“LOOSE ENDS” OF
BONDING
Chapters 8 and 9
Chemistry: Matter and Energy
LATTICE ENERGY OF IONIC COMPOUNDS
 The lattice energy (bond energy) associated with an ionic compound is
based on coulombs law. (Don’t worry about the mathematical
statement of the law):
When comparing two ionic compounds for bond strength, look at the ions’ charges that
make up those compounds. The greater the ions’ charges, the stronger the bond. If the
two compounds contain ions of the same charge, the smaller ion pair will have a
stronger bond.
EXAMPLES OF COMPARING COMPOUNDS
FOR RELATIVE LATTICE ENERGY
 CaS will have a higher lattice energy than KCl because the former has a
2+ ion bonded to a 2- ion, and the latter has a 1+ ion bonded to a 1ion. (The greater the ions’ charges, the stronger the bond.)
 CaS will have a lower lattice energy than MgO, because even though
they both have 2+ ions bonded to 2- ions, Ca2+ and S2- are larger than
Mg2+ and O2-, respectively. (If the two compounds contain ions of the
same charge, the smaller ion pair will have a stronger bond.)
FORMAL CHARGE
 Formal charge is used to evaluate two Lewis Dot Structures that can
both be correct. It is a calculated value for each atom in the structures.
The formula is:
FC= (# of valence e- in a free atom of the element)- (# of e- assigned to the atom in the
structure)
The structure with the lowest formal charges is the best structure!! (Also, any negative
formal charge should be on the most electronegative atom in the molecule.)
EXAMPLE OF DETERMINING FORMAL
CHARGE
Click on the link below and view the video:
http://highered.mcgrawhill.com/olcweb/cgi/pluginpop.cgi?it=swf::100%25::100%25::/sites/dl/free/007
2512644/117354/05_Formal_Charge_Calculations.swf::Formal%20Charge%
20Calculations
RESONANCE
 Resonance occurs when more than one Lewis Structure can be drawn
for a molecule or ion.
The drawing above shows three resonance structures for carbonate ion
(with nonzero formal charges written in).
NOTE: The bond length of all three carbon – oxygen bonds is the same: an
intermediate length between the length of a single bond and the length of a
double bond
BOND DISSOCIATION ENERGY
(COVALENT BONDS)
 Bond length is inversely related to bond strength: longer bonds are
weaker than shorter bonds.
 Furthermore, bond length inversely related to the number of shared
pairs between two atoms: single bonds are longer than double bonds,
double bonds are longer than triple.
 THEREFORE: single bonds are longer and weaker than double bonds,
double bonds are longer and weaker than triple bonds
POLARITY OF MOLECULES
 Polar molecules result when there is an uneven distribution of the
electron density in the molecule. This results in the molecule having a
partial positive charge (δ+) on one end and a partial negative charge
(δ-) on the other end.
Another way of saying this is that the molecule has a “net dipole moment”.
MORE ON POLARITY
 Whether a molecule is polar or nonpolar is dependent on the shape of
the molecule. Molecules that are symmetrical are nonpolar and
molecules that are asymmetrical are polar.
 Ex: PCl3 is a trigonal pyramidal molecule with an unshared pair of
electrons on the P. This is an asymmetrical shape so PCl3 is polar.
 Ex: PCl5 is a trigonal bipyramidal molecule with no unshared pairs of
electrons on the central P. This is a symmetrical shape so PCl5 is
nonpolar.
HYBRIDIZATION
 Hybridization is a theory that explains how atomic orbital “hybridize” to
accommodate for bonding. To determine the hybridization on a
molecule, draw the Lewis Dot Structure and count the total number of
electron domains on the atom in question. The number of domains tells
you how many atomic orbitals were involved in forming the hybrids.
The only atomic orbitals that can be involved are s, p, and d.
DETERMINING THE HYBRIDIZATION
 In CH4, there are 4 electron domains on carbon. Thus, there are 4
atomic orbitals that hybridize – one s and three p, so we say the carbon
is sp3 hybridized.
 In C2H4, there are 3 electron domains on each carbon. Thus, there are
3 atomic orbitals that hybridize – one s and two p, so we say each
carbon is sp2 hybridized
SIGMA (σ) AND PI (π) BONDS
 Molecular orbital theory is another theory that explains what happens
to atomic orbitals when atoms bond. In this theory, overlap of s orbitals
create sigma bonds and overlap of p orbitals create one sigma and two
pi bonds.
 Single bonds are all σ bonds
 Double bonds contain one σ bond and one π bond
 Triple bonds contain one σ bond and two π bonds
BOND ORDER
 Bond order is a quantitative way of expressing bond strength. The
formula for calculating bond order is:
BO = # of shared pairs on the central atom ÷ # of atoms bonded on the central atom
EX: CO2 has a central carbon with two double bonds to oxygens (O=C=O). So there
are 4 shared pairs and two bonded atoms so the bond order is 4 ÷ 2, which is 2.
BOND ORDER AND BOND STRENGTH
 Molecules with a higher bond order contain stronger bonds.
 Carbon monoxide has carbon triple bonded to oxygen. There are three
shared pairs and one bonded atom, so the bond order is 3 ÷ 1, which is
equal to 3. Thus the bonding in CO is stronger than the bonding in
CO2.