Chapter 9
Chemical Bondingg
Chemical Bonds




A CHEMICAL BOND is the attractive force that
holds atoms together in molecules and ionic
compounds
Chemical bonds are electronic in nature - electrons
attracted to another atom’s nucleus
IONIC bonds involve the transfer of one or more
electrons from one atom (or group of atoms) to
another
COVALENT bonds involve the sharing of one or
more electron pairs between 2 atoms
Ionic & Covalent Compounds
Ionic Compounds
Covalent Compounds
g meltingg points
p
High
Gases,, liquids,
q
, solid with
low melting points
Soluble in polar solvents
(eg. Water)
Insoluble in polar
solvents (eg. Water)
Insoluble in non-polar
solvents (oil or gasoline)
Soluble in non-polar
solvents (gasoline)
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Ionic & Covalent Compounds
Ionic Compounds
Covalent Compounds
q ) cmpds
p
Molten ((liquid)
conduct electricity
p do NOT
Molten cmpds
conduct electricity
Aqueous solutions conduct Aqueous solutions do
electricity
NOT conduct electricity
Elements have widely
different electronegativities
Elements have similar
electronegativities
Electronegativity (EN)

Electronegativity (EN) is a measure of the
relative tendency of an atom to attract
electrons to itself when it is chemically
combined with another atom

EN is expressed on the Pauling scale (Pauling
devised a method for calculating EN)

Fluorine is most electronegative and is assigned
a value of 4.0

Oxygen is the 2nd most electronegative (3.5)
4.5
4.0
Transition Metals
– filling d orbitals
3.5
3.0
{
Electronegativity ((Pauling Scale)
Electronegativity vs Atomic No
25
2.5
2.0
1.5
1.0
0.5
0.0
0
10
20
30
40
Atomic Number
2
Electronegativies
Table 66--3 in text
Increases
Using Electronegativities
“Rules of Thumb
Thumb””
A-Bx molecule
If ENB – ENA  ~2.0 Bond is IONIC
(electrons are transferred, not shared)
 If ENB – ENA < ~0.4
04 B
Bond
d iis COVALENT


(molecular – electrons are shared by both atoms)

If ENB – ENA is 0.4 – 1.9, the bond is
considered Polar Covalent (sometimes called
PARTIALLY IONIC)
◦ The higher the difference, the more ionic
Using Electronegativities
CH4
2.5 - 2.1 = 0.4
Covalent
NaCl
3.0 – 1.0 = 2.0
Ionic
HCl
3.0 – 2.1 = 0.9
Polar Covalent
3
Lewis Dot Representations
of Atoms

Chemical bonding usually
involves ONLY the electrons in
the outer shells
(oh, no – not quantum mechanics
again)


These are called VALENCE
electrons
Electrons in the s and p orbitals
for the “Representative”
Elements
Group
# Valence
Electrons
IA
1
IIA
IIIA
2
3
IVA
4
VA
5
VIA
6
VIIA
7
VIIIA –
except He
8
Lewis Dot Representations

Lewis Dot representations have a “dot”
for each outer shell or valence electron
with the chemical symbol
◦ Electron “Bookkeeping
p g System”
y
◦ Write the symbol and use the four sides of
the symbol
◦ Dot count is important
Lewis Dot Representations of
Atoms
Period 2 Elements
3 unpaired
electrons
Each element in the Group will have a similar dot
representation. For example, Ca looks like Be  Ca:
4
Ionic Bonding


IONIC Bonding is attraction of oppositely
charged ions to form a solid
◦ An ion carries an electrical charge (+ or -)
In IONIC crystals every internal ion is surrounded by
i
ions
off the
h opposite
i charge
h
+
-
Monoatomic ion –
only one atom
A cation carries a
positive charge
Polyatomic ion –
more than 1 atom
A anion carries a
negative charge
Sodium Chloride
An example of Ionic Bonding
Cl- is “isoelectronic” with Ar
Na+ is “isoelectronic” with Ne
Lithium Oxide
O2- is “isoelectronic” with Ne
Li+ is “isoelectronic” with He
5
Calcium Oxide
 Draw
the Lewis dot structure for
calcium oxide (CaO)
Lattice Energy of Ionic Compounds

Lattice energy (U) is the energy
required to completely separate one
mole of a solid ionic compound into
ggaseous ions

Lattice energy increases as charge on
ions increase and/or as the distance
between the ions decreases
◦ 2+ ions have greater lattice energy than 1+
Lattice Energy of Ionic Compounds
6
Covalent Bonds
Covalent bonds involve the SHARING of
a pair of electrons between 2 atoms
 Atomic orbitals (probability of finding the
e-) o
bot atoms
ato s must
ust ove
ap to share
s ae
of both
overlap
electrons
 When 2 pairs of electrons are shared
it is a called a double covalent bond
 When 3 pairs of electrons are shared it is
called a triple covalent bond

Lewis Dot Representations

You can draw Lewis Dot Representations of
covalent compounds
H • + • H  H : H (H – H)
The “Octet” Rule
Representative (not transition metals)
elements usually react to yield a stable
noble gas electron configuration.
 Noble
ob e gases (except
(e cept He)
e) have
ave 8 electrons
e ect o s
in their outer shell and are reasonably
inert (not very reactive)

7
Electrons
In covalently bonded atoms, some
electrons are shared, but others are not
 The shared electrons are called
O
G ELECTRONS
C O S
BONDING
 Unshared electrons occur in pairs (spin
paired in one orbital) and called LONE
PAIRS.

Valence Electrons
 F2
Number of valence electrons =
7 x 2F = 14 e
 HF
Number of valence electrons =
1 x 1H + 7 x 1F = 8e
 C6 H 6
Number of valence electrons =
6 x 1H + 4 x 6C = 30e
Valence Electrons

Calculate the number of valence
electrons for H2O and CO2
8
Drawing Lewis Dot Structures

Select a symmetrical "skeleton" for the molecule
or polyatomic ion.
◦ The least electronegative element (except H) is usually
the central atom.
◦ Most times O atoms do not bond to each other
(except : O2, O3, peroxides, superoxides)
◦ Most ternary oxyacids (HNO3, HClO4, H2SO4),
hydrogen bonds to a single O atom, not the central
atoms
◦ For ions or molecules with more than 1 central atom,
choose a symmetrical skeleton.
Drawing Lewis Dot Structures
1.
2.
Count the number of valence electrons
available from all the atoms
For polyatomic ions, add or subtract the
necessary number of elections
◦ If the charge is -2,
2, add 2 electrons to the valance
electron count
3.
4.
Draw a single covalent bond between the
central atom and each surrounding atom
Complete the electron “octets” for atoms
bonded to the central atom
◦ Hydrogen only gets “2” electrons
Drawing Lewis Dot Structures
5.
6.
7.
Non bonding electrons belonging to the
central atom or surrounding atoms are shown
as “lone pairs”
The total number of electrons must be equal
to the total of from the “valence” electron
count
If the central atom has fewer than eight
electrons, try adding double or triple bonds to
central atom using lone pairs from the
surrounding atoms to complete the octet of
the central atom
9
Lewis Structure Tips







Most representative elements follow the octet
rule, except hydrogen (2 electrons)
H only forms one bond to another element
C always forms a total of 4 bonds
In neutral compounds,
compounds N forms 3 bonds & O
forms 2.
Non-metals can form single, double, triple
bonds, but never quadruple bonds.
C forms double or triple bonds to C, N, O, or S
O can form double bonds with many other
elements.
Lewis Structure
CS2
 Number of valance electrons =
1 x 4C + 2 x 6S = 16e

S C S
Lewis Structure


CCl4
Number of valance electrons =
4 x 1C + 7 x 4Cl = 32e
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Lewis Dot Structures
Draw the Lewis Dot Structure for
CO2
Lewis Structures
Carbonate Ion - CO32 Number of valance electrons =
4 x 1C + 6 x 3O + 2 (2- Charge) = 24e

22or
Resonance
Sometimes a molecule
can have more than
one equivalent Lewis
structure.
 When
the
Wh this
hi happens
h
h
bonding is said to
exhibit RESONANCE.
 The real structure is
the average or hybrid
of the equivalent
structures.
2-

2-
2-
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Carbonate Ion
2-
This means the 2
extra electrons are
y distributed
evenly
between all 3
carbon oxygen
bonds - not
localized to any
one C – O bond.
Formal Charge

Formal Charge is the hypothetical charge on an
atom in a molecule or polyatomic ion
FC = {{Groupp Number – (# of bonds + # of unshared e)}
)}
CO32FCc = 4 – (4+0) = 0
2-
FCO =(3) 6 – 1 (2 + 4) - 2 (1 + 6)
= -2
Formal Charge



Sometimes, we can draw several different Lewis
structures that obey the “octet” rule
Formal charges can be used to decide which
Lewis structure is the most reasonable
The most stable structure is where
◦ The atoms bear the smallest formal charges
◦ Any negative charges reside on the more
electronegative atoms
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Formal Charge – CO2
Grp #
6
4
6
6
4
6
# Bonds +
# unshared e-
6
4
6
7
4
5
Formal Charge
0
0
0
-1
0
+1
Limitations of Octet Rule




Be has 2 valence shell electrons (1s22s2). Be forms
2 bonds to other atoms. N = 4 per Be.
Group IIIA for 3 bonds with most other atoms (N
= 6 per B
B, or Al…)
Al )
Compounds containing odd number of electrons
(NO – 11 electrons)
(NO2 – 17 electrons)
3rd period elements or higher can accommodate
more than eight electrons
Limitations of Octet Rule
When the central element needs to share in
more than 8 electrons to hold all the
available electrons (A). Valence shell is
expanded.
d d
 PCl5
 Number of valance electrons =
7 x 5Cl + 5 x 1P = 40e
 But in order to bond to 5 Cl atoms, Cl and P
need to share 10 electrons

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Range of Bonding
The degree of electron sharing depends
upon the electronegativity difference
(EN)
o po a covalent
cova e t bonding
bo
g (equally
(equa y
 Nonpolar
shared electrons) is one extreme
 Ionic bonding (transfer of the electron) is
the other extreme
 There is a range of bonding between
these 2 extremes

Bond Energies/Enthalpy
The amount of energy necessary to break
one mole of bonds in a gaseous covalent
substance to form products in the gaseous
at constant temperature
p
and ppressure
 H2(g)  2H(g) ΔHorxn = + 436 kJ/mol
 HCl(g)  H(g) + Cl(g) ΔHorxn = + 432 kJ/mol
 Note that the energies are positive – energy
input to break the bond

Bond Energies/Enthalpy
Average Bond Energies for single bonds
KJ/mol
H
C
N
O
F
H
C
N
O
F
436
414
393
460
568
347
276
351
193
176
142
157
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Bond Energies/Enthalpy
•
Note that the bond energy of a double bond
is slightly less than that of an equivalent
single bond
Bond Energies/Enthalpy
•
You can estimate the enthalpy of a reaction
from bond energies
H0 = total energy
gy input
p – total energy
gy released
= BE(reactants) – BE(products)
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