Chapters 15/6 Ionic Bonding
• 15.1 Objectives
– Use the periodic table to infer the number
of valence electrons in an atom and draw its
electron dot (Lewis dot) structure.
– Describe formation of cations from metals
and anions from nonmetals
 California Standards
 1d. Students know how to use the periodic table to
determine the number of electrons available for
bonding.
 2e. Students know how to draw Lewis dot structures.
Valence
Electrons:
ELECTRONS
AVAILABLE
FOR
BONDING
(the red ones)
Valence Electrons
• Valence electrons are electrons in the
outmost shell (energy level). They are the
electrons available for bonding.
• The number of valence electrons largely
determines the chemical properties of that
element.
• For Groups 1A-7A, the number 1-7 is the
number of valence electrons for that atom.
• Group 0 is an exception – you can think of it
as group 8A because all the noble gases
(except He) have 8 valence electrons.
Group 1 (alkali metals) have 1
valence electron
Group 2 (alkaline earth metals)
have 2 valence electrons
Group 13 elements have 3
valence electrons
Group 14 elements have 4
valence electrons
Group 15 elements have 5
valence electrons
Group 16 elements have 6
valence electrons
Group 17 (halogens) have 7
valence electrons
Group 18 (Noble gases) have 8
valence electrons, except
helium, which has only 2
Transition metals (“d” block)
have 1 or 2 valence e- . Why?
Lanthanides and actinides
(“f” block) have 1 or 2 valence
electrons
Valence Electrons
• Valence electrons are usually the only eused to bond to other atoms.
– Therefore you usually only show the valence
e- in electron dot structures.
– Electron dot structures are diagrams that
show valence e- as dots.
Generic Dot Notation
An atom’s valence electrons can be represented
by electron dot (AKA Lewis dot) notations.
1 valence e-
2 valence e-
3 valence e-
4 valence e-
X
X
X
X
5 valence e-
6 valence e-
7 valence e-
8 valence e-
X
X
X
X
?
Dot Notations – Period 2
Lewis dot notations for the valence electrons of
the elements of Period 2.
lithium
beryllium
boron
carbon
Li
Be
B
C
nitrogen
oxygen
fluorine
neon
N
O
F
Ne
Electron Dot Structures
•
•
•
Note how you draw two dots per side x four sides = 8
dots maximum.
Note how each side gets one before any side gets two.
See how the number of dots is the same for each
element within a group (column).
Octet Rule
• The Octet Rule was created by Gilbert Lewis in
1916.
• That’s why these diagrams are sometimes called
Lewis dot structures.
• In forming compounds, atoms tend to achieve the
e- configuration of a noble gas, 8 valence e-.
• An octet is a set of 8.
• Each noble gas (except He) has 8 valence
electrons in their highest principle energy level,
and the general configuration is ns2np6
(like 2s22p6 or 3s23p6)
Metallic vs. Nonmetallic Elements
• Atoms of the metallic elements (including
column 1A and 2A) tend to lose their outer
shell valence e- so they can have a complete
octet at the next energy level down.
• Atoms of nonmetallic elements tend to gain e(steal e-) or share e- with another nonmetallic
element to achieve their complete octet.
• There are exceptions but the octet rule
usually applies to most atoms in compounds.
Cations and Anions
• If an atom loses a valence e- = cation
• If an atom gains a valence e- = anion
• Metals create cations because they start with 1 to 3
e- and usually get all of the valence e- stolen so they
can get down to a full lower level octet.
• Example: Sodium loses 1 e-
• Before Na 1s2 2s2 2p6 3s1
• After Na+ 1s2 2s2 2p6
–
(note the 8 e- in the n=2 shell)
• Like
–
Ne 1s2 2s2 2p6
(Neon has 8 e- in the n=2 shell)
• The change is written as follows:
– Na·
Na+ + e-
Cations
• Cations of group 1A alkali metals +1
• Cations of group 2A alkali metals +2
·Mg·
Mg2+ + 2e• For transition metals, the charges on the
cations may vary. Note the Roman Numeral.
• Example: Fe has two: iron(II) or Fe2+
iron(III) or Fe3+
• Some atoms formed by transition metals do
not have noble-gas electron configurations
and are therefore exceptions to the octet
rule.
Exceptions
• Example: Ag Silver
• 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1 4d10 (oddball)
• Silver would have to lose 11 electrons to get down
to noble gas Krypton’s configuration. To gain
enough e- to get to Xenon’s configuration, it
would have to gain 7 electrons. Neither one is
likely.
• But if Ag loses its one 5s1 electron, then it has an
outer shell with 18 e- (the 4 shell), which is a full
shell, and relatively favorable.
• Therefore Ag always forms the Ag+ cation.
Anions
• Anions are atoms or groups with a negative
charge (extra electrons).
• Atoms of nonmetallic elements have relatively
full valence shells and are looking to steal eto make their shells full.
• Cl 1s2 2s2 2p6 3s2 3p5 neutral atom
• Cl- 1s2 2s2 2p6 3s2 3p6 anion
• Ar 1s2 2s2 2p6 3s2 3p6 now Cl- has Ar config.
Cl
+ e-
Cl-
Section 15.2 Ionic Bonding
California Standards
Students know atoms combine to form
molecules by sharing electrons to form
covalent or metallic bonds or by
exchanging electrons to form ionic bonds.
 Students know salt crystals, such as
NaCl, are repeating patterns of positive
and negative ions held together by
electrostatic attraction.
Ionic vs. Covalent Bonds
• Bonds: Forces that hold groups of atoms
together and make them function
as a unit.
 Ionic bonds – transfer of electrons
 Covalent bonds – sharing of electrons
(this will be Ch. 16)
Ionic Bonding
Na: 1s22s22p63s1
now Na+ 1s22s22p6
Cl: 1s22s22p63s23p5
now Cl- 1s22s22p63s23p6
Aluminum has three valence
e- to steal, and the Bromine
atoms would each like to
steal one e-.
So the Aluminum atom gives
up three electrons and the
Bromine atoms each receive
one.
Examples of Ionic Compounds
• Mg2+Cl21• Magnesium chloride: Magnesium loses
two electrons and each chlorine gains
one electron
• Al23+S32• Aluminum sulfide: Each aluminum loses
three electrons (six total) and each
sulfur gains two electrons (six total)
Metal
Lithium
Sodium
Potassium
Magnesium
Calcium
Barium
Aluminum
Monatomic
Cations
Li+
Na+
K+
Mg2+
Ca2+
Ba2+
Al3+
Ion name
Lithium
Sodium
Potassium
Magnesium
Calcium
Barium
Aluminum
Nonmetal
Monatomic
Anions
Fluorine
FChlorine
ClBromine
BrIodine
IOxygen
O2Sulfur
S2Nitrogen
N3Phosphorus P3-
Ion Name
Fluoride
Chloride
Bromide
Iodide
Oxide
Sulfide
Nitride
Phosphide
Recall that anions end in –ide.
Sodium Chloride crystal lattice
• Ionic compounds form
solid crystals at ordinary
temperatures.
• Ionic compounds organize
in a characteristic crystal
lattice of alternating
positive and negative ions.
All salts are ionic compounds and form crystals.
Properties of Ionic Compounds
Structure:
Crystalline solids
Melting point: Generally high
Boiling Point:
Generally high
Electrical
Excellent conductors,
Conductivity: molten and aqueous
Solubility in
Generally soluble
water:
Two K atoms lose 1 e- each => One O atom gains 2 e-
3 Mg atoms lose x 2 e- each => 2 N atoms gain 3 e- each
Ch. 6 – Ionic Naming
The Laws of Definite and Multiple Proportions
• The law of Definite Proportions states that
in samples of any chemical compound, the
masses of the elements are always in the
same proportions.
• The law of Multiple Proportions states that
whenever two elements form more than one
compound, the different masses of one
element that combine with the same mass of
the other element are in the ratio of small
whole numbers.
Ions of Representative Elements
Add – ide to anion name
Ions of metallic elements
Specific list of polyatomic ions you are
accountable to memorize for the test
-1 Ions
-2 Ions
-3 Ions
Name
Formula
Name
Formula
Name
Formula
Acetate
C2H3O2-1
Sulfite
SO3-2
Phosphate
PO4-3
Hydroxide
OH-1
Sulfate
SO4-2
Nitrate
NO3-1
Carbonate
CO3-2
Nitrite
NO2-1
Bicarbonate
(hydrogen
carbonate)
HCO3-1
+ 1 Ions
Ammonium
NH4+
Also know diatomic molecules: I2 Br2 Cl2 F2 O2
N2 H2 and also H2O, NH3 (ammonia), and CH4
(methane)