Chemical Bonds, Molecular
Geometry, and Bond Theory
Brown and LeMay
Chapters 8 and 9
Introduction to
types of bonding
• Three basic types of bonds
– Ionic
• Electrostatic attraction
between ions.
– Covalent
• Sharing of electrons.
– Metallic
• Metal atoms bonded to
each other – involves
delocalized electrons.
8.1 – Lewis Symbols and the Octet Rule
• Octet rule – when forming bonds, atoms try to achieve the
s2p6 configuration of a noble gas
• Lewis symbols – dots represent the valence electrons (no
pairs until more than four)
8.2 – Ionic Bonding
 Ionic bonds = electrostatic forces that hold cations and anions
together
 Ionic crystal = a highly ordered solid collection of ions
 Using Lewis Symbols to show ionic bonding:
• Properties of ionic compounds:
1.
2.
3.
4.
5.
6.
Brittle
High melting points
Crystalline
Can be cleaved (break apart along smooth surfaces)
The strength of the ionic bond can be measured with lattice energy (the
energy required to completely separate a mole of a solid ionic compound
into its gaseous ions). Lattice energies are positive values, so the reverse
process (forming the ionic bonds) is exothermic.
Lattice energy increases with:
– increasing charge on the ions
– decreasing size of ions
– The bigger effect on lattice energy is from the ionic charge
8.3 – Covalent Bonding
• Basic Lewis Structures
– Lewis Structure = using Lewis symbols to represent covalent bonds
between atoms.
– Bonding pairs = the shared pairs of electrons. Shown as dots between
the atoms sharing the electrons, or more commonly as a line for each
shared pair
– Lone pairs (or nonbonding pairs) = the unshared pairs of electrons.
Shown as dots.
– Coordinate covalent bond = when one of the
atoms provides both electrons of the shared pair.
Ex. H3O+, NH4+
– Multiple covalent bonds: ex. CO2, N2
8.4 – Bond Polarity and Electronegativity
• The electrons in a covalent bond are not always
shared equally.
• Ex. Fluorine pulls harder on the electrons it
shares with hydrogen than hydrogen does.
• Therefore, the fluorine end of the molecule has
more electron density than the hydrogen end.
Electronegativity
• Electronegativity is the ability of an atom in a
molecule to attract electrons to itself.
• On the periodic table, electronegativity generally
increases as you go
– from left to right across a period.
– from the bottom to the top of a group.
• Nonpolar covalent bond – when the electrons are shared
equally (diff. in electroneg. is  ~0.4)
• Polar covalent bond – when the electrons are not shared
equally (diff. in electroneg. is between ~0.4 and ~2.0)
• We depict this as follows:
•
•
+
H  Cl
or
H  Cl
Is a Compound Ionic or Covalent?
• Simplest approach: Metal + nonmetal is ionic;
nonmetal + nonmetal is covalent.
• There are many exceptions: It doesn’t take into
account oxidation number of a metal (higher
oxidation numbers can give covalent bonding).
• The electronegativity table also doesn’t take into
account oxidation number.
• Properties of compounds are often best: Lower
melting points mean covalent bonding, for
example.
8.5 – Drawing Lewis Structures
•



Hints for drawing plausible Lewis structures:
Draw a skeletal structure first
H is always a terminal atom
The central atom of a structure usually has the
lowest electronegativity
 In oxoacids, H is usually bonded to O, not the
central atom
 Molecules and polyatomic ion structures are
usually compact and symmetrical
• A method for drawing Lewis structures:
1) Determine the total # of valence electrons in the
molecule or ion
2) Draw a skeletal structure using the hints on the
previous slide and connect atoms with a single
bond
3) Complete the octets around all the atoms bonded
to the central atom
4) Use any remaining electrons for lone pairs around
the central atom(s)
5) If there’s not enough electrons to make an octet
around the center atom, form multiple bonds.
• Formal Charge = the difference between the # of
valence electrons in an isolated atom and the # of
electrons assigned to that atom in a Lewis structure
(it’s a hypothetical charge).
• # of electrons assigned = # of lone pair electrons + ½
the # of electrons in bonds.
• Example:
• Try to draw Lewis structures so that the formal charge is
0, or with a minimum formal charge
• Formal charges of adjacent atoms should be of opposite
sign
• The atoms with the greatest electroneg. should have neg.
formal charges, if any
• The total of the formal charges must equal the charge of
the molecule
• Example:
• The middle one would be the “dominant” Lewis structure
8.6 – Resonance Structures
• Resonance (9.8): When we can draw different Lewis
structures for the same molecule that differ only in the
distribution of electrons (resonance structures), the actual
molecule is a hybrid of the different structures (resonance
hybrid). These hybrids have delocalized electrons (electrons
shared by multiple atoms).
• Ex.
• The organic compound
benzene, C6H6, has two
resonance structures.
• It is commonly depicted
as a hexagon with a
circle inside to signify
the delocalized
electrons in the ring.
8.7 – Exceptions to the Octet Rule
• 3 categories:
1) Odd number of valence electrons – most of these
are free radicals (very reactive molecular
fragments). A bold dot () is sometimes used to
represent the unpaired electron.
• 2) Incomplete Octets – usually the central atom is
Be, B, or Al
• The general rule is: If filling the octet of the central
atom results in a negative formal charge on the
central atom and a positive formal charge on the
more electronegative outer atom, don’t fill the octet
of the central atom.
3) Expanded Octets – happens with elements in Period
3 and higher, because a d sublevel is needed in
order to have more than 8 valence electrons.
Ex. PF5
ex. Phosphate ion – dominant
structure is the one with less
formal charges and an 5 bonds
on the phosphorous
8.8 – Strengths and Lengths of Covalent Bonds
• Strength of a Covalent Bond
• The strength of a bond is measured by determining how much
energy is required to break the bond.
• This is called the bond enthalpy.
• For ex. the bond enthalpy for a Cl—Cl bond, D(Cl— Cl), is
measured to be 242 kJ/mol.
• Bond enthalpies are positive, because bond breaking is an
endothermic process. Bond forming is exothermic, so the
negative value of the bond enthalpy is used.
• Average bond enthalpies are used because the actual bond
enthalpy is influenced by the other bonds in the molecule
• The greater the bond enthalpy, the stronger the bond. A
molecule with strong bonds generally is less reactive than one
with weak bonds.
• One way to estimate H for a reaction is to use the bond
enthalpies of bonds broken and the new bonds formed:
• Hrxn = (bond enthalpies of all bonds broken) − (bond
enthalpies of all bonds formed).
• Example: CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g)
• In this example, one C—H bond and one Cl—Cl bond
are broken; one C—Cl and one H—Cl bond are formed.
H = [D(C—H) + D(Cl—Cl)] − [D(C—Cl) + D(H—Cl)]
= [(413 kJ) + (242 kJ)] − [(328 kJ) + (431 kJ)]
= (655 kJ) − (759 kJ)
= −104 kJ
• Bond Length
• We can also measure an average bond length
(distance between the 2 nuclei) for different bond
types.
• There’s a relationship between bond enthalpy, bond
length, and number of bonds between 2 atoms.
• In general, as the number of bonds between 2 atoms
increases, the bond grows shorter and stronger.
9.1, 9.2 – Molecular Shapes and VSEPR Theory
• Electron group geometry (or electron domain) = the
arrangement of all the electron groups around the central
atom.
• Electron group (or domain) = a collection of valence electrons
in a certain region around the central atom. An electron
group could be a single e-, a lone pair, a single bond, a double
bond, or a triple bond.
• 2 groups = linear electron group geometry (has bond angles of
180°)
• 3 groups = trigonal planar (120°)
• 4 groups = tetrahedral (109.5°)
• 5 groups = trigonal bipyramidal (contains both 90° and 120°)
• 6 groups = octahedral (90°)
• see pg 345
• VSEPR Theory (valence-shell-electron-pair- repulsion)
• Simply put, electron pairs, whether they be bonding or
nonbonding, repel each other.
• By assuming the electron pairs are placed as far as possible
from each other, we can predict the shape of the molecule.
• VSEPR notation: central atom = A, terminal atoms = B,
nonbonding electrons = E
• Ex. AB3E is a molecule with 3 terminal atoms and a lone pair.
Ex. NH3
• Molecular geometry = the shape formed from the bonded
atoms (nonbonding electrons are not included)
• Once you have determined the electron-domain
geometry, use the arrangement of the bonded atoms to
determine the molecular geometry.
• Tables 9.2 and 9.3 show the potential molecular
geometries. We will look at each electron domain
to see what molecular geometries are possible.
Linear Electron Domain
• In the linear domain, there is only one
molecular geometry: linear.
• NOTE: If there are only two atoms in the
molecule, the molecule will be linear no
matter what the electron domain is.
Trigonal Planar Electron Domain
• There are two molecular geometries:
1) trigonal planar, if all electron domains are
bonding
2) bent, if one of the domains is a
nonbonding pair.
Tetrahedral Electron Domain
• There are three molecular geometries:
1. tetrahedral, if all are bonding pairs,
2. trigonal pyramidal, if one is a nonbonding
pair, and
3. bent, if there are two nonbonding pairs.
Trigonal Bipyramidal Electron Domain
• There are two distinct
positions in this
geometry:
– Axial
– Equatorial
• Lone pairs occupy
equatorial positions.
Trigonal Bipyramidal
Electron Domain
• There are four distinct
molecular geometries in
this domain:
• 1. Trigonal bipyramidal (all
bonding pairs)
• 2. Seesaw (1 nonbonding
pair)
• 3. T-shaped (2 nonbonding
pairs)
4. Linear (3 nonbonding
pairs)
Octahedral Electron Domain
• All positions are equivalent
in the octahedral domain.
• There are three molecular
geometries:
• 1. Octahedral (0
nonbonding pairs
• 2. Square pyramidal (1
nonbonding pairs)
• 3. Square planar (2
nonbonding pairs)
Effect of Nonbonding Electrons and Multiple Bonds
on the Ideal Bond Angles
• Nonbonding pairs are physically larger than bonding pairs.
• Therefore, their repulsions are greater; this tends to compress bond
angles.
• Double and triple bonds have larger electron domains than single
bonds.
• They exert a greater repulsive force than single bonds, making their
bond angles greater.
Shapes of Larger Molecules
For larger molecules,
look at the geometry
about each atom
rather than the
molecule as a whole.
9.3 – Molecular Shape and Polarity
• For a molecule to be polar, it has to meet 2 criteria:
• 1. it has to have polar bonds, and
• 2. it has to have a shape in which the bond dipoles don’t
cancel each other out (it has a dipole moment)
9.4 - 9.6: Valence Bond Theory (Hybrid Orbitals)
• Valence Bond Theory looks
at a covalent bond in terms
of the overlap of atomic
orbitals of the bonded
atoms. The overlap of
orbitals allows the bonding
electrons to share the space
between the nuclei, and
they are simultaneously
attracted to both nuclei,
forming a covalent bond.
• The picture below shows how the potential energy of 2
hydrogen atoms changes as they come together to form
H2. As the electrons and nuclei come closer together, a
balance is reached between the like charge repulsions
and the electron-nucleus attraction.
• The internuclear distance at the lowest point of the PE
curve is the bond length.
• Hybridization: the mathematical mixing of atomic
orbitals. This happens with the central atom of
molecules. Terminal atoms do not hybridize.
• Hybrid orbitals form by “mixing” of atomic orbitals to
create new orbitals of equal energy, called degenerate
orbitals.
• The shape of a hybrid orbital is different than the shapes
of the original atomic orbitals. The shapes of the hybrid
orbitals mesh with the molecular geometry that VSEPR
determines.
• When 2 atomic orbitals “mix” they create 2 hybrid
orbitals; when 3 atomic orbitals mix, they create 3 hybrid
orbitals; etc.
sp hybrids
• sp hybridization is the “mixing” of one s atomic orbital and
one p atomic orbital. It leads to linear e- pair geometry and
180o bond angles.
• This is consistent with the observed geometry of
molecules like BeF2
sp2 hybridization
• sp2 hybridization is the “mixing” of one s atomic orbital and
two p atomic orbitals. It leads to trigonal planar e- pair
geometry and 120o bond angles.
sp3 hybridization
sp3 hybridization is the “mixing” of one s atomic orbital and three
p atomic orbitals. It leads to tetrahedral e- pair geometry and
109.5o bond angles.
Multiple Bonds
• Sigma () bond = an end-to-end overlap of simple or hybrid
orbitals along a line between the 2 nuclei (i.e. all the bonds
we’ve talked about so far)
• Pi () bond = a side-to-side overlap of p orbitals, producing
high e- charge density above and below a line between the 2
nuclei. Pi bonds are weaker than sigma bonds.
 Single bonds are  bonds. Double bonds consist of
1  and 1  bond. Triple bonds consist of 1  and
2  bonds.
 A  bond involves more overlap (i.e. is stronger)
than a  bond.
 The shape of the molecule is determined only by
the  bonds.
 Rotation about the double bond is very restricted.
• Geometric Isomers = organic molecules that differ in
the positions of attachment of substituent groups at
a double bond.
• Cis isomers have the substituent groups on the same side of
the molecule
• Trans isomers have the substituent groups of opposite sides.
Localized or Delocalized Electrons
• Bonding electrons (σ or π) that are specifically shared
between two atoms are called localized electrons.
• In molecules that have resonance structures, we have
electrons shared by multiple atoms. They are called
delocalized electrons. Example: benzene
9.7, 9.8: Molecular Orbital (MO) Theory
• Molecular Orbital Theory views a molecule as a whole
instead of a collection of individual atoms.
• Uses the wave functions of a molecule
• Molecular orbitals have many characteristics like atomic
orbitals:
– maximum of two electrons per orbital
– electrons in the same orbital have opposite spin.
– have specific energy levels
• Whenever two atomic orbitals overlap, two molecular orbitals are
formed: one bonding, one antibonding.
• Bonding orbitals are constructive combinations of atomic orbitals.
• Antibonding orbitals are destructive combinations of atomic orbitals.
They have a new feature unseen before: A nodal plane occurs where
electron density equals zero.
Whenever there is direct overlap
of orbitals (the electron density is
centered about the internuclear
axis), forming a bonding and an
antibonding orbital, they are called
sigma (σ) molecular orbitals.
The antibonding orbital is
distinguished with an asterisk as
σ*. Here is an example for the
formation of a hydrogen molecule
from two atoms.
MO Diagram: has the interacting atomic orbitals on the left and right,
and the MOs in the middle. It shows how orbitals combine to form the
molecule.
• In H2 the two electrons go into the
bonding molecular orbital (lower in
energy).
• Bond order = ½(# of bonding
electrons – # of antibonding
electrons) = ½(2 – 0) = 1 bond
• A bond order of 1 is a single bond.
• A bond order of 2 is a double bond.
• A bond order of 3 is a triple bond.
• 1/2, 3/2, or 5/2 are possible
(molecules containing an odd # of
electrons)
Can He2 Form? Use MO Diagram and Bond
Order to Decide!
• Bond Order = ½(2 – 2) = 0
bonds
• A bond order of 0 means
that the bond doesn’t exist.
• Therefore, He2 does
not exist.
σ and π bonds
• Molecular Orbitals from 2p atomic
orbitals:
• The p orbitals that face each
other overlap in  fashion, like the
s orbitals.
• The other two sets of p orbitals
overlap in  fashion.
The MO diagram for the 2nd energy level
 There are σ and σ* orbitals from
the 2s and 2p atomic orbitals.
 There are also π and π* orbitals
from 2p atomic orbitals.
 Since direct overlap is stronger,
the effect of raising and
lowering energy is greater for σ
and σ*
 For O2, F2, and Ne2, the order
of energy for the 2p sublevel is
σ2p < π2p < π*2p < σ*2p
 For B2, C2, and N2, the order of
energy is π 2p < σ 2p < π*2p <
σ*2p because of interaction
between the 2s and 2p.
MO Diagrams for Diatomic Molecules of 2nd
Period Elements
MO Diagrams and Magnetism
• Diamagnetism is the result of all electrons in every orbital
being spin paired. These substances are weakly repelled
by a magnetic field.
• Paramagnetism is the result of the presence of one or
more unpaired electrons in an orbital. These substances
are attracted to a magnetic field.
• Is oxygen (O2) paramagnetic or diamagnetic? Look back at
the MO diagram! It is paramagnetic.
• Lewis structures would not predict that O2 is paramagnetic,
but experimental evidence shows that it is.
Heteronuclear Diatomic Molecules
• Diatomic molecules can consist of
atoms from different elements.
• How does a MO diagram reflect
differences?
• The atomic orbitals have different
energy, so the interactions change
slightly.
• The more electronegative atom has
orbitals lower in energy, so the
bonding orbitals will more resemble
them in energy.