Chapter 6: Chemical
Bonds
6.1 – Ionic Bonding
Stable Electron Configurations

Atoms are stable when the highest energy
level is filled with electrons


Atom is not likely to react
8 valence electrons is the “magic” number


Atom is stable when there are 8 electrons in the
outer orbit
Exception – Helium (He)
 Why?
 Helium has only 2 total electrons
Stable Electron Configurations
Neon
•8 Valence
Electrons
•Stable
Aluminum
•3 Valence
Electrons
•Unstable
Helium
•2 Valence
Electrons
•Stable
Stable Electron Configurations

Remember how to find the number of
valence electrons?
2
1
2
1
2
Stable Electron Configurations
Electron Dot Diagram – model of an
atom that shows number of valence
electrons
Atomic symbol with dots outside the symbol
that represent valence electrons
:
. Cl
:
Na
.

:

Ionic Bonds

Ionic Bond – electrons are transferred
from one atom to another


Electrons are lost/gained
Ions are formed

Ion – charged particle
 Ex: Na+, Cl-
Na Cl
:
:
Protons: 17
Electrons: 17
Neutrons: 18
.
: .
:
.
Protons: 11
Electrons: 11
Neutrons: 12
. Cl
:
+
:
Na
+
_
Protons: 11
Electrons: 10
Neutrons: 12
Protons: 17
Electrons: 18
Neutrons: 18
Ionic Bonds

2 types of ions:

Cation – positively charged ion


Na+
Anion – negatively charged ion


ClNamed by using part of the element’s name and
adding the suffix -ide
Chlorine
Chloride
Ionic Bonds

Chemical Bond – force that holds atoms
or ions together as a unit
Ionic Bonds

Ionization Energy – amount of energy
used to remove an electron

Increases left to right and bottom to top of
periodic table
Ionic Compounds

Ionic Compound – compound formed
from an ionic bond


Atoms are held together by the charge of the
atoms after electrons are lost/gained
Chemical Formula – notation that shows
what elements a compound contains and
the ratio of atoms or ions of these
elements in the compound

Ex: NaCl (1:1 ratio)
MgCl2 (1:2 ratio)
Ionic Compounds

Some ionic compounds form crystal
lattices


Ions are in a fixed position called a lattice
Crystals are formed when the particles of a
solid form a lattice
Writing Ionic Formulas
Ex: calcium chloride



Write the symbol for each
atom
Identify the oxidation
numbers
Do oxidation numbers cancel
out?


YES – write symbols of atoms
NO – balance charges by using
subscripts (usually use the criss
cross method)
Ca
Cl
+2
-1
Ca+2
Cl-1
Ca+2
Cl-1 2
Ca+2
Cl-1
1
CaCl2
2
Writing Ionic Formulas

Rewrite the symbols after the numbers
have been switched




Make sure you include the new subscripts
Do NOT include the + or – symbols after you
criss-cross
If there is a 1 as a subscript, just write the
symbol and do NOT write the 1
If the subscripts are the SAME number (i.e.
Ca2O2), simplify the formula by removing the
numbers (CaO)
Writing Ionic Formulas

NOTE: If the name of an ion ends in “ite” or “-ate”, this is a polyatomic ions (an
ion that has more than one atom. There
is a list of polyatomic ions on the back of
your reference table. Use this list for the
symbols and oxidation numbers of those
ions. Keep the polyatomic ions in
parentheses while writing the formula.
Writing Ionic Formulas
Example 1: potassium sulfate
K
(SO4)
K+1
(SO4)-2
K2
(SO4)1
K2(SO4)
Writing Ionic Formulas

NOTE: Most transition metals will have
roman numerals written with the name of
that element. This represents the
oxidation number of the element.

Example:
iron (II)
iron (III)
Fe+2
Fe+3
Chapter 6: Chemical
Bonds
6.2 – Covalent Bonding
Covalent Bonds

Covalent Bond – chemical bond in which
atoms share electrons

Atoms are held together by the attraction of
the protons in the nucleus and the shared
electrons orbiting the nucleus
Molecule – neutral group of atoms that
are joined together by one or more
covalent bonds
 Diatomic Molecules – 2 atoms of the same
element


H2, F2, N2, Cl2, Br2, I2
Covalent Bonds
Diatomic Elements
Covalent Bonds

Single Bond – 1 bond between diatomic
molecules
N–N

Double Bond – 2 bonds between
diatomic molecules
N=N

Triple Bond – 3 bonds between diatomic
molecules
N≡N
Sharing of Electrons

Polar Covalent Bond – electrons not
shared equally

Creates partial charge



Atom with stronger attraction has partial – charge
Atom with weaker attraction has partial + charge
Ex: water, hydrogen fluoride
H2O
HF
Sharing of Electrons

Non-polar Covalent Bond – electrons are
shared equally


No charge present
Ex: Carbon dioxide
CO2
Attraction Between Molecules

Attractions between polar molecules are
stronger than attractions between nonpolar molecules

Why?


Polar molecules have a charge
Non-polar molecules have no charge