ to the
nucleus and E is
released
•  Visible hydrogen line
spectrum comes from
transition from excited
states to the _______
level
http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.htm
So why not Bohr 4-eva?
UNFORTUNATELY...this model ONLY
works for hydrogen...
BUT...we get two important ideas
•  Electrons exist ONLY in discrete energy
levels described by __________________
•  ________________ is involved in moving
electrons between levels
•  Where to next?
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Wave behavior of matter
•  Louis de Broglie wrote a (very short) PhD
thesis in 1925
•  Proposed that if radiation (light) is a wave
that can act like a particle, matter can act
like a wave
–  Wave-particle duality
–  Equation for the de Broglie wavelength of
matter:
λ=
h
mv
m = mass
v = velocity
Example problems
•  A major league pitcher can throw a baseball 90
miles per hour. If the average baseball weighs
145 g, what is the de Broglie wavelength of the
baseball?
–  A:
•  What is the de Broglie wavelength of an electron
moving at a speed of 5.97x106 m/s? The mass
of an electron is 9.11x10-28 g.
–  A:
New model for the atom!
•  Electron is a particle
with some wave-like
properties (de Broglie)
•  Schrodinger proposed a
model that abandoned
the “electron in an orbit”
notion.
•  New model was based
on the wavelike
properties of the
electron.
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Heisenberg Uncertainty Principle
(HUP)
•  German physicist Werner
Heisenberg determined
that the dual nature
(wave & particle) limits
how precisely the
momentum and
location of a particle can
be known
ΔxΔ(mv) ≥
h
4π
Change in position times change in momentum (mass times velocity)
Heisenberg Uncertainty Principle
(HUP)
•  Typically only comes into affect for small
masses (ie, the position or momentum of
an electron)
•  We know me = 9.11x10-31 kg and the
average speed is 5x106 m/s. If the
uncertainty of the speed is 1%, what is the
uncertainty in the position of the electron?
–  A: 1x10-9 m
•  But the diameter of one H atom is 1x10-10!
Let’s try the calculation for a baseball (mass 145g, speed 90 mph) A:9.1x10-34 m
Quantum mechanics
•  QM is a new way of describing the behavior of
subatomic particles (mostly electrons)
•  1926, Austrian physicist Erwin Schrödinger
developed an equation (the wave function or Ψ)
to describe both the wave and particle properties
of the electron
–  Solving this equation is complicated and requires
advanced calculus...so we’ll skip that step. We will,
however, look at the results!
•  The square of the equation Ψ2 (electron density)
tells us the probability of where the electron will
be
–  Remember, we can’t know for certain because of HUP
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Orbitals and Orbits
•  So Ψ gives us an idea of where the
electron is allowed to be in space—called
an _________________
•  Note that QM orbital _____________ the
same as a Bohr orbit
http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm
https://www2.bwdsb.on.ca/
~f_schlenker/4U/4U
%20quantum%20chemistry/
university%20website/
imgres_files/a.htm
http://www.sparknotes.com/chemistry/
organic1/atomicstructure/section1.rhtml
Quantum numbers
•  Describe an orbital & where electrons are
•  Bohr’s model gave us the primary
quantum number, n
•  QM has ___________ quantum numbers,
_______
________
_______
________
•  Let’s take a closer look at these
numbers...
First quantum number, n
•  ____________________ quantum
number: positive integer (1, 2, 3, ...)
–  Also defined as the ____________________
–  As n increases, the size of the orbital
increases and the electron spends more time
away from the nucleus
–  Higher n means higher E
–  Corresponds to periods on PT
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Sizes of orbitals
•  As n ______________,
so does the size of the
orbital
•  As orbital size increases,
the probability of the
electron being found near
the nucleus decreases
http://chemistry.umeche.maine.edu/~amar/fall2007/orbitals.html
http://physchem.ox.ac.uk/~hill/tutorials/qm1_tutorial/atomorb/index.html
Second quantum number, l
•  _____________________ quantum
number : integral values from 0, 1, ..., n-1
–  Represents the _______________ or
__________________ of the electron shell
–  Defines the _____________ of the orbital
–  Also have ______________ designations
Value of l
0
1
2
3
Letter used
Azimuthal Quantum Number, l
•  l =0
s-sublevel - with
spherical orbital
•  l =1
p-sublevel - with dumbbell
orbitals
•  l =2
d-sublevel - with
cloverleaf orbitals
•  l =3
f-sublevel - with eightlobed orbitals
•  Corresponds to the
blocks on the periodic
table.
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Shapes of orbitals
•  Second quantum number, l,
tells us the shape of the
orbital
•  l = 0 (also called s) is a
________________ orbital
–  There is __________________
l = 0 orbital in each electron
shell (for each n)
http://www.sparknotes.com/chemistry/fundamentals/atomicstructure/section1.rhtml
Shapes of orbitals
•  l = 1 (or p) orbitals are bow-tie shaped
–  Remember for l = 1 there are _____ orbitals
–  Same shape, oriented differently in space…
Which
quantum
number is
different
with these
orbitals?
http://chemistry.umeche.maine.edu/~amar/fall2007/orbitals.html
Shapes of orbitals
•  l = 2 (or d) orbitals are complicated
–  4 are clover leaf shaped
–  Remember
l = 2 there are
____ orbitals
l = 3 (or f) are waaaaaaay
too complicated to worry
about the shape…just
know that they’re there
and how many there are!
http://www.chem.ufl.edu/~itl/2045_s00/lectures/lec_10.html
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f orbital shapes
http://www.chem.tamu.edu/rgroup/hughbanks/courses/673/handouts/handouts.html
Third quantum number, ml
•  _________________
quantum number ml : have
integral values between -l
and l, including zero
–  Represents the ____________
in the subshell
–  Describes the orientation in
space
–  The total number of orbitals in
all sublevels of a shell is _____
The Fourth Quantum Number, ms
•  The line spectra of many
electron atoms show each line
as a closely spaced pair of
lines.
•  Stern and Gerlach experiment
placed a beam of atoms
through a slit into a magnetic
field.
•  Two spots were found: one
with electrons spinning in one
direction and one with
electrons spinning in the other
direction.
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________ Quantum Number, ms
•  Electrons are free to
spin clockwise or
counterclockwise
about their axis.
•  ms = ±1/2
•  ms is independent of
the other three
quantum numbers.
Pauli Exclusion Principle
•  PEP: No two electrons can have the exact
same quantum numbers
•  Since electrons can only spin two ways,
(CW or CCW), then two electrons will fit in
one orbital and will have ______________
spins
–  Designate one as +1/2 and the other as -1/2
Quantum Model in a Nutshell
•  Principal quantum number, n, gives the electrons
average distance from the nucleus and the energy level.
•  Azimuthal quantum number, l, describes the sublevels.
One sublevel for each value.
•  Magnetic quantum number, ml , describes the
orientation of the orbital in space. One orbital for each
possible value.
•  Spin quantum number, ms, describes the spin of the
electron in a magnetic field.
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Interpreting Quantum Numbers
2, 1, -1, +1/2
Quantum
Number
Principle
Symbol
Describes
Possible
Values
n
Energy
Level
Whole #s ≥ 1
Azimuthal
l
Sublevel
Whole #s
0 to n-1
Magnetic
ml
Orbital
Spin
ms
Spin
- l to +l
±1/2
Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP
Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning
Center, 2008.
Determining Quantum Numbers
5d5
Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP
Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning
Center, 2008.
Examples
•  Ex: Give the quantum numbers for the
selected arrow.
•  Ex: Give the quantum numbers for the
selected arrow.
Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP
Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning
Center, 2008.
28
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Incorrect Quantum Numbers
•  Identify what is wrong with the following
sets of quantum numbers.
1, 1, 0, 1/2
2, 1, -2, -1/2
3, 2, -1, 1
2, 1, 1, 1/2
Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP
Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning
Center, 2008.
What order do we fill the orbitals?
Want to have ________________possible energy!
•  As n increases, E _____________ (E1s<E2s<E3s)
•  As l increases, E ______________ SLIGHTLY
(E3s<E3p<E3d)
–  Orbitals within the same subshell l (example, px, py,
and pz) have the same energy and are called
______________________
•  Fill all degenerate orbitals in a subshell before
moving on to the next subshell
Order of filling
•  Electrons are arranged based on the energy of
each orbital.
•  The order of filling from lowest to highest E is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
•  Notice that the d-orbitals fill ____ energy level
later than expected and that f-orbitals fill ____
energy levels later than expected
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Energy Levels in Single and Multi-electron Atoms
Electron Configurations
•  Tell us how the electrons are distributed among
the various orbitals of an atom.
»  Li – 1s22s1
»  Na – 1s22s22p63s1
»  K – 1s22s22p63s23p64s1
»  He - 1s2
»  Ne - 1s22s22p6
»  Ar - 1s22s22p63s23p6
•  The most stable configuration or ground state is
that in which the electrons are in the lowest
possible energy state.
Rules for Writing Electron Configurations
•  Electrons fill orbitals of increasing energy with no more
than two electrons per orbital. (___________________)
•  For degenerate orbitals, electrons fill each orbital singly
before any orbital gets a second electron (__________).
–  Another way to think of this: in degenerate orbitals, electrons
prefer to be unpaired.
•  No two electrons can fill each orbital with the same spin
(________________).
•  How do we show spin?
–  An arrow pointing upward has ms = +1/2 (spin up).
–  An arrow pointing downward has ms = -1/2 (spin down).
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Writing electron configurations
•  Determine the number of electrons.
•  Fill orbitals in order of increasing energy
–  s - ____e- maximum
–  p – ____e- maximum
–  d – ____e- maximum
–  f – ____e- maximum
•  d-orbitals fill 1 level later, f-orbitals fill 2 levels
later
•  Continue filling until all electrons are used.
Electron configuration and the periodic table
Writing Configurations Using the
Periodic Table
•  Trace the path you
take to arrive at the
element.
•  d-sublevels fill one
period behind
•  f-sublevels fill two
periods behind
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Examples
•  O
•  Mg
•  Ga
•  Ru
•  V
Configurations Within a Group
•  Li
•  Na
•  K
•  F
•  Cl
•  Br
•  Core electrons vs. valence electrons
•  Elements in the same group have the same outer valence
shell configuration.
Exceptions to Aufbau Ordering
•  Several elements do not follow the Aufbau order of filling.
•  _________ and __________ also follow copper pattern.
•  ________ also follows Cr pattern.
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Exceptions Continued
•  Because of the stability of certain
configurations.
•  Full sublevel > Half-filled > No
particular organization
•  A half-filled d sublevel is more stable than
a full s sublevel.
Abbreviated or Noble Gas
Configurations
•  Way to describe the core
electrons and only show
valence electrons
•  Cl
•  Write the symbol for the
Noble gas that precedes
the element.
•  Trace the path to the
element
•  Sn
More Examples of Abbreviated
Configurations
•  Ca
•  S
•  Rb
•  Ne
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Orbital Configurations
•  Orbital diagrams represent the detail lost by E.C.
•  Each orbital within a sublevel is represented by
a box, line, or circle.
•  Electrons are represented by half-arrows.
•  Used primarily for valence electrons, since all
inner core electrons will be paired.
Hund’s Rule
•  Within a given sublevel,
each orbital must contain
one electron before they
are allowed to pair up.
•  Orbitals in the same
sublevel are said to be
degenerate.
Examples of orbital configurations
•  O
•  Fe
•  K
•  P
•  Ar
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Orbital Configuration Questions
• 
From the orbital diagrams select an
example which demonstrates
–  a violation of Hund's rule –  a violation of the Pauli exclusion
principle –  a ground state orbital diagram –  an excited state orbital diagram
–  a violation of the Aufbau
principle Lewis Dot Diagrams
•  Introduced by Gilbert N.
Lewis in 1916
•  Use to represent
_________________
electrons and bonding
between atoms
•  Used with the
representative elements
(s- and p-block)
Whitley, Kelley. "The Periodic Table." ChemProfessor. 1 Oct. 2008. 30 Nov. 2008
<http://www.chemprofessor.com/aboutme.htm>.
Dot Diagram Examples
•  Lithium
–  Config:
–  Dot Diagram:
•  Aluminum
–  Config:
–  Dot Diagram:
•  Arsenic
–  Config:
–  Dot Diagram:
Whitley, Kelley. "The Periodic Table." ChemProfessor. 1 Oct. 2008. 30 Nov. 2008
<http://www.chemprofessor.com/aboutme.htm>.
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Configuration Comparison
Pro
Con
Standard
Quicker than orbital
Does not show pairing
Abbreviated
(Noble Gas)
Shows only valence e- No core eQuick
No pairing
Orbital
Shows pairing of
electrons
Longest
Dot
Shows only valence e- No core eVery quick
Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP
Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning
Center, 2008.
Review Questions
•  Se: 1s22s22p63s23p64s23d104p4 1.  Atom that has 2
unpaired electrons.
•  Mg: 1s22s22p63s2
•  Al: 1s22s22p63s23p1
•  N: 1s22s22p3
•  F: 1s22s22p5
2.  Atom with exactly 2
electrons in the
highest sublevel.
3.  The atom that must
obtain 3 electrons to
fill a sublevel.
Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP
Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning
Center, 2008.
Review Questions
A. 
1s22s22p63s23p64s23d104p4
B.  1s22s22p63s23p2
C. 
1s22s22p63s23p24s1
D.  1s22s22p63s23p64s23d8
E. 
1s22s22p63s23p64s23d104p6
F. 
1s22s22p6
•  Which atom is in an
excited state?
•  Which configuration
represents the element
Ni?
•  Which configuration(s)
represents a noble gas?
•  Which configuration
represents a sodium ion?
Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP
Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning
Center, 2008.
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The End
For now…
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