Download Hybrid Orbitals

Lots of multisyllabic words, I know 
Overlap and Bonding
 We think of covalent bonds forming through the
sharing of electrons by adjacent atoms.
 In such an approach this can only occur when
orbitals on the two atoms overlap.
Hybridization
 Atomic orbitals (s, p, d, f) cannot adequately explain the
bonding in molecules
 Consider CH4 … how do we get the tetrahedral shape
out of the above orbitals
Hybridization
 An averaging of atomic orbitals into a new set of
“hybrid orbitals”
 Consider beryllium:
 In its ground electronic state, it would not be able to
form bonds because it has no singly-occupied orbitals.
Hybridization
 Consider beryllium:
 if it absorbs the small amount of energy needed to
promote an electron from the 2s to the 2p orbital, it can
form two bonds.
Hybrid Orbitals
 Mixing the s and p orbitals yields two degenerate
orbitals that are hybrids of the two orbitals.
 These sp hybrid orbitals have two lobes like a p orbital.
 One of the lobes is larger and more rounded as is the s
orbital.
Hybrid Orbitals
 These two degenerate orbitals would align themselves 180
from each other.
 This is consistent with the observed geometry of beryllium
compounds: linear.
Hybrid Orbitals
 With hybrid orbitals the orbital diagram for beryllium
would look like this.
 The sp orbitals are higher in energy than the 1s orbital but
lower than the 2p.
Hybrid Orbitals
 Think about Boron
 What’s its electron configuration?
 How many bonds does it normally form?
 How?
Hybrid Orbitals
…three
degenerate
sp2 orbitals.
Hybrid Orbitals
 For Carbon:
Hybrid Orbitals
…four degenerate
sp3 orbitals.
Hybrid Orbitals
 For geometries involving expanded octets on the central
atom, we must use d orbitals in our hybrids.
Hybrid Orbitals
This leads to five degenerate sp3d orbitals…
…or six degenerate sp3d2 orbitals.
Hybrid Orbitals
 Once you know the electron-domain geometry, you know
the hybridization state of the atom.
 Short cut… count the bonded electron domains… that’s the
number of hybridizations.
#
Hybridization Bond
Electron Pair
Hybrid
Angles Geometry
Orbitals
2
sp
180
Linear
3
sp2
120
Trigonal
planar
4
sp3
109
Tetrahedral
5
sp3d
120,90
Trigonal
bipyramidal
6
sp3d2
90
octahedral
Practice Problems
 Specify the electron-pair and molecular geometry for
each underlined atom in the following list. Describe the
hybrid orbital set used by this atom in each molecule
or ion.
a) BBr3
b) CO2
c) CH2Cl2
d) CO3-2
Practice Problems Answers
Electron-Pair
Geometry
Molecular
Geometry
Hybrid
Orbital Set
a) Trigonal Planar
Trigonal Planar
Linear
tetrahedral
trigonal planar
sp2
sp
sp3
sp2
b) Linear
c) Tetrahedral
d) Trigonal planar
Types of Covalent Bonds
 According to the valence bond theory, bond formation
requires that two orbitals on adjacent atoms to
overlap….
 So what about multiple bonds…
 Draw C2H4

Is there room in between the two atoms for two pairs of
electrons to share space?
 NO
Types of Covalent Bonds
 Sigma (σ)
 Formed from the overlap of hybrid orbitals. Electron
density along the internuclear axis… between the
centers of the atoms…
Types of Covalent Bonds
 Pi (π)
 Formed from the overlap of
unhybridized p-orbitals.
Electron density is above
and below the internuclear
axis, between the atomic
centers
Types of Covalent Bonds
 Draw Aceylene:C2H2
 In triple bonds, two sp
orbitals form a  bond
between the carbons, and
two pairs of p orbitals
overlap in  fashion to form
the two  bonds.
 Or there is always one sigma
bond… anything more is a pi
bond
Molecular Orbital Theory
ATOMS
Electrons in
Orbitals
MOLECULES
Molecular Orbital Theory:
Atomic orbitals from all atoms
combine to form new orbitals
resulting in known geometries
Hybridization: Atomic orbitals
from single atoms combine to
form new orbitals resulting in
known geometries
s, p, d, f, etc
VESPER: Electrons repel;
moleular shapes result.
Experimental
Geometries
Linear, trigonal
planar, tetrahedral
Molecular Orbital Theory
 Molecular orbitals are formed by combinations of
atomic orbitals from different atoms.
 There are bonding and antibonding orbitals
 Without significant electrons density between them, the
nuclei will repel each other… creating antibonding
orbitals
Molecular Orbital Theory
 First Principle of MO Theory
 the total number of molecular orbitals is always equal to
the total number of atomic orbitals contributed by the
atoms that have combined.
 Second Principle of MO Theory
 The bonding molecular orbital is lower in energy than
the parent orbitals, and the antibonding orbital is higher
in energy
Molecular Orbital Theory
 Third Principle of MO Theory
 Electrons of the molecule are assigned to orbitals of
successively higher energy
 Fourth Principle of MO Theory
 Atomic orbitals combine to form molecular orbitals most
effectively when the atomic orbitals are of similar energy
Molecular Orbital Theory
# of molecular orbitals = # of
atomic orbitals
Equal # of bonding and
antibonding orbitals
Antibonding orbitals ALWAYS
higher in energy than its BONDING
COUNTERPART
Basically, it means that these guys
will fill in from the bottom up… from
lower energy to higher energy…
Molecular Orbital Theory
 Can use this to calcualte bond order
 Bond Order = ½ [(# e- in bonding orbitals) –
(# of e- in antibonding orbitals)]
What is the bond order of Ne2+?
Bond order = (8-7)/2 = ½