Bonding and Molecular
Structure
Chapters 8 & 9
Molecular Structure
• The way atoms are arranged in space
is referred to as molecular
structure, and bonding describes the
forces that hold adjacent atoms
together.
Valence Electrons
• outermost electrons are called valence
electrons; inner electrons are called core
electrons
• Lewis Dot Structures can be used to show
valence electrons, usually main group
elements (group # = valence).
– All noble gases (except He) have 8 valence
electrons – the octet rule!
8.1 Chemical Bond
Formation
• Lewis Dot Structures can be used to
depict ionic and covalent bonds
– Ionic bonds are the attractive forces
between the positive and negative ions.
• One or more valence electrons is
transferred!
– Covalent bonds is the sharing of valence
electrons between atoms.
Bonding in Ionic
Compounds
• Tendancy to achieve a noble gas
configuration by gain or loss of
electrons is important to the
chemistry of main group elements…
– formation of the ion pair must outweigh
the formation of the individual ions.
8.2 Covalent Bonding
• Most compounds are covalently bonded,
especially carbon compounds.
• Number of bond pairs – the Octet Rule
– Predict # of bonds by counting the number of
unpaired electrons in a Lewis structure
– A dash represents a bond pair, a colon is used
to represent a lone pair (nonbonding electrons)
Octet Rule
• each atom (except H) achieves a
noble gas configuration; the element
bonds to have 8 valence electrons
• The tendency of molecules and
polyatomic ions to have structures in
which 8 electrons surround each
atom is known as the octet rule.
Drawing Lewis
Structures
• Decide on central atom (usually
lowest electron affinity):
– H is always a terminal atom. It is
ALWAYS connected to only one other
atom!
– LOWEST electronegativity is central
atom in molecule. (often C,N,P,S)
Drawing Lewis
Structures
• Determine total number of valence
electrons:
– In neutral molecules, total # of valence
electrons by adding up group #s of the
elements. FOR IONS add for negative and
subtract for positive charges. Divide by two to
get the number of electron pairs.
• Place one pair of electrons, a σ bond,
between each pair of bonded atoms.
Drawing Lewis
Structures
• Subtract from the total the number of
bonds you just used.
• Place lone pairs about each terminal atom
(except H) to satisfy the octet rule. Left
over pairs are assigned to the central
atom. If the central atom is from the 3rd
or higher period, it can accommodate more
than four electron pairs.
Drawing Lewis
Structures
• If the central atom is not yet
surrounded by four electron pairs,
convert one or more terminal atom
lone pairs to multiple bond pairs.
Only C-NOPS form multiple bonds!
Practice Problem
• Draw Lewis Structures for NH4+, CO,
NO+, and SO42-.
• Predict Lewis structures for
methanol, CH3OH, and hydroxylamine,
NH2OH.
Isoelectronic Species
• NO+, N2, CO, CN• Each has 2 atoms and the same
number of valence electrons (10),
which means they all have the same
Lewis structure
– Isoelectronic!
8.3 Atom Formal Charges
• Formal charge is the charge on an
atom/molecule/ion and the sum of
the charges equals overall charge on
an ion or is zero (for uncharged
molecules)
• Formal charge = group # - [LPE + ½ (BE)]
Practice Problem
•
Calculate the formal charge on
(a) CN(b) SO32-
8.4 Resonance
Structures
• Ozone, O3 has equal bond lengths, implying
that there is an equal number of bond
pairs on each side of the central O atom.
• Resonance structures… represent bonding
when a single Lewis structures fails to
accurately describe actual structure
• Single composite picture… is a resonance
hybrid; an actual structure
• Carbonate ion…
Resonance Structures
• Resonance structures differ only in
the assignment of electron pair
positions, NEVER atom positions.
• Resonance structures differ in the
number of bond pairs between a given
pair of atoms.
Practice Problems
• Draw resonance structures for the
nitrate ion, NO3-. Sketch a Lewis dot
structure for nitric acid, HNO3.
Homework
• After reading sections 8.1 – 8.4, you
should be able to do the following…
• P. 395 (2-16 even)
8.5 Exceptions to the
Octet Rule
• Fewer than eight – H at most only 2 electrons!
BeH2, only 4 valence electrons around Be! Boron
only has 6! (ammonia and boron trifluoride)
– When central atom is from Group 2A or 3A, the
electrons around central atom is twice the group #
• More than eight – only elements of the 3rd and
higher period (SF4, XeF2)
• Odd-electron compounds – a few stable
compounds contain an odd number of valence
electrons and cannot obey the octet rule (NO,
NO2, ClO2)  FREE RADICALS!
Coordinate Covalent
Bonds
• Some atoms such as N and P tend to
share a lone pair with another atom
that is short of electrons, leading to
a coordinate covalent bond.
– Ex. ammonia and boron trifluoride
8.6 Molecular Shape
• VSEPR – valence shell electron pair
repulsion theory
• Molecular shape changes with the numbers
of sigma bonds plus lone pairs about the
central atom
• Each lone pair or bond pair repels all other
lone pairs and bond pairs – they avoid each
other by making as wide an angle possible
Molecular Shape
σ bonds + lone
pairs on central
atom
2
Structure of
Molecule
Structural Pairs
Linear
180o
3
Trigonal planar
120o
4
Tetrahedral
109.5o
5
Trigonal
bipyramidal
120 & 90o
6
octahedral
90o
Geometry
• Electron-pair geometry: geometry
taken up by ALL the valence electron
pairs around a central atom
• Molecular geometry: arrangement in
space of the central atom and the
atoms attached to it.
Effect of Lone Pairs on
Bond Angles
• Strength of repulsion
• Lone pair > Bond pair
Practice Problem
• Give the electron pair geometry and
the molecular geometry for BF3 and
BF4-.
8.7 Bond Polarity and
Electronegativity
• Polar covalent bonds
– 2 dissimilar atoms = unequal sharing
– creates partial charges
• Electronegativity
– The ability of an atom in a molecule to
attract electrons to itself
– Table p. 376
Bond Polarity
• Bond polarity and electronegativity –
calculating the difference in EN can
determine how polar the bond is
– < 0.4 is NONPOLAR
– >1.67 is IONIC
– Between is POLAR and electrons are not
shared equally
Practice Problem
• Consider all possible resonance
structures for SO2. What are the
formal charges on each atom in each
structure? What are the bond
polarities? Do they agree?
8.8 Molecular Polarity
• Since most molecules have polar bonds,
molecules as a whole can be polar!
• Dipole moment is the product of the
partial charges (δ +/-) and the distance by
which they are separated.
– CO2 is not polar, while H2O is!
– A molecule will NOT be polar if all terminal
atoms (groups) are identical and all are
arranged symmetrically around central atom.
Practice Problem
• For each of the following, decide
whether the molecule is polar and
which side is positive and which
negative: BFCl2, NH2Cl, and SCl2.
Homework
• After reading sections 8.5-8.8, you
should be able to do the following…
• P. 396 (22-23, 27-29, 37-39, 42-43)
8.9 Bond Properties
• Bond order - # of bonding electron
pairs shared by 2 atoms in a molecule
–
–
–
–
–
1 – only a sigma bond
2 – 2 shared pairs
3 – 3 shared pairs
Fractional – resonance
Bond order = # shared pairs/# of links
Bond Properties
• Bond length – distance between the
nuclei of 2 bonded atoms
– Effect of bond order is evident when
comparing bonds between the same two
atoms
– Reduced by multiple bonds
Multiple Bonds
• Combinations of sigma and pi bonds
are stronger than pi alone. Pi bonds
are weaker than sigma but never
exist alone.
• Bond length is shorter for a double
than a single, and triple bonds are
the shortest of all!
Bond Dissociation
Enthalpy
• Bond energy – greater the # of
bonding electron pairs between a pair
of atoms, the shorter the bond.
Atoms are held together more tightly
when there are multiple bonds, so
there is a relation between bond
order and the energy required to
separate them.
Bond Dissociation
Enthalpy
• Bond dissociation energy (D) – energy
required to break a chemical bond
– D is + and breaking bonds is
endothermic!
• Energy supplied to break bonds must be the
same as the energy released when the same
bonds form.
– Bonds in reactants are broken and bonds
in products are formed
Bond Dissociation
Enthalpy
• ΔH = ΣΔH(bonds broken) – ΣΔH (bonds formed)
Practice Problem
• Using the bond energies in Table 8.9
(p. 389), estimate the heat of
combustion of gaseous methane, CH4.
9.1 Bonding and Molecular
Structure
• Valence Bond (VB) Theory:
– provides a qualitative picture of structure and
bonding; useful for molecules made of many
atoms; focuses on bonding/lone pairs being
localized
– good description of molecules in ground state
• Molecular Orbital (MO) Theory:
– describes molecules in higher energy states
– focuses on delocalized orbitals
9.2 Valence Bond Theory
• the closer two atoms become, the more
attraction there is between electrons of
one atom and nucleus of the other;
electron clouds are distorted
• more attraction means more overlap
– orbital overlap increases the probability of
finding bonding electrons in the region of space
between the two nuclei
Valence Bond Theory
• The covalent bond that arises from
the overlap of 2 s orbitals is called a
sigma (σ) bond.
• The electron density of a sigma bond
is greatest along the axis of the
bond.
Valence Bond Theory
• Orbitals overlap to form a bond between two
atoms.
• Two electrons, of opposite spin, can be
accommodated in the overlapping orbitals. Usually
one electron is supplied by each of the 2 bonded
atoms.
• Because of overlap, the bonding electrons have a
higher probability of being found within a region
of space influenced by both nuclei.
Hybridization of Atomic
Orbitals
• Orbital hybridization
– Hybrid orbitals could be created by mixing the
s, p, and sometimes d orbitals
– The number of hybrid orbitals is always the
same as the number of atomic orbitals that are
mixed to create the hybrid orbital set.
– The hybrid orbitals are more directed from
the central atom toward the terminal atoms.
Hybrid Orbitals
1. Draw Lewis structure.
2. Determine electron domain
geometry using VSEPR.
3. Specify hybrid orbitals needed to
accommodate electron pairs based
on geometric arrangement.
Hybrid Orbitals
2 pairs (sp)
Linear
180o
3 pairs (sp2)
Trigonalplanar
Tetrahedral
120o
4 pairs (sp3)
5 pairs (sp3d)
6 pairs (sp3d2)
Trigonalbipyramidal
Octahedral
109.5o
120o / 90o
90o
Practice Problems
• Identify the hybridization of the central
atom in the following compounds and ions:
• BH4• SF5• OSF4
• BCl3
• XeO64-
Multiple Bonds
• According to VB, double bonds
require two sets of overlapping
orbitals and two electron pairs.
• Triple bonds require three sets of
atomic orbitals.
Multiple Bonds
• A overlap of p atomic orbitals results
in a pi (π) bond. The pi bond requires
that the molecule be planar.
• Triple bonds consist of one sigma and
two pi bonds (due to two
unhybridized p orbitals).
Multiple Bonds
• Pi bonds do not occur without the bonded
atoms also being joined by a sigma bond.
• A pi bond may form only if unhybridized p
orbitals remain on the bonded atom.
• If a Lewis structure shows multiple bonds,
the atoms involved must therefore be
either sp2 or sp hybridized.
Cis-Tran Isomerism
• Rotation may occur around a single
bond.
• Restricted rotation around multiple
bonds result in isomers, compounds
with the same formula but different
structures.
Cis-Trans Isomerism
• Cis and trans-1,2-dichloroethylene
9.3 Molecular Orbital
Theory
• MO assumes that pure atomic
orbitals of the atoms combine to
produce orbitals that are spread out,
or delocalized.
• The new orbitals are called molecular
orbitals.
Molecular Orbital Theory
• The total number of molecular orbitals is
always equal to the total number of atomic
orbitals contributed by the atoms that
have combined.
• The bonding molecular orbital is lower in
energy that the parent orbitals.
• Electrons of the molecule are assigned to
orbitals of successively higher energy.
Homework
• After reading Chapter 9, you should
be able to do the following:
• P.434 (3-8,21-24)