George Mason University

1
George Mason University General Chemistry
211 Chapter 11 Theories of Covalent
Bonding Acknowledgements Course Text Chemistry
the Molecular Nature of Matter and Change, 6th
edition, 2011, Martin S. Silberberg,
McGraw-Hill The Chemistry 211/212 General
Chemistry courses taught at George Mason are
intended for those students enrolled in a science
/engineering oriented curricula, with particular
emphasis on chemistry, biochemistry, and biology
The material on these slides is taken primarily
from the course text but the instructor has
modified, condensed, or otherwise reorganized
selected material.Additional material from other
sources may also be included. Interpretation of
course material to clarify concepts and solutions
to problems is the sole responsibility of this
instructor.
2
Molecular Structure - Summary

• Atomic theory
• Molecular Weight (MW) Neutrons Protons
• Mass, Atomic Mass units, Law of Definite
  Proportions
• Moles, Chemical Equations, Stoichiometry
• Gas Laws, Thermodynamics (reaction energy)
• Quantum Theory waves vs particles,
  electronic structure of atoms
  energy absorption,
  emission electronic
  energy levels quantum
  numbers, electron shells
• Periodicity orbital diagrams
  Pauli exclusion principle
  Aufbau Principle for populating subshells

3
Molecular Structure - Summary

• Bonding Valence electrons
  Periodic table Ionic Bonds
  Covalent Bonds
  Electronic Configuration Lattice
  Energy, Born-Haber cycle, Bond energy
• Geometry Lewis diagrams
  Resonance, Octet Rule Formal
  Charge (valence electrons
  unbonded electrons
  ½ bonded
  electrons)
• Valence-Shell Electron Pair Repulsion Model
  (VSEPR) Molecular Notation AXaEb Xa
  Bonding pairs Eb Nonbonding pairs
  sum(a b) determines geometry (linear,
  tetrahedral) if b gt 0 molecule may form
  dipole (polar)

4
Bond Theories

• Quantum Numbers Electron Configuration
• In Chapters 8 9 an electron was defined as a
  unique set of 4 quantum numbers
• The first 3 quantum numbers (n, l, ml) defined an
  atomic orbital, which could contain a maximum of
  2 electrons (1/2 -1/2 spin (ms))
• Each orbital (s, p, d) has a unique shape
• spherical (s), dumbell(p), pear shaped(d)
• All of the orbitals defined by a unique set ofn,
  l, ml quantum numbers, have the same energy

5
p orbitals
s orbital
d orbitals
6
Bonding Theories

• Valence Bond (VB) theory is one of two basic
  theories, along with Molecular Orbital (MO)
  theory, that were developed to use the methods of
  quantum mechanics to explain chemical bonding
• Valence Bond Theory is a chemical bonding theory
  that explains the bonding between two atoms
  caused by the overlap of the half-filled atomic
  orbital from each atom
• It focuses on how the atomic orbitals of the
  dissociated atoms combine to give individual
  chemical bonds when a molecule is formed
• The two atoms from the bonding atoms share each
  other's unpaired electron to form a filled
  orbital to form a hybrid orbital and bond
  together.

7
Bonding Theories

• Molecular Orbital (MO) theory is a method for
  determining molecular structure in which
  electrons are not assigned to individual bonds
  between atoms, but are treated as moving under
  the influence of the nuclei in the whole molecule
• In this theory, each molecule has a set of
  molecular orbitals, in which it is assumed that
  the molecular orbital wave function ?j may be
  written as a simple weighted sum of the n
  constituent atomic orbitals
• A given Atomic Orbital (s, p, d) takes the form
  of a subset of Molecular Orbitals
• ? bonding ?? antibonding bonds and
• ? bonding ?? antibonding bonds
• Each has its own energy
• Molecular Orbital orbitals cover the whole
  molecule

8
Valence Bond Theory

• Valence Bond Theory is an attempt to explain the
  Covalent bond from a Quantum Mechanical view
• All orbitals of the same type (s, p, d, f) have
  the same energy
• According to this theory, a bond forms when two
  atomic orbitals (s/s s/p p/p) overlap
• The space formed by the overlapping orbitals has
  a capacity for two electrons that have opposite
  spins, 1/2 -1/2 (exclusion principle)
• Note Each orbital forming the bond has at least
  one unfilled slot to accommodate the electron
  being shared from the other bonding orbital
• The bond strength depends on the attraction of
  the nuclei for the shared electrons

9
Valence Bond Theory

• Valence bond theory (cont)
• The greater the orbital overlap, the stronger
  (more stable) the bond
• The extent of the overlap depends on the shapes
  and directions of the orbitals
• An s orbital is spherical, but p and d orbitals
  have more electron density in one direction than
  in another
• Whenever possible, a bond involving p or d
  electrons will be oriented in the direction that
  maximizes overlap

10
Valence Bond Theory
Hydrogen, H2 1s1
Hydrogen Fluoride, HF He2s22p5 To maximize
overlap, half-filled H 1s and F 2p orbitals
overlap along the long axis of the 2p orbital
Fluorine, F2 He 2s22p5 In F2, the
half-filled 2 px orbital on one F atom points end
to end toward the half-filled 2px of the other F
to maximize overlap
11
Hybrid Orbitals

• One might expect the number of bonds formed by an
  atom would equal its unpaired electrons
• Chlorine, for example, generally forms one bond
  as it has one unpaired electron - 1s22s22p5
• Oxygen, with two unpaired electrons, usually
  forms two bonds - 1s22s22p4
• However, Carbon, with only two unpaired
  electrons, generally forms four (4) bonds
• C (1s22s22p2) He 2s22p2
• The four bonds come from the 2 (2s) paired
  electrons and the 2 (2p) unpaired electrons
• For example, Methane, CH4, is well known
• The uniqueness of these bonds is described next

12
Hybrid Orbitals

• Linus Pauling proposed that the valence atomic
  orbitals in a molecule are different from those
  of the isolated atoms forming the molecule
• Quantum mechanical computations show that if
  specific combinations of orbitals are mixed
  mathematically, new atomic orbitals are
  obtained
• The spatial orientation of these new orbitals
  lead to more stable bonds and are consistent
  with observed molecular shapes
• These new orbitals are called
• Hybrid Orbitals

13
Hybrid Orbitals

• Types of Hybrid Orbitals
• Each type a unique geometric arrangement
• The hybrid type is derived from the number of s,
  p, d atomic orbitals used to form the Hybrid

14
SP Hybrid Orbitals

• SP Hybridization
• 2 electron groups surround central atom
• Linear shape, 180o apart
• VB theory proposes the mixing of two
  nonequivalent orbitals, one s and one p, to
  form two equivalent sp hybrid orbitals
• Orientation of hybrid orbitals extend electron
  density in the bonding direction
• Minimizes repulsions between electrons
• Both shape and orientation maximize overlap
  between the atoms

15
sp Hybrid Orbitals
Ex BeCl2 The Be-Cl bonds in BeCl2 are neither
spherical (s orbitals) nor dumbell (p
orbitals) The Be-Cl bonds have a hybrid shape In
the Beryllium atom the 2s orbital and one of the
2p orbitals mix to form 2 sp hybrid orbitals Each
Be Hybrid sp orbital overlaps a Chlorine 3p
orbital in BeCl2
16
sp2 Hybridization

• sp2 - Trigonal Planar geometry (Central atom
  bonded to three ligands)
• The three bonds have equivalent hybridized shapes
• The sp2 hybridized orbitals are formed from
• 1 s orbital and 2 p orbitals
• Note Of the 4 orbitals available (1 s 3 p)
  only the s orbital and 2 of the p orbitals are
  used to form hybrid orbitals
• Note Unlike electron configuration notation,
  hybrid orbital notation uses superscripts for the
  number of atomic orbitals of a given type that
  are mixed, NOT for the number of electrons in the
  orbital, thus,
• sp2 (3 orbitals), sp3 (4 orbitals), sp3d (5
  orbitals)

17
sp2 Hybridization
Hybrid Orbital Diagram
BF3
The 3 B-F bonds are neither spherical nor
dumbell shaped They are all of identical shape In
Boron, the 2s orbital and two of the 2p
orbitals mix to form 3 sp2 hybrid orbitals, each
containing one the 3 total valence electrons Each
of the Boron hybrid sp2 orbitals overlaps with a
2p orbital of a Fluorine atom
Boron (B) 1s22p1 Forms 3 sp2 hybrid orbitals
BF3
18
sp3 Hybrid Orbitals

• sp3 (4 bonds, thus, Tetrahedral geometry)
• The sp3 hybridized orbitals are formed from
• 1 s orbital and 3 p orbitals
• Example
• Carbon is the basis for Organic Chemistry
• Carbon is in group 4 of the Periodic Chart and
  has 4 valence electrons 2s22p2
• The hybridization of these 4 electrons is
  critical in the formation of the many millions of
  organic compounds and as the basis of life as we
  know it
• The following slides show 3 different forms of
  the electronic structure and explains why the
  hybridized form reflects the observed structure
  of organic compounds

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SP3 Hybrid Orbitals
2p
2p
This structure implies different shapes and
energies for the s and p bonds in carbon
compounds. Observations indicate that all fours
bonds are equivalent
2s
2s
Energy
1s
1s
C atom (ground state)
C atom (promoted)
20
SP3 Hybrid Orbitals

• One bond on Carbon would form using the 2s
  orbital while the other three bonds would use 3
  2p orbitals
• This does not explain the fact that the four
  bonds in CH4 appear to be identical
• Valence bond theory assumes that the four
  available atomic orbitals (2s22p2) in carbon
  combine to make four equivalent hybrid orbitals

21
Hybrid Orbitals

• Hybrid orbitals are orbitals used to describe
  bonding that is obtained by taking combinations
  of atomic orbitals of an isolated atom
• In the case of Carbon, one s orbital and three
  p orbitals, are combined to form 4 sp3 hybrid
  orbitals
• The carbon atom in a typical sp3 hybrid structure
  has 4 bonded pairs and zero unshared electrons,
  therefore, Tetrahedral structure
• AXaEb (a b) 4 0 AX4
• The four sp3 hybrid orbitals take the shape of a
  tetrahedron

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Hybridization of Carbon in CH4
4 sp3 orbitals formed
2p
sp3
sp3
C-H bonds
Energy
1s
1s
1s
C atom (ground state)
C atom (hybridized state)
C atom (in CH4)
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Spatial Arrangement ofsp3 Hybrid Orbitals
Shape of sp3 hybrid orbital different than either
s or p
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sp3d Hybrid Orbitals

• sp3d (5 molecules, thus, Trigonal Bypyramidal
  geometry)
• Molecules with central atoms from Period 3 or
  higher, can utilize d orbitals in the formation
  of hybrid orbitals
• The sp3d hybridized orbitals are formed from
• 1 s orbital, 3 p orbitals, 1 d
  orbital
• PCl5
• AXaEb
• AX5E0
• hybrid orbitals 5 (sp3d)

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SP3d Hybrid Orbitals
Hybridized Orbital Diagram for PCl5

• 5 equivalent (hybrid) orbitals are required
• The one 3s orbital, the 3 3p orbitals and one of
  the unused 3d orbitals of the Phosphorus atom mix
  to form the 5 sp3d hybrid orbitals
• The remaining 4 empty 3d orbitals (unhybridized)
  are not used

26
Diagrams of Hybrid Orbitals Showing their Spatial
Arrangements
27
Hybrid Orbitals

• To obtain the bonding description of any atom in
  a molecule, you proceed as follows
• Write the Lewis electron-dot formula for the
  molecule
• From the Lewis formula, use the VSEPR theory to
  determine the arrangement of electron pairs
  around the central atom, i.e., the geometry
• From the geometric arrangement (AXaEb) of the
  electron pairs, obtain the hybridization type
• Assign valence electrons to the hybrid orbitals
  of this atom one at a time, pairing only when
  necessary
• Form bonds to the central atom by overlapping
  singly occupied orbitals of other atoms with the
  singly occupied hybrid orbitals of the central
  atom

28
Oxygen Atom Bonding in H2O
4 sp3 Hybridized Orbitals
??
H
H
O
??
a b 2 2 4 Tetrahedral AX2E2 bent
1s
OCentral Atom (ground state)
O atom (hybridized state)
O atom (in H2O)
29
Practice Problem

• What hybrid orbitals of Sulfur are involved in
  the bonding in Sulfur Trioxide (SO3)?
• a. sp
• b. sp2
• c. sp3
• d. sp2d
• e. sp3d2
• Ans b

Total Valence e- - 3 x 6 6 24 Bonded
Pairs 3 x 2 6 Distribute e-
about O atoms 3 x 6 18 Unshared e-
about S atom 24 - 6 -18 0 Move e- pairs
from Cl to S to form alternative forms of SO3
Compute formal charge on S select form with
least formal charge (D) AXaEb 3 0 3 AX3
(trigonal Planar) 3 O-S hybridized orbitals are
required one s orbital blended with
2 p orbitals (sp2)
30
Sulfur Trioxide Hybrid Orbitals
3p
3p
3p
VSEPR AX3 Trigonal Planar 3 sp2 orbitals
required
sp2
sp2
S atom (in SO3)
S atom (ground state)
S atom (hybridized state)
31
Nitrogen Atom Bonding in NH3
4 sp3 orbitals required
sp3
sp3
N-H bonds
lone pair
a b 3 1 4 Tetrahedral AX3E1 trigonal pyramid
al
Tetrahedral
1s
N atom (ground state)
N atom (hybridized state)
N atom (in NH3)
32
Multiple Bonds

• Types of Covalent Bond Orbital Overlap
• Orbitals can overlap two ways
• Side to Side or End to End
• Two types of Covalent Bonds
• Sigma Bonds (C-C)
• pi (?) Bonds (CC)
• Multiple Bonds

Ethane Tetrahedral (both carbons)
Ethylene Trigonal planar(both carbons)
Acetylene Linear (both carbons)
109.5o
120o
sp3
sp2
double bond acts as single electron group
triple bond acts as single electron group
33
Multiple Bonds

• End-to-End overlap Sigma Bonds
• The C C bond in Ethane (C2H6) involves overlap
  of 1 sp3 orbital from each carbon
• Each of the six (6) C H bonds involves the
  overlap of a Carbon sp3 and a Hydrogen 1 s
  orbital
• All bonds involve overlap of one end of orbital
  with the end of the other orbital
• The bond formed from end-to-end overlap is called
  a sigma bond (symbol - ?)

34
Multiple Bonding

• According to Valence Bond theory, one hybrid
  orbital is needed for each bond (whether a single
  or multiple) and for each lone pair
• For example, consider the molecule
• Ethene (or Ethylene)

35
Multiple Bonding

• Each Carbon atom is bonded to three other atoms
  and no lone pairs, which indicates the need for
  three hybrid orbitals
• This implies AX3E0 (Trigonal) sp2 hybridization
• 1 2s 2 2p orbitals
• The third 2p orbital is left unhybridized and
  lies perpendicular to the plane of the trigonal
  sp2 hybrids
• The following slide represents the sp2
  hybridization of the Carbon atoms

36
Multiple Bonding
(unhybridized)
2p
2p
sp2
Energy
C atom (ground state)
C atom (hybridized)
37
Multiple Bonding

• Each carbon atom is sp2 hybridized
• Each of the carbon atoms 4 valence electrons
  fill ½ its 3 sp2 orbitals and its unhybridized 2p
  orbital, which lies perpendicular to sp2 plane
• Two sp2 orbitals of each carbon form C H sigma
  (?) bonds by overlapping the 1 s orbitals of the
  two H atoms
• The 3rd sp2 orbital of one carbon forms a C C
  (?) bond with the sp2 orbital of the other carbon
  with end-to-end overlap
• A pi (?) bond is formed when the two unhybridized
  2p orbitals (one from each carbon) overlap
  side-to-side, forming two regions of electron
  density, one above and one below the ?-bond axis
• A double bond always consists of
• one ?-bond and one ? bond

38
Multiple Bonding

• Two of the sp2 hybrid orbitals of each carbon
  overlap end-to-end with the 1s orbitals of the 2
  hydrogen atoms forming a sigma bond
• The remaining sp2 hybrid orbital, one on each
  carbon, overlap end-to-end to form a sigma bond

39
Multiple Bonding

• The remaining unhybridized 2p orbitals, one on
  each of the carbon atoms, overlap side-to-side,
  one on top of the sigma bond and one on the
  bottom of the sigma bond, forming a p bond

The carbon-carbon double bond is described as one
s bond and one ? bond
The two electron pairs in a double bond act as a
single electron group The electron pairs do not
repulse each other because each electron pair
occupies a distinct orbital, a specific region of
electron density, thus repulsions are reduced
40
Practice Problem

• Use valence bond theory to describe the bonding
  in CO2
• Ans
• 1. Draw Lewis structure
• 2. Determine hybridization
• 3. Draw diagram of hybrid atomic orbitals
• 4. Pair electrons (O) with hybrid C orbitals
  forming sigma bonds
• 5. Pair electrons (O) with unpaired p electrons
  in C atom to form pi (?) bonds

Cont on next slide
41
Practice Problem (Cont)
Use valence bond theory to describe the bonding
in CO2
2 bonding pairs 0 non-bonding pairs AXaEb a
b 2 0 2 (Linear) Hybridization sp (2
hybrid orbitals required
42
Molecular Orbital (MO) Theory

• Molecular Orbital (MO) theory is a theory of the
  electronic structure of molecules in terms of
  molecular orbitals, which may spread over several
  atoms or the entire molecule
• MO theory explains the observed and computed
  energy differences among orbitals, which Valence
  Bond theory does not
• As atoms approach each other and their atomic
  orbitals overlap, molecular orbitals (MO) are
  formed
• Note Only outer (valence) Atomic orbitals (AO)
  interact enough to form Molecular Orbitals (MO)
• Electron motions are complex making solutions to
  the Schroedinger equation approximations
• Mathematically, the combination of atomic
  orbitals to form molecular orbitals involves
  adding or subtracting atomic wave functions

43
Molecular Orbital (MO) Theory

• Adding Wave Functions
• Forms a Bonding (?) molecular orbital (MO)
• Region of high electron density between nuclei
• Electron charge between nuclei is dispersed over
  a larger area than in atomic orbitals (AO)
• MO orbital energy is lower than in the AO because
  of the reduction in electron repulsion
• Bonding MO is more stable than AO

44
Molecular Orbital (MO) Theory

• Subtracting Wave Functions
• Forms a Nonbonding (?) molecular orbital
• The node between the nuclei has most of the
  electron density outside the node with very
  little density (zero) between the nuclei
• Thus, the electrons do not shield one nuclei from
  the other resulting in increased nucleus-nucleus
  repulsion
• Therefore, the antibonding MO has a higher energy
  than the corresponding atom orbitals (AO)
• When the antibonding orbital is occupied, the
  molecule is less stable than when the orbital is
  not occupied

45
Molecular Orbital Theory

• Example The bonding of two Hydrogen atoms
• s1s (bonding) molecular orbital is formed
• s1s (antibonding) molecular orbital is formed
• The following slide illustrates the relative
  energies of the molecular orbitals (MO) compared
  to the original atomic orbitals (AO)
• Because the energy of the two electrons in the
  bonding orbital is lower than the energy of the
  individual atoms, the molecule is stable

46
Molecular Orbital Theory
Atomic orbital
Molecular Orbital
Atomic orbital
H atom
H atom
H2 molecule
s1s
1s
1s
s1s
More Stable
47
Bonding and Antibonding Orbitals from 1s Hydrogen
Atom Orbitals
48
Bond Order

• The term bond order refers to the number of
  electron pairs shared between two atoms
• The bond order of a diatomic molecule is defined
  as one-half the difference between the number of
  electrons in bonding orbitals, nb, and the number
  of electrons in antibonding orbitals, na

For example, try H2 and He2. Determine bond
orders
49
Bond OrderH2
H 1s1
Bond Order BO ½(2) (0) ½2 1
50
Bond OrderHe2
He 1s2
Bond Order BO ½(2) (2) ½0 0
51
Diatomic HomonuclearSubstances in 2p period

• The 2p orbitals can overlap in two ways
• End-to-End gives ?2p and ?2p molecular orbitals
  (MO)
• Side-to-Side gives a pair of ?2p and ?2p MOs

52
Diatomic HomonuclearSubstances in 2p period

• The order of MO energy levels, whether bonding or
  nonbonding, is based on the AO (atomic orbital)
  energy levels and on the mode of the p orbital
  combination
• MOs formed from 2s orbitals are lower in energy
  than 2p orbitals because 2s AOs are lower in
  energy than 2p AOs
• Bonding MOs are lower in energy than antibonding
  MOs
• ?2p is lower in energy than ?2p
• ?2p is lower in energy than ?2p

53
Diatomic HomonuclearSubstances in 2p period

• Atomic p orbitals (AO) can interact more
  extensively End-to-End than Side-to-Side
• Thus, ?2p MO is lower in energy than ?2p
• The destabilizing effect of the ?2p MO is
  greater than that of the ?2p MO
• The energy order for MOs derived from 2p orbitals
  is
• ?2p lt ?2p lt ?2p lt ?2p

Most Stable
Least Stable
54
Diatomic HomonuclearSubstances in 2p period

• Several factors are involved in the relative
  energies of the various molecular orbitals (MO)
• Bond length
• Bond energy
• Bond order
• Magnetic properties
• Electron valence shell configuration

55
Diatomic HomonuclearSubstances in 2p period

• Factors that affect the MO energy level order
• There are three (3) mutually perpendicular 2p
  orbitals in each atom of a diatomic molecule (2px
  2py 2pz)
• When the 6 p orbitals (3 from each element)
  combine, only one orbital from each element can
  interact end-to-end forming a ? (bond) and a
  ? (antibonding) Molecular Orbital (MO)
• The other two pairs of orbitals interact side to
  side to form two ? MOs and two ? MOs of
  the same energy giving the expected MO diagrams
  for the p-block Period 2 homonuclear diatomic
  molecules

56
Diatomic HomonuclearSubstances in 2p period
End to end
Side to side
57
Diatomic HomonuclearSubstances in 2p period

• Other factors influence the MO energy level order
• s" and p AOs can be similar in energy or
  differ considerably in energy, which determines
  whether the orbitals mix or dont mix
• O, F, Ne atoms are relatively small and
  electron repulsions raise the energy of 2p
  orbitals high enough above 2s orbitals to
  minimize orbital mixing
• Atoms, such as B, C, N, are larger in size and
  the s and p AOs have less electron repulsion
  and the energy difference between 2s 2p is
  less, resulting in mixing of the s p
  orbitals

58
Diatomic HomonuclearSubstances in 2p period

• This smaller difference in energy of the 2p 2s
  orbitals in the P, C, N atoms permits some mixing
  of the orbitals between the 2s orbital of one
  atom and the endon of the 2p orbital of the
  other atom
• This orbital mixing
• lowers the energy of the ?2s and ?2s MOs and
• raises the energy of the ?2p and ?2p MOs
• The ? MOs are not affected
• The effect of the mixing is the reversal of the
  ?2s and ?2p MOs
• ?2p lt ?2p lt ?2p lt ?2p
• The next slide illustrates these differences

59
Diatomic HomonuclearSubstances in 2p period
Without2s -2p mixing
With2s -2p mixing

• Effect of Mixing
• The ? MOs are not affected
• Reversal of the ?2s and ?2p MOs

MOenergy levels O2, F2, Ne2
MOenergy levels B2, C2, N2
60
Diatomic HomonuclearSubstances in 2p period
MO occupancy and molecular properties for B2
through Ne2 Bond energy and bond order are
inversely related to bond length Orbitals with
unpaired electrons are paramagnetic and are
attracted to an external magnetic field Note
reversal of ?2p ?2p energy levels
61
Diatomic Homonuclearsubstances in 2p period
The arrows show the occupation of molecular
orbitals by the valence electrons in N2 Bond
Order BO ½(Nb - Na) BO ½(242) - (2)
½8-2 3
62
Sample Problem
Which of the following species has a bond order
of 2.5? a. O2 b. O2 c. O22 d. O2 e. NO
Ans d or e for NO BO
½(242)-(21) ½(8-3) 2.5
Nitrogen 1s22s22p3
Oxygen 1s22s22p4
Cont on next slide
63
Sample Problem
Which of the following species has a bond order
of 2.5? a. O2 b. O2 c. O22 d. O2 e. NO
Ans d or e for O2 BO ½(242)-(21)
½(8-3) 2.5
64
Diatomic HeteronuclearSubstances in 2p period

• Heteronuclear diatomic molecules are composed of
  two different atoms HF NO , etc.
• Heteronuclear molecules have Asymmetric MO
  diagrams
• Atoms with greater effective nuclear charge
  (Zeff) draw their electrons closer to the
  nucleus, thus, they have higher electronegativity

65
Diatomic HeteronuclearSubstances in 2p period

• Example Hydrogen Fluoride(HF)
• Higher effective nuclear charge of Fluorine
  nucleus holds electrons more tightly than H
  (proton nucleus)
• All occupied atomic orbitals of F have lower
  energy than the 1s orbital of Hydrogen
• The Hydrogen 1s orbital reacts with the F 2p
  orbitals
• Only one of the 3 F 2p orbitals, 2pz leads to
  end-to-end overlap with Hydrogen 1s orbital
  producing a ? MO
• The other two p orbitals (2px 2py) are not
  involved in the bonding and are called
  nonbonding MOs)

Hydrogen
Fluorine
2s AO of F is not shown
66
Diatomic HeteronuclearSubstances in 2p period

• Example
• Bonding in Heteronuclear Nitrogen Monoxide, NO
• Highly reactive compound because it has a lone
  electron
• Two possible Lewis Structures
• Not clear where lone electron resides (N or O)
• Lower Formal charge on N suggests structure I
• MO theory predicts electron resides closer to the
  Nitrogen atom
• Measured bond energy suggests bond order higher
  than 2

I
II
Formal charge ?
? Formal charge
67
Diatomic HeteronuclearSubstances in 2p period

• Example Nitrogen Monoxide
• The 11 valence electrons of NO fill MOs in order
  of increasing energy, leaving the lone electron
  in one of the ?2p orbitals
• Atomic orbitals of O have lower energy than those
  of N Oxygen is more electronegative Nitrogen
• The 8 bonding electrons and the 3 non-bonding
  electrons give a bond order of ½(8-3) 2.5
• The bonding electrons lie in MOs closer in energy
  to the AOs of the Oxygen atom
• The lone unshared electron occupies an
  anti-bonding orbital (?2p)
• Because this orbital receives a greater
  contribution from the 2p orbitals of the N atom,
  it resides closer to the Nitrogen atom

Nitrogen
Oxygen
68
Molecular Orbital Template(without 2s 2p
mixing)
?2p
?2p
2p
2p
?2p
?2p
?2s
2s
2s
?2s
69
Molecular Orbital Template(with 2s 2p mixing)
?2p
?2p
2p
2p
?2p
?2p
?2s
2s
2s
?2s
70
Equation Summary
71
Equation Summary

• VESPR Model Molecular Notation
• AXaEb
• A The Central Atom (Least Electronegative
  atom)
• X The Ligands (Bonding Pairs)
• a The Number of Ligands
• E Non-Bonding Electron Pairs
• b The Number of Non-Bonding Electron Pairs
• Double Triple Bonds count as a single
  electron pair
• The Geometric arrangement is determined by
• sum (a b)