Chapter 4
Chemical Bonding
Covalent Bonds
Covalent Bonds
 Both nonmetal atoms “want” to gain e- to become stable
o Electrons are shared in order to allow this to happen
o The number of e- shared depends on the element

Compounds formed with covalent bonds are neutral and follow the OCTET
rule!
o All atoms are most stable when they have the electron configuration of
a noble gases (THIS MEANS HAVING EIGHT* VALENCE
ELECTRONS)
o For nonmetals the most common exception to the octet rule is
HYDROGEN.
o Stable = Noble gas configuration

Hydrogen - will share 1 e Oxygen - will share 2 e Nitrogen - will share 3 e Chlorine - will share 1 e Carbon - will share 4 e-
Formation of a Covalent Bond
 Covalent Bonding in Hydrogen
o Examples
H2


F2

HCl

PH
Types of Covalent Bond
o Single - ONE pair of e- shared
o Double - TWO pair of e- shared
o Triple - THREE pair of e- shared
o
o Coordinate - Both electrons being shared originate from a single atom
 A “regular” covalent single bond is the result of overlap of two
half-filled orbitals.
 A coordinate covalent single bond is the result of overlap of a
filled and a vacant orbital.
 - atoms participating in cc bonding generally do not
form their normal # of covalent bonds
o Ex.: HOCl2, CO, N2O
Electron Dot Structures of HCl
Lewis Structures
o Structures which represent in a drawing the arrangement of the atoms
and the types of covalent bonds
o There are SEVEN basic steps to follow.
 Step 1: Arrange the atoms – hydrogen on periphery; use
expected bonding pattern
Step 2: Count the electrons
Step 3: Add bonds & lone pairs (each atom gets octet
 Step 4: Use multiple bonds as needed – convert one
lone pair to bonding pair
 Step 5: Exceptions to Octet rule
(H, B, P, S, noble gases)


Lewis Structures - examples
CCl4
o
o
NH4+
PBr3
o
F2
o
SO4-2
o
o H2S
o HOCl
o HO2Cl
o O3
Double Bonds
O2
o
C 2 H4
o
CO2
o
CH2CHCHCH2
o
o NO3CO3-2
o


Triple bonds
o N2
o C 2 H2
o HCN
Exceptions - Only elements in rows 3 and beyond . Why?
o PCl3 and NCl3
o
PCl5 but not NCl5
o XeF4
o SF6

Resonance
o Some molecules have measured values of bond lengths which do not
support the Lewis structure drawn for the molecule
Example: Ozone, O3
o
o To adequately represent such molecules with Lewis structures, you
must draw all possible arrangements of ELECTRONS.
Naming diatomic Molecular Compounds

“mono” is used only with the second element

Ex. Dihydrogen monoxide

Ex. Tetraphosphorus decoxide
Common vs. Chemical Names
Chemical Name
P4O10
3-D arrangement of electron pairs
o Arrangement of valence electron pairs about a central atom that
minimize repulsions between the pairs.
o Since double & triple bonds are multiple electron pairs in the same
location, they act like a single pair when determining the geometry of
the molecule
Bond Polarity & Electronegativity
o The difference in the electronegativities (ability of an atom to attract
electrons) of the two bonded atoms can be used to define the
“polarity” of the bond.
o Bond Polarity
 Bond Lengths
 Bond lengths are measured using
nucleus-nucleus distances.
o For bonds between the same two atoms:
 Single > Double > Triple
 Example: O-O; C-O; N-S
Molecular Geometry - VSEPR
o Molecular Polarity
o Determining Molecular Geometry & Polarity
- a Shortcut!
 Geometry
 Polarity
 Look @ center atom
 All bonds non-polar = NON-POLAR
 Bonds are Polar
 Only VSEPR bonding groups –
go to B.
 Has VSEPR nonbonding groups = POLAR
 Look @ attached atoms
 All attached atoms the same = NON-POLAR
 One or more different element’s atoms attached =
POLAR
Examples
- use Lewis structures to guide you
 Determine the geometry & polarity of these “molecules”:
 HF
SiBr4
CO3-2
 HCN
SeO2
CF3I
 NF3
PH3