Bonding





Outer shell = valence electrons
Octet rule - An atom is most stable if it has an outer
shell of eight electrons and no electrons of higher
energy
Elements at the left of the periodic table generally have
one or two electrons in excess of a stable noble gas
structure (octet of electrons)
These electrons are easily removed - ionization energy
is low
Such elements are electropositive
Li
Li + e
© Prentice Hall 2001
Chapter 1
1
Bonding


Elements at the right of the periodic
table generally are just one or two
electrons short of a noble gas structure
(octet of electrons)
These elements easily add electrons
F + e
© Prentice Hall 2001
F
Chapter 1
2
Ionic Bonds


Opposite charges attract
Attractions between ions hold a crystal
together and are called ionic bonds
© Prentice Hall 2001
Chapter 1
3
Covalent Bonds

F
Instead of giving up or acquiring
electrons, an atom can also achieve
eight electrons by sharing
+
F
F F
covalent bond
© Prentice Hall 2001
Chapter 1
4
Covalent Bonds

Other covalent bonds
2 H
+
© Prentice Hall 2001
H O
H
O
Chapter 1
5
Polar Covalent Bonds

In the fluorine-fluorine bond as well as in the
hydrogen-hydrogen bond, electrons are shared
equally
In hydrogen fluoride and in water, electrons are
attracted more toward one atom

© Prentice Hall 2001
Chapter 1
6
Polar Covalent Bonds


This tendency of atoms to attract electrons is
known as electronegativity
There is a continuum of bonding types
ionic bond
polar covalent bond
© Prentice Hall 2001
Chapter 1
nonpolar covalent bond
7
Electronegativity
© Prentice Hall 2001
Chapter 1
8
Lewis Structures

The chemical symbols we have been
using in which valence electrons are
shown as dots are called Lewis
structures
© Prentice Hall 2001
Chapter 1
9
Drawing Lewis Structures

Write the symbols for the elements in the
correct structural order

Consider nitric acid, HNO3
© Prentice Hall 2001
Chapter 1
10
Drawing Lewis Structures

Calculate the number of valence electrons for all
atoms in the compound
1
H
@
1 electron
=
1
3
O
@
6 electrons
=
18
1
N
@
5 electrons
=
5
=
24
Total
© Prentice Hall 2001
Chapter 1
11
Drawing Lewis Structures

Put a pair of electrons between each
symbol - at least one bond needed
between each atom and its neighbor
© Prentice Hall 2001
Chapter 1
12
Drawing Lewis Structures

Beginning with the outer atoms, place
remaining electrons in pairs around
atoms until each has eight (except for
hydrogen)
© Prentice Hall 2001
Chapter 1
13
Drawing Lewis Structures

If you run out of electrons before each
atom (other than hydrogen) has eight
electrons, move unshared pairs to form
multiple bonds
© Prentice Hall 2001
Chapter 1
14
Formal Charges

A bookkeeping system for electrons

Used to show the approximate
distribution of electron density in a
molecule or polyatomic ion
© Prentice Hall 2001
Chapter 1
15
Formal Charges

Assign each atom half of the electrons
in each pair it shares
© Prentice Hall 2001
Chapter 1
16
Formal Charges

Also give each atom all electrons from its
unshared pairs
6
7
1
6
© Prentice Hall 2001
4
Chapter 1
17
Formal Charges

Subtract the number of assigned electrons from
the number of valence electrons for an
uncombined atom of the same element
6-6=0
1-1=0
6-6=0
© Prentice Hall 2001
6 - 7 = -1
5 - 4 = +1
Chapter 1
18
Formal Charges

The algebraic sum of all formal charges
on a species (molecule or ion) must
equal the actual charge on the species



Zero for molecules
Positive for cations
Negative for anions
© Prentice Hall 2001
Chapter 1
19
Kekulé Structures


In most cases we show electron pairs
between atoms as bonds and represent
them with a dash
Also, unless there is a particular need to
show unpaired electrons, we generally do
not show them
H
O
N
O
becomes
H
© Prentice Hall 2001
O
N
Chapter 1
O
20
Condensed Structural
Formulas


Kekulé formulas also are called structural
formulas
Often, structural formulas are condensed
H
H
H
H
C
C
C
H
H
H
H
becomes
CH3CH2CH3
© Prentice Hall 2001
Chapter 1
21