Transcript First Law of Thermodynamics:

Ch 2 Structure & Bonding
An amazing thing about the universe - It works in a way that
sometimes when things come together, they stick…
• Sections 2.4-2.8
H
H
H
H
“spin-pairing”
Valence Bonding (Localized) vs.
Molecular Orbital Theory (Delocalized)
Orbital Overlap (Localized Bonding)
Bonding orbitals are constructed by combining atomic
orbitals from adjacent atoms.

In VB Theory, usually only worry about the valence electrons
From Quantum Mechanics: orbitals can add or
subtract; therefore constructive or destructive
interference is possible
Localized bonding… H2
Orbital Overlap: As two H atoms approach, the overlap
of their 1s atomic orbitals increases. The wave amplitudes
add, generating a new orbital with high electron density
between the nuclei.
Sigma bonding: The s bond is actually symmetry notation
• means bonding is directed along internuclear axis
• sigma bonds have a C∞ axis
Pi bonding: The p bond is also found in symmetry notation
• means bonding is above & below the internuclear axis
PH3
Phosphine is a colorless, highly toxic gas
with bond angles of 93.6°. Describe the
bonding in PH3.
Oh shoot… What about methane?
What is the electron
configuration of methane?
s and p hybridization
Promotion: excitation of an electron to a higher energy orbital
in the course of bond formation – not real exactly
Hybridization: mathematical mixing (linear combinations) of
valence atomic orbitals to achieve new equal energy
degenerate orbitals
Methane hybridization
CH4 is tetrahedral
Therefore, 2s and the 2px, 2py, & 2pz must hybridize

new orbitals are called sp3
An inner atom with a steric number of 4 has tetrahedral
electron group geometry and can be described using sp3
hybrid orbitals.
General Features of Hybridization
1.
2.
3.
4.
The # of valence orbitals generated by hybridization
equals the # of valence AOs participating in
hybridization.
The steric number of an inner atom uniquely
determines the number and type of hybrid orbitals.
Hybrid orbitals form localized bonds by overlap with
atomic orbitals or with other hybrid orbitals.
There is no need to hybridize orbitals on outer atoms,
because atoms do not have limiting geometries.
The bonds formed by all other outer atoms can be described
using valence p orbitals.
Isolobal: When analogous fragments on differing
molecules have closely similar bonding patterns
Molecular Orbital Theory:
When overlapping AOs just won’t cut it.
So far, bonding has been described as overlapping AOs or hybrid
orbitals
However, *all* electrons from each bonding atom feel the presence of
the others
Bonding in the diatomic molecules second row elements can be
explained in two ways.

Localized Bonding (LB) Theory, a.k.a. Valence Bond Theory

Molecular Orbital (MO) Theory
MO Theory

assumes pure s and p AOs of the atoms in a molecule combine to
produce orbitals that are spread out, or delocalized, over several
atoms, leading to MOs

One advantage over VB Theory: correctly explains electronic
structures of molecules which do not follow Lewis Dot structure.
4 Principles of MO Theory
1st Principle: the total # of MOs produced by a set
of interacting AOs is equal to the # of interacting
orbitals
2nd Principle: the bonding MO is lower in energy
than the parent AOs & the antibonding MO is
higher in energy (LCAO)
To explain this, let’s look at the hydrogen molecule,
H2
Bonding and Antibonding in H2
Each hydrogen atom contributes a 1s orbital
These orbitals can be added or subtracted


Addition: Bonding MO (s1s)
Subtraction: Bonding MO (s1s*)
Principles of MO Theory
3rd Principle: electrons of the molecule are
assigned to orbitals of successively higher
energy according to the Aufbau principle
and Hund’s Rule


Electrons occupy the lowest energy orbitals
first.
Atoms are most stable with the highest
number of unpaired electrons
molecular orbital diagram of He2
NOW: Use a molecular orbital diagram to predict if
it is possible to form the He2+ cation.
Bond Order in MO Theory
Bond order allows us to represent the net
amount of bonding between two atoms.

The higher the bond order, the more stable the
structure
1
BO  # electrons in bonding MOs  # electrons in antibondin g MOs
2
One More Principle…
4th Principle: atomic orbitals combine to
form molecular orbitals most effectively
when the atomic orbitals are of similar
energy & symmetry

i.e. a 1s will not bond with a 2s.
Second-Row Diatomic Molecules
-
+
subtractive
additive
NOTE NUMBERING!
Orbital Mixing
In B2, the overlap of 2s
and 2pz orbitals stabilizes
ss and destabilizes sp
The amount of mixing
depends on the energy
difference between the 2s
and 2p atomic orbitals.
Mixing is largest when the
energies of the orbitals are
nearly the same
2 cases: 1) Zave ≤ 7, 2) Zave > 7 (qualitative)
Heteronuclear Diatomic Molecules
Let’s examine the MO diagram of NO
…and of CO…
Evidence for Antibonding Orbitals