2 Shapes of Atomic Orbitals for Electrons Four different kinds
of orbitals for electrons denoted s, p, d, and f s and p orbitals
most important in organic and biological chemistry s orbitals:
spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus
at middle d orbitals: elongated dumbbell-shaped, nucleus at
center
Slide 3
The Nature of Chemical Bonds: Covalent bond forms when two
atoms approach each other closely so that a singly occupied orbital
on one atom overlaps a singly occupied orbital on the other atom
Two models to describe covalent bonding. Valence bond theory,
Molecular orbital theory
Slide 4
Central Themes of Valence Bond Theory 1) Opposing spins of the
electron pair. The region of space formed by the overlapping
orbitals has a maximum capacity of two electrons that must have
opposite spins. Basic Principle of Valence Bond Theory: a covalent
bond forms when the orbitals from two atoms overlap and a pair of
electrons occupies the region between the nuclei.
Slide 5
2) Maximum overlap of bonding orbitals. The bond strength
depends on the attraction of nuclei for the shared electrons, so
the greater the orbital overlap, the stronger the bond.
Slide 6
Central Themes of Valence Bond Theory 3) Hybridization of
atomic orbitals. bonding in simple diatomic molecules: Example1: HF
(direct overlap of the s and p orbitals of isolated ground state
atoms). Example 2: CH 4 (4 hydrogen atoms are bonded to a central
carbon atom)- hybridization happens to obtain the correct bond
angles. Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms. We call
this Hybridization!
Slide 7
Slide 8
Fig. 11.1
Slide 9
What Is A Hybrid Orbital? are a type of atomic orbital that
results when two or more atomic orbitals of an isolated atom mix
(the number of hybrid orbitals on a covalentlyatomcovalently bonded
bonded atom is equal to the number of atomic orbitals used to form
the hybrid orbitals) are used to describe the orbitals in
covalently bonded atoms (hybrid orbitals do not exist in isolated
atoms), have shapes and orientations that are very different from
those of atomic orbitals in isolated atoms in a set are equivalent,
and form identical bonds (when the bonds are to a set of identical
atoms), and are usually involved in sigma bonds in polyatomic
molecules; pi bondsmolecules usually involve the overlap of
unhybridized orbitals.
Slide 10
In Hybridization there is mixing or blending of atomic orbitals
to accommodate the spatial requirements in a molecule.
Hybridization occurs to minimize electron pair repulsions when
atoms are brought together to form molecules.
Slide 11
1. sp 3 Orbitals and the Structure of Methane Carbon has 4
valence electrons (2s 2 2p 2 ) In CH 4, all CH bonds are identical
(tetrahedral) sp 3 hybrid orbitals: s orbital and three p orbitals
combine to form four equivalent, unsymmetrical, tetrahedral
orbitals (sppp = sp 3 ), Pauling (1931) 11 Types of Hybrid
Orbitals
Slide 12
Carbon: 2s 2 2p x 1 2p y 1 one electron in each of four sp
3
Slide 13
The Structure of Methane sp 3 orbitals on C overlap with 1s
orbitals on 4 H atoms to form four identical C-H bonds Each CH bond
has a strength of 436 (438) kJ/mol and length of 109 pm Bond angle:
each HCH is 109.5, the tetrahedral angle. 13
Slide 14
The sp 3 Hybrid Orbitals in NH 3 and H 2 O Fig. 11.5
Slide 15
Hybridization of Nitrogen and Oxygen Elements other than C can
have hybridized orbitals HNH bond angle in ammonia (NH 3 ) 107.3
C-N-H bond angle is 110.3 Ns orbitals (sppp) hybridize to form four
sp 3 orbitals One sp 3 orbital is occupied by two nonbonding
electrons, and three sp 3 orbitals have one electron each, forming
bonds to H and CH 3. 15
Slide 16
sp 3 Orbitals and the Structure of Ethane Two Cs bond to each
other by overlap of an sp 3 orbital from each Three sp 3 orbitals
on each C overlap with H 1s orbitals to form six CH bonds CH bond
strength in ethane 423 kJ/mol CC bond is 154 pm long and strength
is 376 kJ/mol All bond angles of ethane are tetrahedral 16
Slide 17
2. sp 2 Orbitals and the Structure of Ethylene sp 2 hybrid
orbitals: 2s orbital combines with two 2p orbitals, giving 3
orbitals (spp = sp 2 ). This results in a double bond. sp 2
orbitals are in a plane with 120 angles Remaining p orbital is
perpendicular to the plane 17
Slide 18
Bonds From sp 2 Hybrid Orbitals Two sp 2 -hybridized orbitals
overlap to form a bond p orbitals overlap side-to-side to formation
a pi ( ) bond sp 2 sp 2 bond and 2p2p bond result in sharing four
electrons and formation of C-C double bond Electrons in the bond
are centered between nuclei Electrons in the bond occupy regions
are on either side of a line between nuclei 18
Slide 19
Structure of Ethylene H atoms form bonds with four sp 2
orbitals HCH and HCC bond angles of about 120 CC double bond in
ethylene shorter and stronger than single bond in ethane Ethylene
C=C bond length 134 pm (CC 154 pm) 19
Slide 20
Fig. 11.3
Slide 21
3. sp Orbitals and the Structure of Acetylene C-C a triple bond
sharing six electrons Carbon 2s orbital hybridizes with a single p
orbital giving two sp hybrids two p orbitals remain unchanged sp
orbitals are linear, 180 apart on x-axis Two p orbitals are
perpendicular on the y-axis and the z-axis 21
Slide 22
Orbitals of Acetylene Two sp hybrid orbitals from each C form
spsp bond p z orbitals from each C form a p z p z bond by sideways
overlap and p y orbitals overlap similarly 22
Slide 23
Bonding in Acetylene Sharing of six electrons forms C C Two sp
orbitals form bonds with hydrogens 23
Slide 24
4. The sp 3 d Hybrid Orbitals in PCl 5 Fig. 11.6
Slide 25
5. The sp 3 d 2 Hybrid Orbitals in SF 6 Sulfur Hexafluoride --
SF 6 Fig. 11.7
Slide 26
Slide 27
Molecular Orbital Theory A molecular orbital (MO): where
electrons are most likely to be found (specific energy and general
shape) in a molecule Additive combination (bonding) MO is lower in
energy Subtractive combination (antibonding) MO is higher energy
27
Slide 28
Molecular Orbitals in Ethylene The bonding MO is from combining
p orbital lobes with the same algebraic sign The antibonding MO is
from combining lobes with opposite signs Only bonding MO is
occupied 28
Slide 29
29 Valence Bond Theory vs. MO Theory VB Theory begins with two
steps: hybridization (where necessary to get atomic orbitals that
point at each other) combination of hybrid orbitals to make bonds
with electron density localized between the two bonding atoms Key
differences between MO and VB theory: MO theory has electrons
distributed over molecule VB theory localizes an electron pair
between two atoms MO theory combines AOs on DIFFERENT atoms to make
MOs (LCAO) VB theory combines AOs on the SAME atom to make
hybridized atomic orbitals (hybridization) In MO theory, the
symmetry (or antisymmetry) must be retained in each orbital. In VB
theory, all orbitals must be looked at once to see retention of the
molecules symmetry.
Slide 30
Sigma and Pi Bonds
Slide 31
Difference between sigma and pi bond Sigma bond() Formed by
head to head overlap of AOs The 2 AOs that overlap are symmetrical
about the x axis joining the 2 nuclei. Has free rotation Lower
energy Only one bond can exist between two atoms( a single covalent
bond) Pi bond() Formed by side-to-side overlap of AOs No free
rotation Has higher energy One or two bonds can exist between two
atoms Have nodal plane on the molecular axis and no longer
symmetrical about the molecular axis.
Slide 32
Formation of Sigma Molecular Orbitals: 1. Overlapping of two 1s
atomic orbital Example: H 2 2. Overlapping of two px atomic orbital
+ bonding antibonding bonding antibonding
Slide 33
3. Overlapping of an s and px atomic orbitals 4. Overlapping of
px and dz 2 or dx 2 -y 2 + + bonding antibonding bonding
antibonding
Slide 34
Formation of Pi Molecular Orbitals: 1. Overlapping of two py
atomic orbitals + bonding antibonding 2. Overlapping of two pz
atomic orbitals + bonding antibonding
Slide 35
3. Overlapping of py or pz with dxz or dxy + bonding
antibonding
Slide 36
Fig. 11.10
Slide 37
Fig. 11.11
Slide 38
Comparison between sigma and pi electrons in ethylene and
acetylene 1. Pi electron are made exposed to the environment than
the sigma electrons. 2. Pi electrons are more reactive than sigma
electrons. 3. The looseness of the pi electrons in C 2 H 2 is less
than in C 2 H 4. This is the result of the greater s character in
sp hybrids as compared with the s character in sp 2 hydrid. 4. Pi
electrons in C 2 H 2 are attracted more strongly towards the
nucleus than the pi electrons in C 2 H 4. 5. The pi bonds in C 2 H
2 are more susceptible to attack by other chemical entities than in
C 2 H 4.
Slide 39
Valence Bond Theory: Overlap of Atomic Orbitals H-H 1s bond -
overlap of s orbitals H-F F: 2p + bond - overlap of s orbital and p
orbital N2N2 N: 2p 2s bond - head/head overlap of p orbitals 2
bonds - sidewise overlap of p orbitals Bonds with hybridization of
atomic orbitals: CH 4 1s H: C atom: bonds - overlap of H-s orbitals
and sp 3 orbitals Of C C: 2p 2s sp 3 sp 3 hybridization A single sp
3 orbital, each with single electron
Slide 40
sp 2 hybridization H2COH2COC: 2p 2s sp 2 p C atom 3 bond; 1
bond A single sp 2 orbital, each with single electron C H H O sp
hybridization CO2CO2 C: 2p 2s sp p O C O C atom: 2 bonds 2 bond A
single sp orbital, each with single electron
Slide 41
N2N2 N: 2p 2s bond - head/head overlap of p orbitals 2 bonds -
sidewise overlap of p orbitals N: 2p 2s sp N2N2 p with sp
hybridization with all valence electrons N N 1 bonds 2 bond N
Dinitrogren: 2 Models without hybridization with just p electrons 1
st sp orbital one with nonbonded pair 2 nd sp orbital O atom in H 2
CO (previous slide) -geometry around O is trigonal planar -
requires 3 equivalent orbitals from the O - hence, sp 2
hybridization O atom Can form 1 bond; Can form 1 bond Has 2
nonbonded pairs O: 2p 2s sp 2 p