1. General Chemistry Atomic Structure & the Periodic Table
Chemical Bonding
Introduction to Chemical Bonding
The Ionic Bond
The Covalent Bond
The Metallic Bond
Intermolecular Attractive Forces
Structure and Shape
Stoichiometry
Functional Group (Biology)
Class Notes
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2. ATOMIC STRUCTURE Definition of Chemistry:
The study of the properties, composition, and
STRUCTURE of matter, the physical and
chemical changes it undergoes, and the energy
liberated or absorbed during those changes.
The foundation for the STRUCTURE of inorganic
materials is found in the STRUCTURE of the atom.
Material Properties
Class Notes
Bulk Structure
Molecular Structure
Atomic Structure
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3. ATOMIC STRUCTURE Historical Development: Greek Concepts of Matter
Aristotle - Matter is continuous, infinitely
divisible, and is composed of only 4 elements:
Earth, Air, Fire, and Water
Won the philosophical/political battle.
Dominated Western Thought for Centuries.
Seemed very “logical”.
Was totally WRONG!!
Class Notes
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4. ATOMIC STRUCTURE The “Atomists” (Democritus, Lucippus,
Epicurus, et. al.) - Matter consists ultimately
of “indivisible” particles called “atomos” that
canNOT be further subdivided or simplified.
If these “atoms” had space between them,
nothing was in that space - the “void”.
Lost the philosophical/political battle.
Lost to Western Thought until 1417.
Incapable of being tested or verified.
Believed the “four elements” consisted of
“transmutable” atoms.
Was a far more accurate, though quite imperfect
“picture” of reality.
Class Notes
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5. ATOMIC STRUCTURE Modern Concepts of Matter
John Dalton (1803) - An atomist who formalized
the idea of the atom into a viable scientific theory
in order to explain a large amount of empirical
data that could not be explained otherwise.
Matter is composed of small “indivisible” particles
called “atoms”.
The atoms of each element are identical to each
other in mass but different from the atoms of other
elements.
Class Notes
A compound contains atoms of two or more
elements bound together in fixed proportions
by mass.
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6. ATOMIC STRUCTURE Present Concepts - An atom is an electrically
A chemical reaction involves a rearrangement of
of atoms but atoms are not created nor destroyed
during such reactions.
Present Concepts - An atom is an electrically
neutral entity consisting of negatively charged
electrons (e-) situated outside of a dense, posi-
tively charged nucleus consisting of positively
charged protons (p+) and neutral neutrons (n0).
Class Notes
Particle Charge Mass
Electron x g
Proton x g
Neutron x g
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7. ATOMIC STRUCTURE p+no e- e- no p+ Nucleus Model of a Helium-4
(4He) atom
p+no e- e-
no p+
Electron Cloud
How did we get this concept? - This portion of our
program is brought to you by:
Class Notes
Democritus, Dalton, Thompson, Planck, Einstein, Millikan,
Rutherford, Bohr, de Broglie, Heisenberg, Schrödinger,
Chadwick, and many others.
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8. ATOMIC STRUCTURE Democritus - First atomic ideas
Dalton First Atomic Theory
J. J. Thompson s - Measured the charge/mass
ratio of the electron (Cathode Rays)
Fluorescent
Material _
Cathode +
Class Notes
Anode
Electric Field
Source (Off)
With the electric field off, the cathode ray is not deflected.
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9. ATOMIC STRUCTURE - - + + - +
Fluorescent
Material -
Cathode + +
Anode
Electric Field
Source (On)
With the electric field on, the cathode ray is deflected
away from the negative plate. The stronger the electric
field, the greater the amount of deflection.
Class Notes -
Cathode +
Anode
Magnet
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10. ATOMIC STRUCTURE e/m = E/H2r e = the charge on the electron
With the magnetic field present, the cathode ray is
deflected out of the magnetic field. The stronger the
magnetic field, the greater the amount of deflection.
e/m = E/H2r
e = the charge on the electron
m = the mass of the electron
E = the electric field strength
H = the magnetic field strength
r = the radius of curvature of the electron beam
Class Notes
Thompson, thus, measured the charge/mass ratio
of the electron x 108 C/g
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11. ATOMIC STRUCTURE Summary of Thompson’s Findings:
Cathode rays had the same properties no matter
what metal was being used.
Cathode rays appeared to be a constituent of all
matter and, thus, appeared to be a “sub-atomic”
particle.
Cathode rays had a negative charge.
Cathode rays have a charge-to-mass ratio
of x 108 C/g.
Class Notes
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12. ATOMIC STRUCTURE e = 1.602 10 x 10-19 coulomb
R. A. Millikan - Measured the charge of the electron.
In his famous “oil-drop” experiment, Millikan was able to
determine the charge on the electron independently of its
mass. Then using Thompson’s charge-to-mass ratio, he
was able to calculate the mass of the electron.
e = x coulomb
e/m = x 108 coulomb/gram
m = x gram
Class Notes
Goldstein - Conducted “positive” ray experiments that
lead to the identification of the proton. The charge
was found to be identical to that of the electron and
the mass was found to be x g.
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13. ATOMIC STRUCTURE Ernest Rutherford - Developed the “nuclear” model
of the atom.
The Plum Pudding Model of the atom: + + +
A smeared out “pudding”
of positive charge with
negative electron “plums”
imbedded in it. + + + + + + +
The Metal Foil Experiments:
Fluorescent
Screen
a-particles
Radioactive
Material in
Pb box.
Class Notes
Metal
Foil
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14. ATOMIC STRUCTURE If the plum pudding model is correct, then all of
the massive a-particles should pass right through
without being deflected.
In fact, most of the a - particles DID pass right
through. However, a few of them were deflected at
high angles, disproving the “plum pudding” model.
Rutherford concluded from this that the atom con-
sisted of a very dense nucleus containing all of the
positive charge and most of the mass surrounded by
electrons that orbited around the nucleus much as
the planets orbit around the sun.
Class Notes
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15. ATOMIC STRUCTURE Assignment: V = (4/3)pr3 1. 2. 3. 4.
Assume the diameter of the nucleus of a hydrogen
atom is 1 x cm and the diameter of the atom
is 1 x cm.
1. Calculate the volume of the nucleus and the volume
of the atom in cm3 .
2. Calculate the volume of empty space in the atom.
3. Calculate the ratio of the volume of the nucleus to
volume of the whole atom.
4. Calculate the density of the nucleus if the proton’s
mass is x g
V = (4/3)pr3
Class Notes 1. 2. 3. 4.
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16. ATOMIC STRUCTURE Planck Bohr Problems with the Rutherford Model:
It was known from experiment and electromagnetic
theory that when charges are accelerated, they
continuously emit radiation, i.e., they loose energy
continuously. The “orbiting” electrons in the atom
were, obviously, not doing this.
Planck
Atomic spectra and blackbody radiation
were known to be DIScontinuous.
Class Notes
Bohr
The atoms were NOT collapsing.
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17. ATOMIC STRUCTURE Atomic Spectra - Since the 19th century, it had
been known that when elements and compounds
are heated until they emit light (glow) they emit
that light only at discrete frequencies, giving a
line spectrum. -
Class Notes +
Hydrogen
Gas
Line Spectrum
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18. ATOMIC STRUCTURE n/c = R[(1/m2) - (1/n2)]
When white light is passed through a sample of
the vapor of a substance, only discrete frequencies
are absorbed, giving an absorption ban spectrum.
These frequencies are identical to those of the
line spectrum of the same element or compound.
For hydrogen, the spectroscopists of the 19th
Century found that the lines were related by the
Rydberg equation:
Class Notes
n/c = R[(1/m2) - (1/n2)]
n = frequency
R = Rydberg Constant
c = speed of light
m = 1, 2, 3, ….
n = (m+1), (m+2), (m+3), ….
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19. ATOMIC STRUCTURE E = hn WAVE characteristics.
Max Planck - In 1900 he was investigating the nature
of black body radiation and tried to interpret his
findings using accepted theories of electromagnetic
radiation (light). He was NOT successful since these
theories were based on the assumption that light had
WAVE characteristics.
To solve the problem he postulated that light was
emitted from black bodies in discrete packets he
called “quanta”. Einstein later called them
“photons”. By assuming that the atoms of the black
body emitted energy only at discrete frequencies, he
was able to explain black body radiation.
Class Notes
E = hn
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20. ATOMIC STRUCTURE Both spectroscopy and black body radiation
indicated that atoms emitted energy only at
discrete frequencies or energies rather than
continuously.
Is light a particle or a wave??
Why do atoms emit only discrete energies?
What actually happens when light interacts
with matter?
Class Notes
What was wrong with Rutherford’s Model?
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21. ATOMIC STRUCTURE Niels Bohr - Bohr corrected Rutherford’s model
of the atom by formulating the following postulates:
Electrons in atoms move only in discrete orbits
around the nucleus.
When in an orbit, the electron does NOT emit
energy.
They may move from one orbit to another but are
NEVER residing in between orbits.
When an electron moves from one orbit to
another, it absorbs or emits a photon of light with a
specific energy that depends on the difference in
energy between the two orbits.
Class Notes
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22. ATOMIC STRUCTURE + The Bohr Model of the Atom Balmer Series Lyman
(Visible)
Lyman
Series
Paschen
Series +
(UV)
(IR)
Class Notes
The Bohr Model of the Atom
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23. ATOMIC STRUCTURE En = (- 2.179 x 10-18 J)/n2 Ephoton = Ehigh - Elow
The lowest possible energy state for an electron
is called the GROUND STATE. All other states
are called EXCITED STATES.
En = ( x J)/n2
Ephoton = Ehigh - Elow
Ephoton = [( x J)/n2high]
-[( x J)/n2low]
= x J[(1/n2high) - (1/n2low)]
Class Notes
Does this equation look familiar?
n/c = R[(1/m2) - (1/n2)]
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24. ATOMIC STRUCTURE l = h/mv Niels Bohr won the Nobel Prize for his work.
However, the model only worked perfectly for
hydrogen. What about all of those other elements??
Louis de Broglie - Thought that if light, which was
thought to have wave characteristics, could also have
particle characteristics, then perhaps electrons, which
were thought to be particles, could have characteristics
of waves.
Class Notes
l = h/mv
An electron in an atom was a “standing wave”!
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25. ATOMIC STRUCTURE (Dx)(Dp)  h/2p
Werner Heisenberg - Developed the “uncertainty”
principle: It is impossible to make simultaneous and
exact measurements of both the position (location)
and the momentum of a sub-atomic particle such as
an electron.
(Dx)(Dp)  h/2p
Our knowledge of the inner workings of atoms and
molecules must be based on probabilities rather
than on absolute certainties.
Class Notes
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26. ATOMIC STRUCTURE Hy = Ey H = Hamiltonian operator
Erwin Schrödinger - Developed a form of quantum
mechanics known as “wave mechanics”.
Hy = Ey
H = Hamiltonian operator
E = Total energy of the system
y = Wave function
[(-h2)/(8p2m)]2 - [e2/r]  = E
Kinetic Energy
Term
Potential Energy
Term
This is simply a quantum mechanical statement of the Law
of Conservation of Energy
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27. ATOMIC STRUCTURE  + - 
Of the numerous solutions to the Schrödinger equation
for hydrogen, only certain ones are allowed due to the
following boundary conditions:
Y, the wave function, must be continuous and finite.
It must be single-valued at all points (There can’t be
two different probabilities of finding an electron at one
point in space).
The probability of finding the electron, Y2, somewhere
in space must = 1. +

Y2dxdydz = 1
- 
Y has many values that meet these conditions. They are
called “orbitals”.
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28. ATOMIC STRUCTURE Y Y Y Y Y
Wave Function - A mathematical function associated
with each possible state of an electron in an atom or
molecule.
It can be used to calculate the energy of an
electron in the state
the average and most probable distance from the
nucleus
the probability of finding the electron in any
specified region of space. Y Y Y Y Y
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29. ATOMIC STRUCTURE Quantum Numbers:
Principle Quantum Number, n - An integer
greater than zero that represents the principle
energy level or “shell” that an electron occupies.
Energy # of orbitals
n Level Shell n2
1 1st K 1
2 2nd L 2
3 3rd M 9
4 4th N
etc. etc. etc. etc.
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30. ATOMIC STRUCTURE Azimuthal Quantum Number, l - The quantum
number that designates the “subshell” an electron
occupies. It is an indicator of the shape of an orbital
in the subshell. It has integer values from 0 to n-1.
l = 0, 1, 2, 3, …, n - 1
s p d f
Magnetic Quantum Number, ml - The quantum
number that determines the behavior of an electron
in a magnetic field. It designates the orbital and
has integer values from -l to +l including 0.
ml = -l, …, -3, -2, -1, 0, +1, +2, +3, …, +l
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31. ATOMIC STRUCTURE Orbital # of n l Name ml Orbitals 1 0 1s 0 1
p , 0, s
p , 0,
d -2, -1, 0, +1,
etc. etc. etc etc. etc.
Spin Quantum Number, ms - The quantum number
that designates the orientation of an electron in a
magnetic field. It has half-integer values, +½ or -½.
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32. ATOMIC STRUCTURE So what do atoms look like?
A. Interpretation of Y: The probability of finding
an electron in a small volume of space centered
around some point is proportional to the value of
Y2 at that point.
B. Electron Probability Density vs. r
C. Dot Density Representation: Imagine super-
imposing millions of photographs taken of an
electron in rapid succession.
D. Radial Densities
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33. ATOMIC STRUCTURE Electron Configuration
A. Many-electron atom: An atom that contains
two or more electrons.
B. Problems with the Bohr model:
1. It “assumed” quantization of the energy
levels in hydrogen.
2. It failed to describe or predict the spectra
of more complicated atoms.
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34. ATOMIC STRUCTURE Energy C. What are the differences in electron energy
levels in hydrogen vs. more complicated atoms?
3s 3p d
Energy 2s 2p
Ground State Hydrogen Atom 1s
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35. ATOMIC STRUCTURE Energy Splitting of the Degeneracy 2s 2p 2p 2s 1s 1s
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36. ATOMIC STRUCTURE Splitting of the Degeneracy
1. In hydrogen, all subshells and orbitals in a
given principal energy level have the same energy.
They are said to be Degenerate.
2. In many-electron atoms, s-orbitals have lower
energy than p-orbitals which have lower energy
than d-orbitals which have lower energy than
f-orbitals, etc., etc.
3. Reason: Complex electrostatic interactions.
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37. A. Shielding Effect - A decrease in the nuclear force + - ++
Hydrogen
+++
Helium
Lithium
A. Shielding Effect - A decrease in the nuclear force
of attraction for an electron caused by the presence
of other electrons in underlying orbitals.
B. Effective Nuclear Charge - A positive charge
that may be less than the atomic number. It is the
charge “felt” by outer electrons due to shielding by
electrons in underlying orbitals.
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38. ATOMIC STRUCTURE The Pauli Exclusion Principle - No two electron in
the same atom can have the same four quantum
numbers.
H e-  H -
Quantum Electron 1 Electron 2
Number n l ml
ms / /2
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39. The Aufbau Principle - A procedure for “building up”
the electronic configuration of many-electron atoms
wherein each electron is added consecutively to the
lowest energy orbital available, taking into account
the Pauli exclusion principle.
Order of Filling -
1s 2s 2p 3s 3p 4s 3d 4p 5s
Increasing Energy 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f 5g
mnemonic device
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40. Designating Electron Configurations - Standard Designation
H 1s1
He 1s2
Li 1s2 2s1
Be 1s2 2s2
B s2 2s2 2p1
C s2 2s2 2p2
Orbital Diagram Designation H He Li Be B C 1s 2s 1s 1s 2s 2p 1s 2s 2p 1s 2s 1s
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41. Core Designation - A designation of electronic
configuration wherein the outer shell electrons
are shown along with the “core” configuration of
the closest previous noble gas. Li Na K Rb
[He] 2s1 Be Mg Ca Sr
[He] 2s2
[Ne] 3s1
[Ne] 3s2
[Ar] 4s1
[Ar] 4s2
[Kr] 5s1
[Kr] 5s2
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42. Hund’s Rule of Maximum Multiplicity - Electrons
occupy a given subshell singly and with parallel spins
until each orbital in the subshell has one electron.
“Electrons try to stay as far apart as possible”
Elevator Analogy
Bus Seat Analogy
[He] 2s2 2p1
[He] B C N
[He] 2s2 2p2
[He]
[He] 2s2 2p3
[He] 2s 2p
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43. Assignment: Write the electron configuration using
all three types of designation for lead (Pb).
Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
4f14 5d10 6p2
Pb [Xe] 6s2 4f14 5d10 6p2
Electronic Configuration for postive ions (cations) -
Cations are formed by removing electrons in order
of decreasing n value. Electrons with the same n
value are removed in order of decreasing l value.
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44. Questions? Assignment: What are the electron configurations for
Fe2+ Fe3+ Cr Cr3+ Se2- ?
Questions?
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