
Types of Bonds
› Ionic
› Covalent
› Metallic

Lewis Symbols

Octet Rule
Energy of bond formation is usually
negative
 Form lattice or 3-D array of + and – ions

› Lattice energy – energy required to separate
a mole of ionic solid into gaseous ions
 Based on charges, sizes, arrangement

Electron Configurations
› Representative Elements
› Transition Elements

Determines:
›
›
›
›

Lattice energy
Way ions are arranged
Properties
Biological activity
Determined by:
› Nuclear charge
› Number of electrons
› Orbital in which the outer-shell electrons are
found

Share electrons
› Sharing pattern

Bond polarity and electronegativity
› Equal vs. unequal sharing
 Nonpolar covalent
 Polar covalent
› Electronegativity used to predict bond polarity
 Difference in electronegativity is bond polarity
 <0.5 Nonpolar covalent
 0.5-2.0 Polar covalent
 >2.0 Ionic bond

Polar Molecules
› 2 atom molecules
› >2 atom molecules
Sum valence electrons from all atoms
Write symbols and connect with single line
Complete octets of atoms bonded to
central atom
4. Place any leftover electrons around the
central atom
1.
2.
3.
›
5.
If there are not enough electrons around the
central atom, try multiple bonds
Be and B do not satisfy octet. Elements
from the 3rd period and higher may have
expanded octets.

Occur when a single Lewis structure does
not represent the molecule.

Possible whenever a Lewis structure has
a multiple bond and an adjacent atom
with at least one unshared pair of
electrons

Odd numbers of electrons
› ClO2, NO, NO2

Less than an octet
› B

More than an octet
› Pd. 3 and higher – uses d orbitals
Bond Enthalpy - Always positive
 Triple > Double > Single
 Bond enthalpies can be used to
estimate enthalpies of reactions

› ΔHrxn = Σ(bonds broken) – Σ(bonds formed)

Higher bond enthalpies means shorter
bond length

VSEPR model
› Includes bonding and non-bonding pairs of
electrons

Shapes
Shapes of Molecules
•Linear
•Bent
•Trigonal planar
•Trigonal pyramid
•T-shape
•Tetrahedral
•Square planar
•See-saw
•Trigonal bipyramidal
•Square pyramidal
•Octahedral
•Examples
Distinguish between bonding and nonbonding pairs
 Define electron pair geometry

› 5 electron pair geometries: linear, trigonal
planar, tetrahedral, trigonal bipyramid,
octahedral

Determine 3-D structure of molecule
› Lone pairs are ignored for molecular geometry

Bond angle decreases as the number of
lone pairs increases
› Tetrahedral: 109.5, 107, 104.5

VSEPR is extended and each central
atom is determined independently
›
›
›
›
Methanol, CH3OH
Acetic acid CH3COOH
Urea, (NH2)2CO
Nitrogen peroxide, N2O4
Determine Lewis Structure, Resonance structures (if
applicable), Geometry, Polarity
Bond polarity – differences in
electronegativity
 Molecule polarity – bond polarity in
combination with molecular geometry

› Bond dipoles are equal in magnitude and
opposite in direction – NONPOLAR molecule
› Non-cancelling dipoles – POLAR molecule

Covalent bonds form when orbitals on
two atoms overlap
› As overlap increases, energy of interaction
decreases until min. energy is reached
(bond length)
› At bonding distance, attractive forces
balance repulsive forces

Hybrid orbitals explain molecular
geometry

Sigma Bonds are the first bond that forms
between two atoms.
› Sigma bonds are the only bonds that are used to
determine the electron domain and molecular
shape.
› Sigma bonds form in the space directly between two
atomic nuclei.

Pi Bonds are the second or third bond that
forms between the same two atoms.
› Pi bonds are NOT used to determine molecular
shape. However, they do shorten the length of the
bond – pulling the nuclei closer together.
› Pi bonds form around the sigma bond – either top
and bottom or front and back.

Single bond
› 1 sigma bond
› Ethane, C2H6

Double bond
› 1 sigma bond
› 1 pi bond
› Ethene, C2H4

Triple bond
› 1 sigma bond
› 2 pi bonds
› Ethyne, C2H2
Every pair of bonded atoms shares one
or more pairs of electrons
 Two electrons shared on the same axis as
the nuclei are σ bonds

› σ bonds are always localized in the region
between the two bonded atoms
If two atoms share more than one pair of
electrons, the additional pairs form π
bonds
 If resonance is possible, delocalized
electrons are also possible.


Use Molecular Orbital theory to explain
molecules that aren’t explained by
VSEPR or Hybrid theory

Electrons in molecules are found in
molecular orbitals.
› Each contains a max. of 2 electrons
› Each has a definite energy
› Associated with ENTIRE molecule (not single
atom)

When two Atomic Orbitals overlap, two
Molecuar Orbitals form.
› One with high electron density (bonding orbital)
› One with low density (antibonding orbital)

Bonding orbitals have lower energy than
antibonding orbitals
As applies to homonuclear diatomic molecules. If
heteronuclear diatomic molecules have a small
electronegativity difference, MO will be similar

Bond Order = ½ (bonding – antibonding
electrons)
› Single bond – Bond order = 1
› Double bond – Bond order = 2
› Triple bond – Bond order = 3
 Bond order = 0 is unstable (doesn’t exist)

As bond order increases,
› Bond length decreases
› Bond energy increases

Two Types
› Paramagnetism (unpaired electrons) – strong
attraction between magnetic field and
molecule
› Diamagnetism (no unpaired electrons) – weak
repulsion between magnetic field and molecule

Behavior is determined by measuring the
mass of a sample in the presence/absence
of magnetic field
› Large increase indicates paramagnetism
› Small decrease indicates diamagnetism

Experimentally, O2 is paramagnetic.
› Lewis structure
› Molecular Orbital diagram

Experimentally, O2 has a short bond
length and high bond dissociation
energy
› Suggests double bond

Molecular Orbital theory predicts both.