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OVERVIEW
Kossel-Lewis Approach to chemical bonding Octet Rule
Covalent Bond
Lewis Structures Simple Molecules
Ionic or Electrovalent Bond
Bond Parameters
VSEPR Theory
Orbital Overlap Hybridisation
Molecular Orbital Theory
Summary
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Most bonds are somewhere in between.
Forms of ChemicalBonds
There are TWO extreme forms ofconnecting or bonding atoms:
Ioniccomplete transfer of electronsfrom one atom to another
Covalentelectrons shared betweenatoms
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Electron Distribution in
Molecules Electron distribution is
depicted with Lewis
electron dotstructures
Electrons aredistributed as:
shared or BONDPAIRS and
unshared or LONEPAIRS.
G. N. Lewis1875 - 1946
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Chemical bond: attractive force holding two or more
atoms together.
Covalent bond results from sharing electrons between the
atoms. Usually found between nonmetals.
Ionic bond results from the transfer of electrons from a
metal to a nonmetal.
Metallic bond: attractive force holding pure metalstogether.
Chemical Bonds, Lewis Symbols,and the Octet Rule
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All noble gases except He has an s2p6 configuration.
Octet rule: atoms tend to gain, lose, or shareelectrons until they are surrounded by 8 valence
electrons (4 electron pairs).
Caution: there are many exceptions to the octet
rule.
The Octet Rule
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Usually occurs with B and elementsof higher periods. Commonexceptions are: Be, B, P, S, andXe.
BF3BF3
SF4SF4
Be: 4
B: 6
P: 8 OR 10
S: 8, 10, OR 12
Xe: 8, 10, OR 12
Voilations of Octet Rule
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Central Atoms Having Less than an Octet
Relatively rare.
Molecules with less than an octet are typical for
compounds of Groups 1A, 2A, and 3A.
Most typical example is BF3.
Formal charges indicate that the Lewis structure with
an incomplete octet is more important than the oneswith double bonds.
Exceptions of Octet Rule
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Central Atoms Having More than an Octet
This is the largest class of exceptions.
Atoms from the 3rd period onwards can accommodate
more than an octet.
Beyond the third period, the d-orbitals are low enough in
energy to participate in bonding and accept the extra
electron density.
Exceptions of Octet Rule
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Covalent Bonding
C l
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CovalentBonding
Covalent bond is the sharing of the VALENCEELECTRONS of each atom in a bond
Recall: Electrons are divided between
core and valence electrons.ATOM core
valenceNa 1s2 2s2 2p6 3s1 [Ne]
3s1[Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5
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Valence Electrons1A
2A 3A 4A 5A 6A 7A
8A
Number of valence electrons is equalto the Group number.
C t
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Cova entBondingThe bond arises from the mutual attraction of 2 nuclei for the same electrons.
HB+ H
A
HB
HA
A covalent bond is a balance
of attractive and repulsive forces.
+
-
-
+
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BondFormation
A bond can result from a head-to-head overlap of
atomic orbitals on neighboring atoms.
H H Cl
Cl
+
Overlap of H (1s) and Cl (2p)
his type of overlap places bonding electrons in aOLECULAR ORBITAL along the line between
the two atoms and forms a SIGMA BOND (s).
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More sharing examples
O2
N2
O O
N N
O O O O
N N N N N N
double bond (2 pairs)
triple bond (3 pairs)
Share until octet is complete.
octet complete
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Bond Energy
F2
single bond BE = 142 kJ/mole
O2 double bond BE = 494
N2
triple bond BE = 942
X2 + energy X + X
increasing
bon
d
stren
gth
Is breaking a bond an endothermic or exothermic process?
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NH3NH
H
H
NH
H
HH+
NH4+ NH3 + H
+ NH4+
coordinate covalent bond(the pair of electronsfrom the same atom)
normal covalent bond(each atom supplies
an electron)
Some more sharing examples
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Follow Step by Step Method1. Total all valence electrons. [Consider Charge]
2. Write symbols for the atoms and guess skeleton
structure [ define a central atom ].
3. Place a pair of electrons in each bond.
4. Complete octets of surrounding atoms. [ H = 2 only ]
5. Place leftover electrons in pairs on the central atom.
6. If there are not enough electrons to give the central atom
an octet, look for multiple bonds by transferring
electrons until each atom has eight electrons around it.
Drawing Lewis Structure
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Resonance Structures
Drawing LewisStructures
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To determine the electron pair geometry: draw the Lewis structure,
count the total number of electron pairs around the central
atom,
arrange the electron pairs in one of the above geometries to
minimize e-e repulsion, and count multiple bonds as one
bonding pair.
VSEPR Model
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VSEPR Model
d l
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VSEPR Model
d l ( i
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VSEPR Model (Domainsaround Central Atom)
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Drawing Lewis Structures
Formal Charge
Consider:
For C:
There are 4 valence electrons (from periodic table).
In the Lewis structure there are 2 nonbonding electrons and 3
from the triple bond. There are 5 electrons from the Lewis
structure.
Formal charge: 4 - 5 = -1.
C N
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Drawing Lewis Structures
Formal Charge Consider:
For N: There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3
from the triple bond. There are 5 electrons from the Lewis
structure.
Formal charge = 5 - 5 = 0. We write:
C N
C N
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Bond and LonePairs
Electrons are distributed as shared or BONDPAIRS and unshared or LONE PAIRS.
H Cl
This is a LEWIS ELECTRON DOT structure.
shared or bond pair
Unshared orlone pair (LP)
Rules of Lewis
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his observation is called the OCTET RULE
Rules of LewisStructures
No. of valence electrons of an atom =Group number
xcept for H
(and atoms of 3rd and higher periods),#Bond Pairs + #Lone Pairs = 4
For Groups 5A-7A (N - F),o. of BOND PAIRS = 8 - group No.
or Groups 1A-4A (Li - C),o. of BOND PAIRS = group number
Building a Dot
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2. Count valence electronsH = 1 and N = 5
Total = (3 x 1) + 5= 8 electrons or
Decide on the central atom; never H.entral atom is atom of lowest affinity for electrons
n ammonia, N is central
Building a DotStructure
Ammonia, NH3
4 pairs
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4. Remaining electrons formLONE PAIRS to completeoctet as needed.
3. Form a sigma bondbetween the central atomand surrounding atoms.
H H
H
N
Building a Dot Structure
H H
H
N
3 BOND PAIRS and 1 LONE PAIR.e that N has a share in 4 pairs (8 electrons), whileh H shares 1 pair.
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Step 2. Count valence electronsS = 63 x O = 3 x 6 = 18
Negative charge = 2TOTAL = 6 + 18 + 2 = 26 e-
or 13 pairs
Step 1. Central atom = S
10 pairs of electrons are left.
Sulfite ion, SO32-
Step 3. Form sigma bonds
O O
O
S
Sulfite ion SO 2-
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Remaining pairs become lone pairs,first on outside atomsthen on central atom.
Sulfite ion, SO32
(2)
Each atom is surrounded by an octet ofelectrons.
O O
O
S
TE - must add formal charges (O-, S+) for complet diagram
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Carbon Dioxide, CO2
1. Central atom = __C____2. Valence electrons = _16_ or _8_
pairs
3. Form sigma bonds.O OC
O OC
This leaves __6__ pairs.4. Place lone pairs on outer atoms.
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O OC
Carbon Dioxide, CO2 (2)
4. Place lone pairs on outer atoms.
O OC
O OC
The second bonding pair forms a pi (p) bond.
. To give C an octet, form DOUBLE BONDSetween C and O.
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SO3
H2CO
Double and eventriple bonds arecommonly observedfor C, N, P, O, and S
O OC
C2F4
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Sulfur Dioxide, SO2
1. Central atom = S2. Valence electrons = 6 + 2*6 = 18 electrons
or 9 pairs
O OS
O OS
bring in
left pair
OR bring in
right pair
3. Form pi ( ) bond so that S has an octet note that there are two ways of doing
this.
Sulfur
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SulfurDioxide, SO2
O OS
bring inleft pair
OR bring in
right pair
O OS
O OS
Equivalent structures
called:RESONANCESTRUCTURES
The proper Lewis struct
is a HYBRID of the two.
ETTER representation of SO2
made by forming 2 double bonds
O = S = O
Each atom has- OCTET- formal charge = 0
Violations of the
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Violations of theOctet Rule
Usually occurs with:
Boron
BF3SF
4
Elements of higher periods.
Boron
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BoronTrifluoride
Central atom = B
Valence electrons = 3 + 3*7 = 24
or electron pairs = 12
Assemble dot structure
F
F
F
B
The B atom has a sharein only 6 electrons (or
3 pairs). B atom inmany molecules iselectron deficient.
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Central atom = S
Valence electrons = 6 + 4 * 7 = 34 e -
or 17 pairs.
Form sigma bonds and distributeelectron pairs.
F
F
F
S
F
5 pairs around the S atom. A5 pairs around the S atom. A
common occurrence outsidecommon occurrence outsidethe 2nd period.the 2nd period.
Sulfur Tetrafluoride, SF4
Formal Atom
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Formal charge = Group no.
- 1/2 (no. bond electrons)
- (no. of LP electrons)
Formal AtomCharges
he most important dominant resonance structuref a molecule is the one with formal chargess close to 0 as possible.
Carbon
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04 - (1/ 2)(8) - 0 =
6 - (1/2)(4) - 4 = 0
CarbonDioxide, CO2
At OXYGEN
O C O
At CARBON
Carbon Dioxide
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C atom chargeis 0
6 - (1/2)(6) - 2 = +1
6 - (1/2)(2) - 6 = -1
Carbon Dioxide,CO2 (2)
O C O
n alternate Lewis structure is:
AND the correspondingresonance form
+
O C O
+
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Boron Trifluoride, BF3
F
F
F
B
What if we form a BF doublebond to satisfy the B atomoctet?
Boron Trifluoride
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Boron Trifluoride,BF3 (2)
To have +1 charge on F, with its very high electronaffinity is not good. -ve charges best placed onatoms with high EA.
Similarly -1 charge on B is bad
NOT important Lewis structure
fc = 7 - 2 - 4 = +1 Fluorine
fc = 3 - 4 - 0 = -1 BoronF
F
F
B
+
St t D t i ti b
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Water, H2O
H O H
H O H
2 bond
pairs2 lone
pairs
Structure Determination byVSEPR
The electronpair
geometry isTETRAHEDRAL
Themoleculargeometry isBENT.
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Ammonia, NH3
The electron pair geometry is
TETRAHEDRAL.The Molecular geometry the positionsof the atoms is TRIGONAL PYRAMID.
Structure Determination byVSEPR
H
H
H
lone pair of electronsin tetrahedral position
N
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TAKE A BREAK
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Ionic Bonding
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Ionic Bonds
Ionic compounds
Essentially complete electron transfer froman element of low IE (metal) to an elementof high electron affinity (EA) (nonmetal)
Na(s) + 1/2 Cl2(g) Na+ + Cl-
NaCl (s)
Ionic bonding
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ClNa+ Cl
Ionic bonding Ionic bonding involves 3 steps (3 energies)
1) loss of an electron(s) by one element, 2) gain of
electron(s) by a second element, 3) attraction betweenpositive and negative
Na Cl
e1) 2)
3)
Na+
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NON-DIRECTIONALbonding
via Coulomb (charge)interaction
rimarily between metalsrps 1A, 2A and transition metals)
d nonmetals (esp O and halogens)
Ionic Bonds
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Bond Type Single Double Triple
# of es 2 4 6
Notation =
Bondorder
1 2 3
Bondstrength
Increases from Single to Triple
Bondlength
Decreases from Single to Triple
Chemical Bonds
Average Bond Enthalpies
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Average Bond Enthalpies(KJ/mol)
Average Bond Lengths of some Single
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Average Bond Lengths of some Single,Double and Triple Bonds
L i S b l /
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Lewis Symbols/Electronic Configuration
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Electronegativity: The ability of one atoms in a
molecule to attract electrons to itself.
Pauling set electronegativities on a scale from0.7 (Cs) to 4.0 (F).
Electronegativity increases
across a period and
down a group.
Electronegativity
Electronegativity of
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g yElements
Electronegativity and
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There is no sharp distinction between bonding types.
The positive end (or pole) in a polar bond is represented
+ and the negative pole -.
Electronegativity andBond Polarity
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HCl is POLAR because ithas a positive end and anegative end. (difference
in electronegativity)
Cl has a greater share inbonding electrons than
does H.
Cl has a greater share inbonding electrons than
does H.
Cl has slight negative charge (-d) and H has
slight positive charge (+ d)
H Cl
+ -
H Cl
+ -
Bond Polarity
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why oil and water will not mix! Oil isnonpolar, and water is polar.
The two will repel each other, and so youcan not dissolve one in the other
Bond Polarity
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Valence bond theory
Valence bond theory utilizes orbitals andelectrons coming together to form thecovalent bonds in a molecule.
According to valence bond theory, a bondbetween two atoms is formed when twoelectrons with their spins paired are
shared by overlapping atomic orbitals
Molecular Shapes
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There are five fundamental geometries for molecular
shape:
Molecular Shapes(VSEPR)
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Summary of VSEPR shapes
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e-pairs Notation Name of VSEPRshape
Examples
2 AX2 Linear HgCl2 , ZnI2 , CS2 , CO2
3 AX3 Trigonal planar BF3 , GaI3
AX2E Non-linear (Bent) SO2 , SnCl2
4 AX4 Tetrahedral CCl4 , CH4 , BF4-
AX3E (Trigonal) Pyramidal NH3 , OH3-
AX2E2 Non-Linear (Bent) H2O , SeCl2
5 AX5 Trigonal bipyramidal PCl5 , PF5
AX4E Distorted tetrahedral(see-sawed)
TeCl4 , SF4
AX3E2 T-Shaped ClF3 , BrF3
AX2E3 Linear I3- , ICl2
-
6 AX6 Octahedral SF6 , PF6-
y p
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Sigma Bond Formation by
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Sigma Bond Formation byOrbital Overlap
sigma bond ( )
+HH
Two s AtomicOrbitals (A.O.s)
overlap to form ans (sigma)Molecular Orbital
(M.O.)
Sigma Bond Formation by
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Orbital Overlap
sigma bond ( )
+HH
Two s A.O.s overlap tofrom an s M.O.
Similarly, two p A.O.s
can overlap end-on tofrom a p M.O.
eg F2
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Hybridization
In 1931, Linus Pauling proposed thewave functions
for the s and p atomic orbitals .
The mathematical process ofreplacing pure atomic
orbitals with reformulated atomicorbitals for bonded
atoms is called hybridization. In a hybridization scheme, the
number of hybrid
orbitals equals to the total number of
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H b idi ti
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3
signifies one s and three p orbitals are combined
Mixing one s orbital with three p orbitals yieldsfour equivalent sp3 hybrid orbitals.
Hybridization
Hybridization
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The formation of four sp3 hybrid orbitals by combinationof an atomic s orbital with three atomic p orbitals. Eachsp3 hybrid orbital has two lobes, one of which is largerthan the other. The four large lobes are oriented toward
the corners of a tetrahedron at angles of109.5.
Hybridization
Hybridization
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The bonding in methane. Each of the four C-H bonds
results from head-on (s) overlap of a singly occupiedcarbon sp3 hybrid orbital with a singly occupiedhydrogen 1s orbital. Sigma bonds are formed byhead-to-headoverlap between the hydrogen s orbital
and a singly occupied sp3
hybrid orbital of carbon.
Hybridization
sp hybridization
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sp hybridizationNow consider BeCl
2which has linear molecular
geometry determined experimentally.
The combination of one s and one p orbital gives twosp hybrid orbitals oriented 180 apart. Two
unhybridized p orbitals remain and are oriented at 90
sp2 hybridization
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sp hybridization
E.g. the molecular
geometry is trigonal
planar with bond angle =120. To explain its
geometry, we can use the
following rational.
sp2 signifies one s and two
p orbitals are combined.
sp3d hybrid Orbitals
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sp d hybrid Orbitals
For Hybridization scheme to correspond to the 5- and 6- electron-
group geometries of VSEPR theory, we need to go beyond s and porbitals and traditionally this meant including d orbitals.
We can achieve the five half-filled orbitals and trigonal-bipyramidal
molecular geometry through the hybridization of one s, three p and
one d orbitals of valence shell into five sp3
d hybrid orbitals.
d h b id bi l
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sp3d2 hybrid Orbitals
In the same way, we can achieve the six half-filledorbitals and octahedral geometry through thehybridization ofone s, three p and two d orbitals ofvalence shell into six sp3d2 hybrid orbitals.
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Hybridisation of Elements
Shape ofmolecules/ions
Hybridisation type
AtomicOrbitals
Examples
Square
Planar
dsp2 d + s + p(2) [Ni(CN)4]2-
[Pt(Cl)4]2-
Trigonal
bipyramidal
sp3d s + p(3) + d PF5, PCl5
Squarepyramidal
sp3d2 s + p(3) +d(2)
BrF5
Octahedral sp3
d2
s + p(3) + SF6 ,
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Covalent Bonding and Orbital
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Lewis structures and VSEPR do not explain why a bond
forms.
How do we account for shape in terms of quantum
mechanics?
What are the orbitals that are involved in bonding?
We use Valence Bond Theory:
Bonds form when orbitals on atoms overlap.
There are two electrons of opposite spin in the orbital overlap.
Covalent Bonding and OrbitalOverlap
Covalent Bonding and Orbital
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Covalent Bonding and OrbitalOverlap
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Molecular Shapes
Linear: Atoms lie in a straight line,bond angle is 180o
Planar triangular: Atoms are locatedon the corners of a triangle. Bondangles are all 120o
Tetrahedral: Atoms are located on
the corners of a tetrahedron. Bondangles are all 109.5o
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Trigonal bipyramid: Consists of two triangularpyramids that share a common base.
Consists of two types of bonds: a) equatorial bonds:bond angles are 120o and b) axial bonds: bonds are
180o from each other, but 90o between each equatorialbond
Octahedral: Two square pyramids sharing a commonbase. All bond angles are 90o from each other.
Molecular Shapes
Electron Domain Geometry
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Electron Domain Geometry
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Molecular Shape and Molecular
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Molecular Shape and MolecularPolarity
Molecular Shape and Molecular
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o ecu a S ape a d o ecu aPolarity
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l d i
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Electron domains
Regions in space where groups ofelectrons can be found
Two types of electron domains:a) bonding: Electrons that are
involved between pairs of atoms
b) Nonbonding: contains valenceelectrons bonded to a single atom
B di d i
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Bonding domains
All the electrons within a given
single, double or triple bond areconsidered to be in the samedomain
N b di d i
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Nonbonding domains
These can consist of either alone (unshared) pair of
electrons, or a single unpairedelectron
Nonbonding domains affect the
shape of the molecule
Molecular shapes with four
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Molecular shapes with fourdomains
Tetrahedral
Trigonal pyramidalBent
Molecular shapes with five
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Molecular shapes with fivedomains
Trigonal bipyramid
Distorted tetrahedral(seesaw)
T-shaped
linear
Molecular shapes with six
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Molecular shapes with sixdomains
Octahedral
Square pyramidalSquare planar
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Polarity of polyatomic
molecules/ions
If all the atoms attached to the
central atom are not the same,or there are lone pairs in thecentral atom, the molecule is
usually polar.
M l l Sh
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A molecule is usually non-polar if
a) all the bonds are non-polar or
b) There are n lone pairs in the valenceshell of the central atom and all atomsattached to the central atom are thesame
Molecular Shapes
H d B di
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Hydrogen Bonding
Hydrogen bond is formed when a hydrogen atom findsitself between two highly electronegative atoms suchas F,O,N.
It may be intermolecular (existing between two ormore molecules of the same or different substances)or intramolecular (present within the same molecule).
Hydrogen bonds have a powerful effect on thestructure and properties of many compounds
VSEPR Model-HydrogenBonding
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The Effect of Nonbonding Electrons
By experiment, the H-X-H bond angle decreases on
moving from C to N to O:
Since electrons in a bond are attracted by two nuclei, they do
not repel as much as lone pairs.
Therefore, the bond angle decreases as the number of lone pairs
increases
Bonding
104.5O
107O
NH
HH
C
H
HHH
109.5O
OHH
VSEPR Model
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Shapes of Larger Molecules In acetic acid, CH3COOH, there are three central atoms.
VSEPR Model
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Ionic Bonding
Resonance Structures
VSEPR
Basic Shapes
3-D Notation
Hybridization
Molecular Geometries
Octet Rule Polar Molecules
Lewis Structures Covalent Bonding
Types of Bonds
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