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7/27/2019 2013 Lect2b Chemical Properties- Electronic Structure and Chemical Bonding
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Structure BASICS
of Organic Molecules
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Objectives
Objectives
Know how to use the periodic table
Understand atomic structure of an atom including its massnumber, isotopes, and orbitals
Know how atomic orbitals overlap to form molecular orbitals Understand orbital hybridization
Using the VSEPR model, predict the geometry of molecules
Understand the formation of molecular orbitals
Know how to draw Lewis structures
Predict the direction and approximate strength of a bond dipole
Using a Lewis structure, find any atom or atoms in a moleculethat has a formal charge
Understand how to draw resonance structures
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ELECTRON CONFGURATION
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The basic unit of matter is an atom. It consists dense centralnucleus surrounded by negatively charged electrons.
The nucleus contains a mix of positively charge proton and electrically
neutral neutrons.
The electrons of an atom are bounded to the nucleus by
electromagnetic force
The electrons determine the chemical properties of an
element .
ATOM AND ITS ELECTRONS
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It means the arrangement of electrons of an atom.
The knowledge of electron configuration of different
atoms is useful in understanding the structure of
elements in periodic table.
The concept also useful for describing chemical bonds
that hold atoms together
ELECTRON CONFIGURATION
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Chemical BondFormation
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A chemical bondis an attraction between atoms that allowsthe formation of chemical substances that contain two or more atoms.
The bond is caused by the electrostatic force of attractionbetween opposite charges, either between electrons and
nuclei, or as the result of a dipole attraction.
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the negatively charged electrons that are orbiting the nucleus and the
positively charged protons in the nucleus attract each other.
Also, an electron positioned between two nuclei will
be attracted to both of them
These electrons cause the nuclei to be attracted toeach other, and this attraction results in the bond
However, this assembly cannot collapse to a size dictated by the
volumes of these individual particles. Due to the matter wave natureof electrons and their smaller mass, they occupy a much larger
amount of volume compared with the nuclei, and this volume
occupied by the electrons keeps the atomic nuclei relatively far apart,
as compared with the size of the nuclei themselves.
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In general there are two types of chemical bonds:
Ionic Bonds - Bonds formed when one or more
electrons are transferred from one atom
to another atom.
This attraction may be seen as the result of different
behaviors of the outermost electrons (VE) of atoms.
The strength of chemical bonds varies considerably;
there are "strong bonds" such as covalent or ionic bonds
and "weak bonds" such as dipoledipole interactions, theLondon dispersion force and hydrogen bonding.
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In this type of bond, the outer atomic orbital of one atom has a vacancy
which allows addition of one or more electrons.
These newly added electrons potentially occupy a lower energy-state than
they experience in a different atom. Thus, one nucleus offers a more tightly
bound position to an electron than does another nucleus, with the result that
one atom may transfer an electron to the other.
This transfer causes one atom to assume a net positive charge, and theother to assume a net negative charge. The bondthen results from
electrostatic attraction between atoms, and the atoms become positive or
negatively charged ions.
Often, such bonds have no particular orientation in space, since they result
from equal electrostatic attraction of each ion to all ions around them.
Ionic bonds are strong since the forces between ions are short-range, and do
not easily bridge cracks and fractures. This type of bond gives a charactistic
physical character to crystals of classic mineral salts, such as table salt.
a simplified view of an ionicbond,
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Ionic Bonding
Na F
Sodium Atom Fluorine Atom
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Ionic Bonding (2)
Na F
Attraction between the two ions is electrostatic --
I onic Bond
Sodium ion Fluoride ion
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Covalent Bonds -
Bonds formed when two
atoms share one or more
pairs of valence electrons.
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a pair of electrons are drawn into the space between the two atomic
nuclei. Here the negatively charged electrons are attracted to the
positive charges ofboth nuclei, instead of just their own. This
overcomes the repulsion between the two positively charged nuclei of
the two atoms, and so this ing attraction holds the two nuclei in a fixed
configuration of equilibrium,
Thus, covalent bonding involves sharing of electrons in which thepositively charged nuclei of two or more atoms simultaneously attract
the negatively charged electrons that are being shared between them.
These bonds exist between two particular identifiable atoms, and have
a direction in space, allowing them to be shown as single connectinglines between atoms in drawings, or modeled as sticks between
spheres in models.
In a polar covalent bond, one or more electrons are unequally shared
between two nuclei.
the simplest view of a so-called 'covalent' bond,
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Covalent bonds often result in the formation of small collections of better-
connected atoms called molecules, which in solids and liquids are bound toother molecules by forces that are often much weaker than the covalent
bonds that hold the molecules internally together. Such weak intermolecular
bonds give organic molecular substances, such as waxes and oils, their
soft bulk character, and their low melting points (in liquids, molecules must
cease most structured or oriented contact with each other).
When covalent bonds link long chains of atoms in large molecules, however
(as in polymers such as nylon), or when covalent bonds extend in networks
though solids that are not composed of discrete molecules (such as
diamond or quartz or the silicate minerals in many types of rock) then the
structures that result may be both strong and tough, at least in the direction
oriented correctly with networks of covalent bonds. Also, the melting points
of such covalent polymers and networks increase greatly.
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A SIMPLE COVALENT BOND
H . H.
A pair of electrons is shared between the two bonded atoms.
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When two atoms move into close proximity, they experience a change in energy. At the
distance ofthe bond length, they achieve minimum energy.
energy is released during the formationof the bond.
DISTANCE AND ENERGY RELATIONSHIP
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BOND DISSOCIATION ENERGIES.
Conversely, breaking bond of two atoms in a molecule requires an input
of energy because the energy level of the molecule is lower than the
energy level of the two atoms.
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The electrons of the atoms possess a set of stable energy levels
or ORBITALS, and can undergo transitions between orbital by
absorbing or emitting photons that match the energy differences
between the level.
ELECTRON CONFIGURATION
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Electrons of atom occupy a set of allowed states which called SHELL
the first shell can accommodate 2 electrons, the second shell
8 electrons, and the third shell 18 electrons
(An atom's nth electron shell can accommodate 2n2 electrons)
A SUBSHELL is the set of states defined by azimuthalquantum number, l = 0, 1, 2, 3 which correspond to the s,p, d,
and flabels. The max number of electrons which can be placed
in a subshell is given by 2(2l+ 1).
This gives :2 electrons in a s subshell,
6 electrons in a p subshell,
10 electrons in a d subshell
14 electrons in a f subshell.
SHELL AND SUBSHELL of ATOM
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Electron configuration
The outermost electron shell is often referred to as
the "VALENCE SHELL" and determines the chemical
properties
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LEWIS STRUCTURE
OF MOLECULES
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Know how to draw Lewis structures
Using a Lewis structure, find any atom
or atoms in a molecule that has a formal
charge
Understand how to draw resonance
structures
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Lewis structure
diagrams that show the bonding between atoms of a
molecule and the lone pairs of electrons that may exist
in the molecule.
A Lewis structure can be drawn for any covalently
bonded molecule, as well as coordination
compounds.
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C. Lewis Symbols for VE
The elements symbol represents the nucleus and the core
electrons.
Dots are placed around the symbol to represent the VE.
The electrons are paired if there are more than four VE.
A maximum of 4 pairs of electrons (8 e) areaccommodated by a main group element.
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Lewis Structures
show the arrangement of atoms and VE in a molecule.
A pair of electrons that are shared in a covalent bond is
represented by a line between a pair of atoms.
Electrons that are shared by two atoms are called bonding
pairs. Electrons that reside on a single atom and are not shared
are called lone pairsornonbonding electrons.
Some elements can share more than 1 pair of e in a
covalent bond called double bondsand triple bonds.
Most atoms in a molecule are surrounded by 8 electrons to
achieve noble gas configuration. The tendency of these
atoms to be surrounded by 8e is known as the octet rule.
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Construction of the Lewis Dot diagram
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The total number of electrons represented in a Lewis
structure is equal to the sum of the numbers of valence
electrons on each individual atom.
Non-valence electrons are not represented in Lewis
structures.
The octet rule states atoms with eight electrons in theirvalence shell will be stable, regardless of whether these
electrons are bonding or nonbonding.
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Lewis structures for oxygen, fluorine, thehydrogen sulfate anion, and formamide
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Drawing Lewis Structures
1. Decide on the central atom.
2. Determine the total number of valence electrons in the
molecule or ion.
3. Place one pair of electrons between each pair ofbonded atoms to form a single bond.
4. Use the remaining electrons as lone pairs around each
terminal atom (except H) and then the central atom so
that each is surrounded by eight electrons.
5. If the central atom has fewer than eight electrons, move
one or more lone pairs from terminal atoms to form
multiple bonds to the central atom.
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Predicting Lewis Structures
1. Hydrogen Compounds
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Example: Lewis structure of
the nitrite ion
The formula of the nitrite ion is NO2.
Step one: Nitrogen is the least electronegative atom, so it is the
central atom by multiple criteria.
Step two: Count valence electrons. Nitrogen has 5 valence electrons;
each oxygen has 6, for a total of (6 2) + 5 = 17. The ion has a charge of
1, which indicates an extra electron, so the total number of electrons is 18.
Step three: Place ion pairs. Each oxygen must be bonded to the nitrogen,
which uses four electrons two in each bond. The 14 remaining electrons
should initially be placed as 7 lone pairs. Each oxygen may take amaximum of 3 lone pairs, giving each oxygen 8 electrons including the
bonding pair. The seventh lone pair must be placed on the nitrogen atom.
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Step four: Satisfy the octet rule. Both oxygen atoms currently have
8 electrons assigned to them. The nitrogen atom has only 6 electrons
assigned to it. One of the lone pairs on an oxygen atom must form a
double bond, but either atom will work equally well. We therefore must
have a resonance structure.
Step five: Tie up loose ends. Two Lewis structures must be drawn:
one with each oxygen atom double-bonded to the nitrogen atom. The
second oxygen atom in each structure will be single-bonded to thenitrogen atom. Place brackets around each structure, and add the
charge () to the upper right outside the brackets. Draw a double-
headed arrow between the two resonance forms.
http://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.pnghttp://en.wikipedia.org/wiki/File:Nitrite-ion-lewis-canonical.png7/27/2019 2013 Lect2b Chemical Properties- Electronic Structure and Chemical Bonding
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Predicting Lewis Structures
2. Oxo Acids and Their Anions
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Predicting Lewis Structures
3. Isoelectronic species have the same number of valence
electrons and the same Lewis structure.
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RESONANCE
Some compounds can have more than one valid Lewisstructure.
Multiple Lewis structures that represent the same
molecule are called resonance structures.
Often the actual structure of a molecule
with resonance structures is a
composite of all the resonance
structures called a resonance hybrid.
RESONANCE HYBRIDE
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RESONANCE HYBRIDEMultiple Lewis structures
1) sulfur dioxide
2) nitric acid
3) formaldehyde
resonance
hybrid
major
contributor
minor
contributor
non
contributor
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4) carbon
monoxide
5) azide anion
RESONANCE HYBRIDE
Multiple Lewis structures
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D. Exceptions to the Octet Rule
1. Compounds in which an atom has fewer than eight
electrons. Group 3 atoms (B, Al, Ga, In, Tl) are
commonly surrounded by six
electron or three electron pairs.
These atoms can formcoordinate covalent bonds in
which the pair of electrons in
the covalent bond originate
from one atom.
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D. Exceptions to the Octet Rule
2. Compounds in which an atom has more than eight
electrons.
Element from the third period or higher can formmolecules or ions in which the octet rule is exceeded.
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D. Exceptions to the Octet Rule
3. Molecules with an Odd Number of Electrons Compounds containing an odd number of nitrogens
will have an odd number of electrons making it
impossible to draw a structure that obeys the octet
rule.
N
O O
N
OO
Chemical species with an unpaired electron are called
free radicals and are highly reactive.
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CHARGE DISTRIBUTION
IN COVALENT BONDS AND
MOLECULES
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We need to determine these loci of charge concentration,
to figure out if an atom is negative, positive, or neutral?
in order to understand chemical reactivity.
Many important organic species are ions.
In these ions, charges appear to be preferentially
concentrated on certain atoms.
C
O
OO
2-
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Calculation of Formal Charges
Formal
charge
Number ofValence
electrons
Number ofnonbonding
electron
Half number ofelectron in
covalent bond
= - -
Heres the formula for figuring out the formal
charge of an atom:
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Determination of Formal Charges
Determine the numberof valence electrons
which are present in the neutralatom.
Subtract the non-bonding electrons. Theseelectrons belong to the atom.
Subtract one-halfof the bonding electrons.
This formula explicitly spells out how many
electron are formally owned by the atom.
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is the charge assigned to an atom in a
molecule, assuming that electrons in a chemicalbond are shared equally between atoms,
regardless of relative electronegativity.
Formal charge
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Charge Distribution in Covalent Bonds
A. Formal Charge
LPE = lone pair electrons
BE = bonding electrons
BE)](Number21LPE[Number-NumberGroupChargeFormal
NH H
H
H N H
H
H
Formal Charge = 5 - [2 + (6)] = 0
Formal Charge = 1 - [0 + (2)] = 0
Sum of Formal Charges = 0 + 0 + 0 + 0 = 0
Formal Charge = 5 - [0 + (8)] = + 1
Formal Charge = 1 - [0 + (2)] = 0
Sum of Formal Charges = 1 + 0 + 0 + 0 = +1
Charge Distribution in Covalent Bonds
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Charge Distribution in Covalent Bonds
A. Formal Charge
LPE = lone pair electrons
BE = bonding electrons
BE)](Number21LPE[Number-NumberGroupChargeFormal
Formal Charge = 6 - [4 + (4)] = 0
C
O
OO
2-
Formal Charge = 6 - [6 + (2)] = - 1
Formal Charge = 4 - [0 + (8)] = 0
Sum of Formal Charges = -1 + -1 + 0 + 0 = -2
Lets apply it to some examples.
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http://masterorganicchemistry.files.wordpress.com/2010/09/formal-charge-copy.jpghttp://masterorganicchemistry.files.wordpress.com/2010/09/formal-charge-copy.jpg7/27/2019 2013 Lect2b Chemical Properties- Electronic Structure and Chemical Bonding
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More on Formal Charges
The sum of the formal charges must be equalto the total charge on the ion or molecule
If there arent enough electrons to provideevery atom with an octet, consider double ortriple bonds
All structures must obey the rules of valence
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Still More on Formal Charges
Separated formal charges should be
avoided, if possible
Where there are separated formalcharges, the negative formal charge
should reside on the more
electronegative element
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