This document summarizes the key differences between ionic and covalent bonding. Ionic bonds form when electrons are transferred from a metal to a nonmetal, creating oppositely charged ions that are attracted via electrostatic forces. Covalent bonds form when electrons are shared between two or more nonmetal atoms, resulting in overlapping electron clouds between the bonded atoms. Most covalent substances exist as molecules where covalent bonds exist within molecules but intermolecular forces exist between molecules. Electronegativity is used to predict bond polarity, with larger electronegativity differences indicating more ionic character. Bond length and energy are also related, with shorter, stronger bonds having higher energies.

1. Covalent Bonding
Chapter 6-1 (Old Text)
Chapter 9-1 and 9-5 (New Text)

2. Ionic Bond
Ideally, the electron clouds of each
ion do not overlap. Electron transfer
occurs from metal to nonmetal and
the resulting ions attract only by
electrostatic attraction.
Na1+ Cl1-

3. Ionic
Compounds
have a rigid
and strong
crystal lattice
structure in
which the
ions are
arranged in
an orderly,
repeating
pattern.
Cl - ion
Na+ ion

4. Covalent Bond
Electrons are shared between two or more nonmetal
atoms in a molecule. There are no ions! Electron
clouds of adjacent atoms overlap and merge to form
new bonding orbitals (a process called hybridization).

5. Most Covalent Substances Have a
Molecular Form
There are covalent bonds inside molecules but not between them!
The attractions between molecules are called intermolecular forces
(as opposed to intramolecular forces which are true covalent bonds).
The strengths of the intermolecular forces determine the melting and
boiling points of molecular substances. You don’t break any chemical
bonds to melt molecular substances!

6. Electronegativity – A measure of the ability of an atom to attract shared
electrons in a chemical bond , on a relative scale of 0 to 4. (American Chemist
Linus Pauling developed this concept.) The noble gases are not ranked because
they don’t form bonds. Notice that F is the best at attracting electrons in a bond.
Finding the Difference in Electronegativity (ΔEN) between two bonded atoms
can be used to predict whether a specific bond is more ionic or more covalent. In
reality, most bonds are a little of both. ΔEN is calculated for each bond in a
molecule. Just subtract the electronegativity values of the two bonded atoms.

7. EN > 2.1
Ionic Bond – Electron Transfer Between 2 Atoms
EN = 0.4 to 2.1
Polar Covalent Bond – Unequal Sharing of the
Shared Electrons Between 2 Atoms
EN < 0.4
Pure Covalent Bond (Nonpolar Covalent) –
Essentially equal Sharing of the Shared Electrons
Between 2 Atoms

8. Examples
Predict the Bond Type based on EN:
(EN values are always positive)
The first one is done for you…
1) Na and Cl EN = 0.9 – 3.0 = 2.1, polar covalent bond
2) Ca and O
3) C and H
4) Al and Cl
5) Mg and H
6) S and O

9. Polar Covalent vs. Pure Covalent
Bonds
The O-H bond (EN = 1.4) in
water is a polar covalent bond.
The shared electrons spend more
time with oxygen than with
hydrogen, thus the electron
sharing is uneven.
The Br-I bond (EN = 0.3) in
bromine iodide is a pure
covalent (nonpolar) bond. The
shared electrons essentially spend
equal time with bromine and
iodine.
Br I
H
O
H
2.1 2.1
3.5
2.5
2.8

10. Covalent Bonds are Like Springs!
The atoms can vibrate back and forth
to some extent without the bond breaking.
The strength of a bond is related to its length,
which is in turn related to the radii of the
bonding atoms.

11. Bond Length
• Bond length is the average distance between the nuclei of two
bonded atoms. It is the position of highest stability that balances
the forces of attraction and repulsion between two bonded
nuclei. Energy is released when bonds form from atoms. The
chemical potential energy of the chemical system is minimized
(lowest) at the ideal bond length distance.
Chemial
Potential

12. Comparing Bond Length vs. Bond Energy
Bond (Dissociation)Energy is the amount of energy needed to
break a bond. It is the same amount of energy as was released
when the bond formed. As bond length increases, bond energy
decreases. In other words, longer bonds are weaker and easier
to break than shorter bonds! (note: pm = picometer or 10-12 m)
Bond Length (pm) Bond Energy
(kJ/mol)
H-H 75 436
H-C 109 418
H-Cl 127 432
H-Br 142 366
H-I 161 298

13. Unstable Molecule – NI3
2NI3(s) ——> N2(g) + 3I2(g)
Iodine is a much bigger atom than Nitrogen! The N-I
bond is long and weak! Tickling from a feather is enough
to break the bond, producing diatomic nitrogen gas and
solid diatomic iodine (the purple smoke)! So, NI3 is very
unstable, but N2 is very stable (triple bond)!
N
I
I
I

14. Table of Average Bond Dissociation
Energies in kJ/mol
(Bond Enthalpies)
These are average
values because:
1) Atoms in a
bond
continually
vibrate (move),
and
2) The same type
of bond may
be measured
inside different
molecules and
an average
strength
calculated

15. To calculate the energy needed to
break all the bonds in a molecule….
• 1st - Draw the Lewis Structure of the molecule in order to see what specific
bonds are in the molecule.
• 2nd - Use a table of average bond energies to look up each type of bond.
Pay specific attention to whether the bonds are single, double, or triple
bonds.
• 3rd - Add together the bond energies for each bond in the molecule.
• Ex: Calculate the energy needed to break the bonds in Hydrocyannic Acid
(HCN):
Lewis Structure
H – C  N :
Total energy to break bonds in HCN is:
1 H-C bond (413 kJ/mol) + 1 CN bond ( 891 kJ/mol ) = 1304 kJ/mol
(see the bond energies from the previous slide)
Make sure you account for each and every bond in the molecule!

16. Single vs. Double vs. Triple Bonds
Bond Length vs. Bond Energy
Bond Bond Length
(pm)
Bond Energy
(kJ/mol)
C - C 154 348
C = C 134 614
C  C 120 839
N - N 145 170
N = N 124 418
N  N 110 945
When you compare bonds made between the same two atoms,
triple bonds are generally stronger and shorter than double bonds
which are generally stronger and shorter than single bonds!

17. You are done!
• Now complete the study sheet using your
freshly completed notes and this Power
Point Presentation (to obtain
electronegativity values and bond energies)