Chemical Structure and Chemical Bonding. Acidity and Basicity

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Chemical Structure and Chemical Bonding. Acidity
and Basicity

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Outline 1. Bonding and antibonding orbitals 2.
Bonding between different elements 3. Chemical
structures in organic chemistry 4. Resonance
structures 5. Acidity and Basicity
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1. Bonding and antibonding orbitals
Comparison of Molecular Orbitals of Hydrogen (c1
and c2 )
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Electrons on bonding orbitals act as glue,
holding positively charged nuclei together. As
opposed to that, electrons on anti-bonding
orbitals pull the nuclei apart.
Three most important examples
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2. Bonding Between Different
elements Different elements have different
abilities to hold electrons around their nuclei.
Such ability is called Electronegativity Electron
egativity depends on the identity of the element,
on the total charge of the species and often on
the direction from the nucleus. From the
viewpoint of quantum chemistry, Electronegativity
- is the ability of an element to host lower in
energy molecular orbitals. It implies that
bonding, usually occupied orbitals, are shifted
toward the more electronegative atom that induces
a partial negative charge on the more
electronegative atom. Analogously, antibonding
orbitals are shifted toward the less
electronegative atom.
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In he Periodic table, electronegativity of
elements increases from the left to the right and
less prominently - from the bottom to the
top. Thats what is especially important
for us OgtNgtCgtH
If the bond is formed between atoms with very
different electronegativities (e.g. Na and F),
the charge separation becomes large and the bond
loses directionality, that is characteristic
for the Ionic bond. Ionic bonds are explained
well on Page 4 of your textbooks.
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Bonding in Methane
(CH4) We are starting from 8 electrons (4
valence electrons of carbon and 1 electron from
each hydrogen), placed on 8 orbitals (one 1s
orbital from each hydrogen, 2s, 2px, 2py and 2pz
orbitals of carbon). First we will mix orbitals
of hydrogens
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3. Chemical Structures in Organic
Chemistry A pair of electrons on a bonding
orbital is considered as a chemical bond. The
bond order (or number of bonds) is the number of
occupied bonding orbitals minus the number of
occupied antibonding orbitals. Thus, the
molecule of methane has four chemical bonds (one
is stronger and three other are a little weaker).
Each bond embraces all five atoms of the
molecule.
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Physical investigations of methane show that this
tetrahedral molecule has four equivalent
hydrogens, located at the same distances from
the central atom of carbon. Is it possible to
redistribute four all-molecule chemical bonds to
four carbon-hydrogen (C-H) chemical bonds?
Yes. In order to do it, we have to mix initial
atomic orbitals of carbon and hydrogens
differently Instead of optimizing them by
searching for the minimum of energy, they can be
optimized by searching for the maximum of the
electron density between pairs of the bonded
atoms of carbon and hydrogen. This procedure
leads us to the set of Natural Orbitals. Each
bonding natural orbital accounts for one specific
C-H bond.
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The diagram of energy for the Natural orbitals
doesnt match the diagram of energy for the
molecular orbitals. That is not
surprising, because the natural orbitals are not
optimized by energy and thus can not be used for
many energy-related conclusions.
Organic chemistry can not exist without chemical
structures. Natural orbitals give us the
theoretical basis to draw them!
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Like the pair of bonding and antibonding orbitals
in H-H is formed from atomic orbitals of
hydrogen, each pair of natural orbitals (bonding
and antibonding) in methane is formed from an
atomic orbital of hydrogen and an appropriate
orbital of carbon (green on the picture).
All green orbitals of carbon have the same
shape. They can be obtained by the linear
combination of s- and p- atomic orbitals of
carbon.
Since the atom of carbon has one s and three
p- orbitals, they are mixed at the ratio 1 3
and called sp3 hybrid orbitals.
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Summary of the Structural
Theory 1. All-molecule chemical bonds are
described in terms of Molecular orbitals. 2. Same
bonds, redistributed to two-center bonds (each
bond between a particular pair of atoms) are
described in terms of Natural orbitals. 3. Each
pair (bonding and antibonding) of Molecular
orbitals is built from Atomic orbitals. 4. Each
pair of Natural orbitals (bonding and
antibonding) is split to Atomic orbitals and
Hybrid orbitals. 5. Hybrid orbitals are built
from Atomic orbitals by simple averaging them. 6.
sp3 Hybridized orbitals are instances of
Hybrid orbitals.
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Next we need to bring unoptimized by energy
convenient chemical structures up to optimized by
energy, but messy systems of molecular orbitals
by applying corrections.
There are two options 1. Quantitative
adjustments Computer quantum chemistry
calculations of the energies of interactions
between Natural orbitals. This option goes beyond
our course. 2. Qualitative considerations of a
whole bunch of effects in organic Chemistry
(inductive effect, resonance effect, steric
hindrances etc.) We will use this option a lot.
A bit misleading, but valid definition A carbon
described in terms of sp3-hybridized orbitals is
said to be sp3-hybridized. Hybridization is a
mathematical trick rather than a real process.
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A little of
history The Structural theory in organic
chemistry was developed by Kekule, Couper and
Butlerov in 1861. The concept of hybridization
was suggested by Pauling in 1928 and brought him
Nobel prize in 1954 The theory of Natural
Orbitals and corresponding computational software
was introduced by Foster and Weinhold in 1980.
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Lewis Structures and Molecular
Geometry Lewis structure is a chemical
structure, where each occupied Natural Orbital is
presented by a pair of dots-electrons, and each
half-occupied natural orbital is presented by a
single dot-electron. In other words, every
localized (two-center) chemical bond is presented
by a pair of electrons. Limitation No
electrons are allowed on antibonding
orbitals. Purpose An excellent tool for 1.
Prediction of both aspects of chemical structures
(connectivity and molecular geometry). 2.
Prediction of chemical and physical properties.
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Building Lewis
Structures First, remember that elements of the
2nd period can not have more than 8 electron
around their nuclei. Since every bond is
presented by a pair of electrons, the maximum
number of bonds (valency) can not exceed 4.
Lewis Bases are donors of a pair of electrons.
Compounds with a lone electron pair (LEP) are
usually quite strong bases.
Shifts of electron pairs result in
redistribution of charge between atoms. On the
scheme on the left, a positive formal charge is
induced on nitrogen. This charge is not the same
as the actual charge.
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Actual and Formal
Charges Actual charges are real and can be
obtained by physical measurements. The negative
charge goes to more electronegative atoms, the
positive charge goes to less electronegative
atoms. For example, in NH4, the actual positive
charge is located on less electronegative
hydrogens. Formal Charge is the difference
between the nuclear charge, corresponding to the
valence electrons and the number of electrons,
contributing to the atomic charge. One way to
figure out the formal charge is to write a
reaction scheme, or it can be calculated, using
the definition.
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• Rules for Writing
  Structural Formulas
• Always show the formal charge
• 2. Dont show non-valence electrons
• 3. Show unshared electrons, if its essential for
  what you are writing the structure for.
• 4. If you have shown any of shared or unshared
  electrons, you must show all unshared electrons.
• 5. The doublet for the elements of the 1st period
  and the octet for the elements of the 2nd period
  may not be exceeded, but not necessarily
• complete.

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Molecular
geometry 1. Bond lengths depend on which atoms
are connected and how they are connected. Its
explained well in your textbooks (page 13). 2.
Bond angles. The rule for prediction (VSEPR
approach) All electron pairs (shared or
unshared) around a nucleus are arranged so that
they are as far from one another as possible.
(Exception double bonds behave like one
electron pair, because those pairs can not be
separated). Limitations There must not be
unpaired (single) electrons in the system.
Possible Cases
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Physical methods to determine molecular
geometry X-Ray crystallography in the solid
state Electron diffraction and Microwave
spectroscopy in the gas phase
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4. Resonance
Structures Remember, that Lewis structures are
based on natural orbitals, not optimized by
energy. Thus its not surprising that for some
chemical species we can not find a good Lewis
structure. What can we do?
We need to describe such problematic structures
with a set of structures rather than with a
single structure. Such set of structures is
called a set of resonance structures. On pages
18, 19 of your textbook there are good examples
of using resonance structures. Dont confuse
chemical equilibrium and resonance structures.
The actual molecule is planar, but doesnt have
such strong charge separation as would be
predicted by the structure on the right.
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5. Acidity and Basicity
Lewis acids accept a pair of electrons from a
low-energy orbital, producing the conjugate
base. Examples AlCl3, BF3, H Lewis bases
donate a pair of electrons to a low-energy
orbital, producing the conjugate acid. Examples
OH-, CH3O-, NH3 Bronsted acids donate H,
producing the conjugate base. Examples HCl,
H Bronsted bases accept H, producing the
conjugate acid. Examples OH-, CH3O-, NH3
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In aqueous solutions Acids produce H and the
conjugate base. Example CH3COOH H CH3COO-
Bases produce OH- and the conjugate
acid. Example NH3 H2O OH- NH4 Acids
react with bases and produce a salt (which can be
another base at the same time) and water (which
also can act as an acid). Example
Most of the compounds exhibit both acidic and
basic properties. Such compounds are called
amphoteric compounds.
This equilibrium is shifted toward the weaker
acid and the weaker base. A stronger acid has a
lower value of pKa and a weaker conjugate base.
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Henderson-Hasselbalch equation pH pKa
lg(A-/HA) When the pH equals pKa, the
solution has the same ability to protonate the
conjugate base A- as the acid HA has the ability
to protonate the solvent, and hence the acid
exists as a 11 ratio of its HA and A- forms
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Hard acids and bases have low polarizability.
Their orbitals do not change their shapes, so the
interaction is mostly driven by electrostatic
forces. Soft acids and bases have high
polarizability. Their orbitals change their
shapes and drive the interaction.
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Electrolytic dissociation
1. Acids produce H and the anion of the
conjugated base. Example HCl H
Cl- Presentation in reaction schemes HCl (or
another acid), H, H3O 2. Hydroxides of alkali
and alkali earth metals produce OH- and the
cation of the corresponding metal. Example NaOH
Na OH- Presentation in reaction schemes
NaOH (or another hydroxide), OH- 3. Salts
produce cations of the metal and anions of the
rest of the molecule. Example 1 Na2Cr2O7 2Na
Cr2O72- Presentation in reaction schemes
Na2Cr2O7, Cr2O72- (reactive anion) Example 2
Hg(CH3COO)2 Hg2 2CH3COO- Presentation in
reaction schemes Hg(CH3COO)2, Hg2 (reactive
cation)