CHEMICAL BONDING
Chapter 6
WHAT IS ELECTRONEGATIVIT Y?
WHY DOES IT MATTER?
INTRODUCTION TO
CHEMICAL BONDING
Section
5.1
SECTION 6.1
Electronegativity- The ability for an atom
to attract an electron to itself
There are 3 types of bonds
Ionic Bond
Covalent Bond
Polar Covalent Bond
T YPES OF BONDS
 Ionic bond- Electrons are transferred forming ions. These ions
are attracted to each other. The bond is a strong force
between the two atoms.
 Covalent bond- two or more atoms sharing pairs of electrons
T YPES OF BONDS (CONT)
 Polar Covalent- two or more atoms sharing a pair of electrons
unevenly
ELECTRONEGATIVIT Y AND BONDING
 The type of bond between two or more atoms depends on the
dif ference in electronegativity of the atoms.
Type of Bond
Difference in
Electronegativity
Ionic
1.8 - 3.3
Polar- Covalent
0.4 - 1.7
Covalent
0.0 - 0.3
SECTION 6.1: LEARNING CHECK
 What is the main distinction between ionic and covalent
bonding?
 How is electronegativity used in determining the ionic or
covalent character of the bonding between two elements?
 What type of bond would be expected between the following
atoms?
 Li and F
 Cu and S
 I and Bf
COVALENT BONDING
AND MOLECULAR
COMPOUNDS
Section
6.2
TERMS TO KNOW
 Molecule- a neutral group of atoms that are held
together by covalent bonds
 Chemical formula- the relative numbers of atoms of
each kind in a chemical compound by using atomic
symbols and numerical subscripts
 Octet- 8 electrons
 Single bond- 1 electron pair shared
 Double bond-
EXAMINE THESE COVALENT MOLECULES.
WHAT DO YOU NOTICE?
FORMATION OF A COVALENT BOND
Overall goal:
 Obtain an octet
 Create a stable atom
 Lowest amount to energy
CHARACTERISTICS OF THE COVALENT
BOND
 Bond length: distance
between two bonded
atoms
 Bond energy: the
energy required to
break a chemical
bond and form neutral
isolated atoms.
THE OCTET RULE
Atoms will gain, lose, or share electrons to get
8 electrons in their highest occupied energy
level
Exceptions:
 Helium and Hydrogen- only need 2 valence electrons
 Boron- has 3 electrons and is stable with 6 electrons
 Example: BF 3
 Expanded Octet- Some atoms can hold more than 8
when they are bonded to extremely electronegative
atoms.
 Example: PF 5 and SF 6
ELECTRON-DOT NOTATION
 An electron configuration notation in which only the valence
electrons of an atom of a particular element are shown.
LEWIS STRUCTURES
 Visual representation of molecules
 Element symbol- nuclei and inner-shell electrons
 Dot-pairs- non-bonding valence electrons
 Dashes- bond between two elements
HOW TO DRAW LEWIS STRUCTURES?
1. Determine the type and number of atoms in the molecule
2. Write the electron-dot notation for each type of atom in the
molecule
3. Determine the total number of valence electrons available to
be combined
4. Arrange the atoms
1.
2.
If carbon is present it is in the center
If not, the least electronegative atom is in the center
5. Add unshared pairs of electrons to each nonmetal atom (except
hydrogen) Each atom needs to be surrounded by 8 electrons.
6. Count the electrons in the structure to be sure that the number
of valence electrons used equals the number available.
7. Place non-bonding valence electrons in pairs around atoms
without an octet
EXAMPLES:
NH 3
H 2S
SiH 4
PF 3
STRUCTURAL FORMULA
 Indicates the kind, number, arrangement, and bonds but not
the unshared electrons.
H-Cl
MULTIPLE COVALENT BONDS
 Double or triple bonds
 Triple bonds are the shortest and strongest covalent
bonds
 The need for a multiple bond becomes obvious if
there are not enough valence electrons to complete
octets by adding unshared pairs.
REVIEW
 Compare the molecules H 2 NNH 2 and HNNH
RESONANCE STRUCTURES
 Bonding in molecules or ions that cannot be correctly
represented by a single Lewis structure.