Chemical Bonding I: Basic
Concepts
Chapter 8
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Lewis Dot Symbols
Consists of the symbol of an element and one dot for each
valence electron in an atom of the element.
Note that (except helium) the number of valance electrons
each atom has is the same as the group number of the
element.
Ionic Bond
• Is the electrostatic
force that holds ions
together in an ionic
compound
• The metal gives the
electrons to the non
metal
• Cation is the metal
• Anion is the non metal
The Covalent Bond
• A bond in which two
electrons are shared by
two atoms.
• Only in covalent
compounds
• Non metal and non
metal
The Covalent Bond
The Covalent Bond
• Pairs of valence
electrons that are not
involved in covalent
bond formation are
called: Lone pairs
• We can only draw
Lewis structures for
compounds that have
covalent bonds
Polar Covalent Bonds
• Though atoms often form
compounds by sharing
electrons, the electrons
are not always shared
equally.
• 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
• The ability of an atom to attract toward itself the
electrons in a chemical bond
Electronegativity
The greater the difference in electronegativity, the
more polar is the bond
Lewis Structures
• Is a representation of covalent bonding in
which shared electron pairs are shown
either as lines or as pairs of dots between
two atoms, and lone pairs of dots on
individual atoms
• Only valence electrons are shown in a
Lewis structure
Writing Lewis Structures
PCl3
1. Find the sum of valence
electrons of all atoms in
the polyatomic ion or
molecule.
Keep track of the electrons:
5 + 3(7) = 26
– If it is an anion, add one
electron for each negative
charge.
– If it is a cation, subtract
one electron for each
positive charge.
Writing Lewis Structures
2. The central atom is the
least electronegative
element that isn’t
hydrogen. Connect the
outer atoms to it by
single bonds.
Keep track of the electrons:
26 − 6 = 20
Writing Lewis Structures
3. Fill the octets of the
outer atoms.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2
Writing Lewis Structures
4. Fill the octet of the
central atom.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0
Writing Lewis Structures
5. If you run out of electrons
before the central atom has
an octet…
…form multiple bonds until
it does.
The "best" Lewis structure for NO3• 1. Determine the total
number of valence
electrons in a molecule
• Draw a skeleton
• Of the 24 valence electrons in NO3-, 6 were required
to make the skeleton. Consider the remaining 18
electrons and place them so as to fill the octets of as
many atoms as possible
• Are the octets of all the atoms filled? If not then fill the
remaining octets by making multiple bonds
• Check that you have the lowest FORMAL CHARGES possible for
all the atoms, without violating the octet rule; (valence e-) - (1/2
bonding e-) - (lone electrons).
Writing Lewis Structures
• The best Lewis structure…
– …is the one with the fewest charges.
– …puts a negative charge on the most
electronegative atom.
The Concept of Resonance
• Resonance means the
use of two or more
Lewis structures to
represent a particular
molecule
• Resonance structure, is
one of two or more
Lewis structures for a
single molecule that
cannot be represented
accurately by only one
Lewis structure
Exceptions to the Octet Rule
• The expanded octet
• Odd- electron molecules
– Atoms in the 2nd period (that
– Some molecules contain an
are in families 3A-7A) cannot
odd number of electrons.
have more than 8 valance
– We need an even number
electrons around the central
of electron for complete
atom
pairing the octet rule
– Atoms of elements in and
clearly cannot be satisfied
beyond the 3rd period (that are
with all the atoms in any of
in families 3A-7A) form some
these molecules
compound in which more than • Incomplete octet
8 electron surround the central
– The number of electrons
atom
surrounding the central
atom in a stable molecule is
fewer than eight
Figure
2.10
+1
+2
The modern periodic
table.
+3 NC -3 -2 -1
0
Bond Enthalpy
• A measure of the stability of a molecule
• Thus bond enthalpy is the enthalpy change
required to break a particular bond in 1 mole of
gaseous molecules
– Ho(reaction) = sum of the bond energies of bonds being broken - sum of
the bond energies of the bonds being formed.
– Ho(reaction) = H(reactant bonds broken) - H(product bonds formed)
Example of Using Bond Energies to
Calculate Heat (enthalpy) of Reaction
Use the bond energies provided in the table above to calculate the heat (enthalpy) of
reaction, Ho, for the reaction:
CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g)
Write the balanced chemical equation, with all reactants and products in the gaseous
state.
CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g)
Write the general equation for the heat (enthalpy) of reaction:
Ho(reaction) = H(reactant bonds broken) - H(product bonds formed)
Substitute bond energy values into the equation and solve for
Ho(reaction
)
Bonds Broken
bond type
4xC–H
4 x Cl – Cl
Bonds Formed
bond energy
4 x 413= 1652
4 x 243 = 972
H(reactant bonds broken) = 2624
bond type
4 x C - Cl
4 x H - Cl
H(product bonds formed)
Ho(reaction) = 2624 - 3040 = -416 kJ mol-1
bond energy
4 x 328 = 1312
4 x 432 = 1728
= 3040