Basic Concepts of Chemical
Bonding
Ionic, Covalent, Metallic
Bonding
Polarity &
Electronegativity
Lewis Structures
Three Types of Chemical
Bonds
 Ionic Bond: electrostatic forces that exist
between ions of opposite charge. Metals
+ nonmetals. Ex: NaCl
 Covalent Bond: sharing of electrons
between two atoms. Nonmetals. Ex: CO2
 Metallic Bond: metal ions in a “sea of
valence electrons.” Ex: Cu
Lewis Symbols
 Only valence electrons are involved in
chemical bonding.
 We use Lewis electron-dot symbols (Lewis
symbols)to show valence electrons.
 Lewis symbol = element symbol + dots
representing electrons. Use all four sides
around element symbol and no more than
two electrons per side. Usually balance one
dot on one side with dot on opposite side.
Examples of Lewis Symbols
The Octet Rule
 Octet Rule: Atoms tend to lose, gain, or
share electrons until they are surrounded
by eight valence electrons.
 Full s & p orbitals = stability
Sodium Chloride Formation
 Na(s) + 1/2 Cl2(g)  NaCl(s)
 There is a transfer of an electron from
the Na atom to the Cl atom.
Ionic Bonding:
Formation of NaCl
NaCl is
composed of
Na+ and Cl- ions
that are
arranged in a
regular threedimensional
array.
Stability of Ionic Compouds
 Very neg. heats of formation of
ionic compounds.
 Ions are drawn together,
releasing energy as a solid
lattice forms.
 Ionic compounds are very
stable.
 Ionic compounds are hard &
brittle with high melting points.
Lattice Energy
 Lattice energy: the energy required to
completely separate a mole of solid
ionic compound into its gaseous ions.
 Note: This is an endothermic process.
 ***The higher the lattice energy, the
stronger the ionic bond.***

NaCl(s)  Na+(g) + Cl-(g)
∆H = +788 kJ/mol
Lattice Energy, Cont.
 Lattice energy depends on:

Eelectric =  (Q1Q2)/d
 Where: Q1 and Q2 = ion charges;
 = constant; d = distance between
center of ions
 Lattice energy increases as the
charges on the ions increase and as
their radii decrease.
1+; 1-
2+; 1-
2+; 2-
3+; 3-
Magnitude of lattice energy
depends mainly on ionic charge
because radii do not vary
greatly.
Practice With Lattice Energy
 Arrange the following in order of
increasing lattice energy: NaF, CsI, CaO
 CsI < NaF < CaO
 Which substance has the greatest lattice
energy? Why? AgCl, CuO, CrN
 CrN, because of larger 3+/3- charges
Born-Haber Cycle
 Lattice energies cannot be
measured experimentally.
 Use Hess’s Law to construct a
Born-Haber cycle.
 A Born-Haber cycle shows the
formation of ionic solids from
elements.
The principles of Hess’s
Law can be applied in
the Born-Haber cycle to
determine the lattice
energy of NaCl.
Covalent Bonding
 Covalent Bond: chemical bond formed
by sharing of electrons to form noble
gas configurations. Ex: H2 molecule
Multiple Bonds
 Single bond: sharing of 2 electrons.
Shown as H-H or H:H
 Double bond : sharing of 4 electrons.
Shown as O::C::O or O=C=O
 Triple bond: sharing of 6 electrons.
Shown as :N:::N: or :NN:
Examples of Multiple Bonds
Bond Lengths
 Average bond distance varies with number of
bonds. Generally, as number of shared electron
pairs increases, bond length decreases.
 N-N (1.47Å) N=N (1.24Å) NN (1.10Å)
 Note: A triple bond is NOT 3x stronger than a
single bond. The first bond in a triple bond is
stronger than a “normal” single bond. The
“second” and “third” bonds in a triple bond are
weaker than the “first” bond.
Bond Polarity
 Bond polarity describes the sharing of electrons
between atoms.
 Nonpolar covalent bond: electrons shared
equally between atoms.
 Polar covalent bond: one of the atoms in a
compound exerts a greater attraction for the
bonding electrons than the other. If the
difference is big enough, we get an ionic bond.
Electronegativity
 Use electronegativity to estimate whether a
bond will be nonpolar covalent, polar
covalent, or ionic.
 Electronegativity: the ability of atoms in a
molecule to attract electrons to itself.
 The greater the electronegativity, the
greater its ability to attract electrons.
 High ionization energy & very negative
electron affinity = high electronegativity.
Electronegativity Scale
 Linus Pauling makes electronegativity
scale.
 Values are unitless
 Scale: 0.7 (cesium) - 4.0 (fluorine)
 Metals are less electronegative;
nonmetals are more electronegative.
Electronegativity &
Bond Polarity
 The greater the difference in electronegativity
between two bonded atoms, the more polar
the bond.
 Difference < 0.5 = nonpolar covalent
Ex: F2 4.0 - 4.0 = 0
 0.5 ≤ difference < 2.0 = polar covalent
Ex: HF
4.0 - 2.1 = 1.9
 Difference ≥ 2.00 = ionic bond
Ex: LiF 4.0 - 1.0 = 3.0
 Use + and - to show partial positive and
negative charges on atoms.
Nonpolar covalent bond
+
-
Polar covalent bond
+
-
Ionic bond
Polarity Practice
 Tell which bond is more polar and
indicate in each case which atom has
the partial negative charge.
 B-Cl or C-Cl?
 B-Cl, with - on Cl
 P-F or P-Cl?
 P-F, with - on F
Polar Molecules
 Entire molecules can be polar, not just
bonds within the molecules. Ex: HF
 Polarity is important--it helps to
determine many properties of a cmpd!
Drawing Lewis Structures
 1. Sum valence e-’s of all atoms. Add e-’s
for anions.
 2. Sum number of e-’s each atom needs
for an octet. All atoms need 8 e-’s except
for H that needs only 2 e-’s.
 3. (e-’s needed-valence e-’s)/2 = minimum
number of bonds in the molecule
 Note: bonds may be single or multiple.
One double bond=two single bonds
Lewis Structures, Cont.
 4. Draw symbols, attaching bonds as
needed. Note: H only takes a single
bond. C takes 4 bonds, is often central.
 5. Complete the octets of atoms bonded
to central atom. Use dots for lone pairs
of e-’s.
 6. Place any leftover e-’s on central
atom even if it has more than an octet.
Draw Structure for PBr3







Which is your central atom?
Determine number of valence e-’s.
P has 5, Br has (3 x 7).
5 + 21 = 26 e-’s
Determine number of e-’s needed for octets.
P has (1 x 8), Br has (3 x 8). 4 x 8 = 32 e-’s
(Needs - valence)/2 = minimum bonds needed
(32-26)/2 = 3 bonds needed at minimum
PBr3 Lewis Structure
 Note: bonds + lone pairs = number of
valence e-’s.
Exceptions to Octet Rule
 Odd number of electrons:
Less than an octet:
Exceptions, Cont.
 More than an octet:
Formal Charge
 Sometimes multiple Lewis structures
can be drawn for a molecule. Formal
charge can help determine correct
one.
 Formal charge: charge that an atom
would have in a molecule if all atoms
had the same electronegativity
(bonding e-’s are shared equally).
Calculating Formal Charge
 Electrons are assigned as follows:
 1) All unshared (nonbonding) e-’s are
assigned to the atom on which they are
found.
 2) Half of the bonding e-’s are assigned to
each atom in the bond.
 Formal charge = # of valence e-’s - # of
e-’s assigned in Lewis structure.
Oxidation
Number
Formal
Charge
Polarity
 It is important
to note that
oxidation
number
overstates the
importance of
electronegativity
and formal
charge ignores
the role of
electronegativity
completely.
Formal Charge as the
“Tie Breaker”
 As a general rule, when several
Lewis structures are possible, the
most stable (and favored) one will be
the one in which…
 1) the atoms bear formal charges
closest to zero.
 2) any negative charges reside on
the more electronegative atoms.
Formal Charge & NO2
Resonance Structures
 Sometimes molecules and ions cannot be
described by a single Lewis structure.
 Resonance structure: atoms keep same
arrangement but placement of e-’s
changes. Look for changes in placement
of double bonds.


Resonance
 All of the resonance
structures are taken
together as a description
of a molecule or ion.
A molecule or ion does not switch among
resonance structures, it is more of a blending.
Not all resonance structures are equivalent.
Formal charges can help to predict favored
structures.
Resonance Structures
Resonance In Benzene
Benzene
(C6H6) is an
important
example of
an aromatic
organic
molecule.
It is the
“poster
child” for
resonance.
Bond Enthalpy
 Stability of a molecule = strengths of its
covalent bonds.
 Bond Enthalpy: ∆H for the breaking of a
particular bond in a mole of gaseous
substance.
 Always endothermic.
 Greater bond enthalpy = stronger bond
 Stronger bonds = less reactivity