Atomic and Molecular Structure
Michael Abosch
Brian Pflaum
2nd Period
Unit Outline
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Atomic and Electronic Structure, and Quantum
Mechanics
Periodic Trends
Molecular Structure
Bonding Theory
The Wave Nature of Light
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Electromagnetic Radiation- all visible light, radio waves,
infrared, X-rays etc.
Electromagnetic Spectrum- shows radiation arranged in
order of increasing wavelength
Visible light is only a small portion of spectrum.
http://steve.files.wordpress.com/2006/03/Elcetromagnetic%20spectrum.jpg
The Wave Nature of Light
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fλ= c
(frequency)(wavelength)= Speed of light (2.9979x108
m/s)
Frequency measured in s-1 (often Hz)
Wavelength measured in meters (often nm,μm)
The Quantization of Energy
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Quantum=The smallest quantity of energy that
can be emitted or absorbed as electromagnetic
radiation.
Energy, E, of a single quantum equals a
constant times the frequency of radiation.
E=hf
h=planck’s constant=6.626X10-34Joule-seconds.
Photoelectric Effect
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When photons of sufficiently high energy
(greater than the individual metal’s threshold
energy) strike a metal surface, electrons are
emitted from the metals
Energy of Photon, E=hf (planck’s
constant)(frequency)
Kinetic Energy of ejected electrons:
KEe=Ephoton-Ethreshold of metal
Wave Behavior of Matter
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Dual nature of radiant energy: both particle and
wave-like properties
DeBroglie wavelength: wavelength=(Planck’s
constant)/(momentum)=(h)/(mv)
Mass in Kg, Velocity in m/s
Orbitals
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An allowed energy state of an electron in the
quantum mechanical model of the atom;
describes the spatial distribution of the electron.
The orbital is defined by the values of quantum
numers n, l, and ml
The Principal Quantum Number
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The principal quantum number, n, can have
positive integral values of 1,2,3 etc…As n
increases, the orbital becomes larger, and the
electron spends more time farther from nucleus
The Azimuthal Quantum
Number
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The azimuthal quantum number, l, can have integral
values from 0 to n-1 for each value of n. This
quantum number defines the shape of an orbital.
The value of l is generally designated by the letters
s,p,d, and f.
Value of l
0123
Letter Used
s pdf
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The Magnetic Quantum Number
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The magnetic quantum number, ml, can have
integral values between -l and l. Describes
orientation of orbital in space.
Relationship amongst Quantum
Numbers
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Spin Magnetic Quantum Number
and Pauli Exclusion Principle
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The Spin Magnetic Quantum Number, ms, has
two possible values: +1/2, -1/2.
No two electrons in an atom can have the same
set of four quantum numbers n, l, ml, and ms
Thus, an orbital can hold a maximum of two
electrons, and they must have opposite spins.
Electron Configurations
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Electron Configuration=A particular
arrangement of electrons in the orbitals of an
atom.
The orbitals are filled in order of increasing
energy, with no more than two electrons per
orbital.
Orbital Diagrams
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Each orbital is denoted by a box,
and each electron by a half arrow
(which represents spin-up or spindown)
Electrons having opposite spins are
said to be paired when they are in
the same orbital
An unpaired electron is one not
accompanied by a partner of
opposite spin.
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=9887
Hund’s Rule
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Hund’s Rule=For orbitals of the same
energy level, the lowest energy is attained
when the number of electrons with the
same spin is maximized.
Note how in the diagram below, all three p
orbitals are filled singularly before an
electron is paired
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The Periodic Table and Electron
Filling Order
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Condensed Electron
Configurations
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The Electron configuration of the most recent nobel gas is
represented by its chemical symbol in brackets. From there,
Just proceed in the normal filling order until you reach the element.
In Potassium, the previous noble gas is argon, and its remaining
Electron occupies just one of the s orbitals, hence why it is denoted
As 4s1
Ions
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Start by writing the electron configuration for
the normal element
Then remove (or add) electrons as necessary,
always taking (or adding) from the highest
principle quantum number first (ignoring the
filling order).
Fe=[Ar]4s23d6
Fe(II)=[Ar]3d6
Anomalous Electron
Configurations
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Electron configurations of certain elements appear to
violate the “rules”
Frequently occurs when there are enough electrons to
lead to precisely half-filled sets of degenerate (same
energy-level) orbitals, or to completely fill an orbital.
This conserves Energy
No universal pattern or predictability
Ex: Chromium is [Ar]4s13d5 instead of [Ar]4s23d4
Practice
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What’s the electron
configuration for Lead?
Answer:
[Xe]6s24f145d106p2
Assign Quantum
numbers to it’s last filled
electron.
Answer: n=6, l=1, ml=0,
ms=+1/2
http://www.elementsdatabase.com/
Periodic Trends
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Atomic Size
Ionic Size
Ionization Energies
Electronegativity
campus.ru.ac.za/full_images/ img05206111510.jpg
Atomic Size
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Within each group, size increases from top to
bottom, results primarily from the increase in
principle quantum number of electrons
In each period, atomic radius tends to decrease from
left to right. Increase in the effective nuclear charge
as we move across a row steadily draws valence
electrons closer to nucleus
Exceptions: The addition of a paired electron
produces increased repulsion that sometimes leads
to an increase in size (Like from a p3 to a p4
element.)
Atomic Size
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Ionic vs. Atomic Size
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Cations: Compared to its neutral atom, cations
are smaller because electrons have vacated the
biggest orbital
Anions: Compared to its neutral atom, anions
are larger because adding electrons increases
repulsions, which leads to more space.
Ionic vs. Atomic Size
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7306
Isoelectronic Series
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Isoelectronic Series=A group all containing the
same number of electrons. As the atomic
number increases, the radius decreases.
Ex: Cl-, Ar, K+
Size: Cl->Ar>K+
Ionization Energy
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Ionization Energy=The minum energy required
to remove an electron from the ground state of
the isolated gaseuous atom or ion
The Greater the ionization energy, the more
difficult it is to remove an electron.
Variations in Successive
Ionization Energies
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I1>I2>I3 etc…
It’s more difficult to pull
away an electron from an
increasingly morepositive ion
There is a sharp increase
in ionization energy to
remove a core electron,
as they are closer to the
nucleus.
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Periodic Trends in First
Ionization Energy
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Within each period, ionization energy generally
increases with increasing atomic number.(Smaller
atomic radius)
Within each group, Ionization generally decreases from
top to bottom (Larger atomic radius).
Irregularities: Added “p” orbital sometimes leads to
decrease in ionization energy because the “p”
orbitals have more space than the “s” orbitals. Adding
a paired electron can also lead to a decrease in
ionization energy, as there is increased electron-electron
repulsion.
Periodic Trends in First
Ionization Energy
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Electronegativity
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Electronegativity=An order
of an atom’s overall ability to
attract electrons. It combines
atomic size and ionization
energy into a single summary
number.
http://www.green-planet-solar-energy.com/images/PT-small-electroneg.gif
Covalent Bonding
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Created when two atoms
share electrons
Strive to fulfill the Octet
rule- “atoms tend to gain,
lose, or share electrons until
they are surrounded by eight
valence electrons”
Many covalent bonds are
exceptions to the octet rule
www.ider.herts.ac.uk/.../covalent_bonding.gif
http://academic.brooklyn.cuny.edu/biology/bio4f
v/page/covalent-hydrogen.jpeg
Lewis Symbols
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Consists of the Atom’s chemical symbol, plus
one dot for every valence electron it has
Anions have extra dots, cations fewer dots
Examples:
..
.
.
H•
: Ar :
:F:
• C•
..
.
.
Drawing Lewis Structures
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Write the Chemical symbols for every atom in the molecule
The atom that makes the most bonds is generally the central
atom
Determine the total amount of Valence Electrons in the
molecule
Place single bonds between all atoms in the molecule that bond
With remaining electrons, fill up octets on all the atoms
If extra electrons exist, place them on the central atom
If too few electrons exist, create double, or triple bonds, keeping
the octet rule in mind.
Drawing Lewis Structures
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Example- CO2
Write all Chemical symbols
Carbon makes more
bonds (4) than oxygen (2)
O+O+C = 6+6+4= 16
O=- C-=O
Place single bonds
Fill all Octets
Not Enough! Must
make double bonds
This Creates 16
electrons, while
satisfying the octet
rule
Formal Charge
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Formal charge= the charge the atom would have if
each bonding electron pair were shared evenly between
its two atoms
To determine formal charge draw Lewis structure, and
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Count all unshared electrons per atom
Add half of the single, double, or triple bonds electrons to
the total (either 1,2, or 3 electrons)
Subtract this number from that atom’s usual amount of
valence electrons
Formal Charge
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Count all
unshared
electrons
Example- CN-
Add half
of bond
total
[:C≡N:]
2
+(6/2)
2
+(6/2)
4- 5 =
5- 5 =
-1
0
Subtract from
Atom’s usual
amount of valence
electrons
Electron Domains
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Any Bond (only single bonds) plus electron pairs (or
last unpaired electron) counts as an electron domain.
Electron Domains are important in understanding
molecular shape
Shapes are categorized by the amount of total electron
domains, then described further by the amount of
bonding domains
If an atom has 5 electron domains, but only 3 are
bonding domains, the other 2 are considered non
bonding domains, and are lone pairs.
Molecular Shapes
5 Basic Shapes
•Shape based on number of electron domains in the molecule
Linear
Trigonal
Bipyramidal
Trigonal Planar
Octahedral
Tetrahedral
All Pictures: chemlab.truman.edu/.../MM1Files/Linear3.gif
Linear
•One or Two electron Domains
•1 or 2 bonding domains
•Bond angles = 180˚
•Example- CO2
www.renewacycle.com/2007_02_01_archive.html
Trigonal Planar
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Three Electron Domains
Bond angle = 120˚
3 bonding domainstrigonal planar
Ex. BF3
2 bonding domains- bent
molecule
Ex. bent- NO2
Trigonal Planar
Bent
Tetrahedral
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Four Electron Domains
Bond Angle109.5˚
4 bonding domains- Tetrahedral
ex. CH4
3 bonding domains- trigonal
pyramidal
ex. NH3
2 bonding domains- bent
Ex. H2O
Tetrahedral
Trigonal pyramidal
Bent
Trigonal Bipyramidal
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Five Electron Domains
Bond Angles- Equatorial 120˚
Polar 180˚
5 bonding domains- trigonal
bipyramidal- ex. PCl5
4 bonding domains- Seesaw-ex.
SF4
•3 bonding domainsT-shaped- ex. ClF3
•2 bonding domainsLinear- ex. XeF2
Linear
See-Saw
Trigonal
Bipyramidal
T-Shaped
Octahedral
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Octahedral
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Square Pyramidal
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Square Planar
6 Electron Domains
Bond Angles- Equatorial- 90˚, Polar 180˚
6 bonding domains- Octahedral
Ex. SF6
5 bonding domains- Square Pyramidal
Ex. BrF5
4 bonding domains- Square Planar
Ex. XeF4
Dodecahedral
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Just Kidding
Bond Order & Length
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Double bond= bond order of 2
Triple bond = bond order of 3
As Bond order increases, bond length decreases
As Bond order increases, greater repulsive forces
exist between adjacent electron domains,
creating bigger angle
As Bond order increases, more energy is needed
to break the bond
Bond Polarity
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Happens when electrons are shared unevenly between atoms
Therefore does not happen between like atoms (i.e. H-H)
Generally, electronegativity differences of .4 or higher are
considered polar
When electronegativity difference is great enough, the bond is
considered ionic, not polar covalent
Ex.
H-H
C≡N
Na-Cl
2.1-2.1-0 2.5-3.0=.5
.9-3.0= 2.1
0<.4
.5>.4
2.1>>.4
Nonpolar
Polar
Ionic
Molecular Polarity
When is a molecule Polar?
No
Polar Bonds
Present?
Yes
Nonpolar
Molecule
Yes
Polar Bonds
arranged
symmetrically?
No
Polar
Molecule
http://bitesizebio.com/wp-content/uploads/2007/12/picture-1.png
Molecular Polarity
Symmetrical Molecules
Linear
Asymmetrical Molecules
Trig. Planar
Bent
Square Pyramidal
Tetrahedral
Trig. Bipyramidal
Pyramidal
Octahedral
Square Planar
Seesaw
T-Shaped
Valence Bond Theory
(hybrid orbitals)
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Bonds occur when electron shells overlap
Since electrons are simultaneously attracted to both
nuclei, bonds occur
Valence bond theory alone does not explain polyatomic
molecules. For this, Hybrid orbitals are needed
sp orbitals
Consider the linear non-polar BeF2 molecule
As it is, the molecule would not bond, since it has a full 2s shell
1s 2s
2p
After promotion…
The molecule can now bond, however it would
make non identical polar bonds
1s 2s
2p
By Hybridizing the
2s and 2p shells…
Now the Be molecule can make 2 identical
bonds
1s
sp
2p
The remaining 2p orbitals end up unhybridized
Additionally, the bigger lobes produced by the sp orbital allow for more overlap, which means
stronger bonds
All 2p orbitals can be hybridized. When this occurs, the same amount of orbitals must be
created. Ex.
1s
2s
2p
promote
1s 2s
2p
hybridize 1s
sp2
2p
More Hybrid Orbitals
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sp makes 2 180˚ orbitals
sp2 makes 3 120˚orbitals
sp3 makes 4 109.5 tetrahedron-arranged
orbitals
http://upload.wikimedia.org/wikipedia/commons/thumb/9/9f/Sp3Orbital.svg/290px-Sp3-Orbital.svg.png
Sigma(σ) and Pi(π) Bonds
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σ bonds- bonds occurring on the internuclear axis
π bonds- bonds occurring between two p orbitals oriented
perpendicularly to the internuclear axis.
π bonds produce a sideways overlap, which is not as
substantial, and therefore, not as strong, as a σ bond
Single bonds are σ bonds, double bonds consist of one σ bond
and one π bond, triple bonds have one σ bond and 2 π bonds
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Molecular Orbital Theory
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Better explains excited states of molecules
Each Molecular Orbital (MO) holds up to two
electrons, of opposing spins
Associated with the entire molecule, not just a
single atom
Molecular Orbital Theory
•Easiest way to analyze is through an energy level
diagram
•Bottom half of each shell is the bonding molecular
orbital, and is lowest energy.
•Top half is the antibonding molecular orbital, and is
higher in energy
•As the name suggests,
antibonding orbitals cancel out
the bonding orbitals
•Because energy increases as
the position on the chart
increases, slots are filled in from
the bottom up
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Molecular Orbital Theory
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Additionally, based on the positioning in the diagram,
the bonds can be analyzed as either σ or π bonds
The diagrams can be used to determine whether or not
an atom would form naturally
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Bond Order
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Bond order = ½(# of bonding electrons- # of
antibonding electrons)
Bond order of H2 = ½(2-0) = 1, Therefore H2 exists
Bond order of He2 = ½(2-2) = 0. Therefore, Helium
is not diatomic in nature
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