Read Sections 4.1, 4.2, 4.3, 4.4, and 4.5 before viewing the slide show.
Unit 14
Binary Ionic Compounds
•Criteria for Stability of Ions (4.1)
•Lewis Symbols for Atoms (4.2)
•Octet Rule (4.4)
•Formula Writing for Binary Ionic Compounds (4.5)
•Nomenclature for Binary Ionic Compounds (4.5)
Criterion for Stability of Ions (4.1)
•Group 18, the noble
gases, is a particularly
unreactive group
•The noble gases each
have eight valence
electrons (except for
helium with two)
Noble gases
•There must be a special stability associated with eight
valence electrons.
•The driving force in electron exchange or electron
sharing to form compounds is the quest for eight valence
electrons
The Octet Rule (4.4)
•The observations on
the previous slide may
be summarized as the
Octet Rule:
•Atoms tend to gain, lose, or share electrons until they are
surrounded by eight valence electrons
(As you might expect, there will be exceptions to this.
Notably hydrogen and helium can have two valence
electrons at the most.)
Practical Aspects of the Octet Rule
Ionic Compounds
•Ionic compounds are made by the transfer of
electrons from one atom to another
•Ionic compounds are formed when a metal (left of
the blue line) accepts electrons from a nonmetal
(right of the blue line).
Practical Aspects of the Octet Rule
Molecular (or Covalent) Compounds (4.6)
•Molecular (or Covalent) compounds are made by
sharing of electrons between atoms
•Molecular compounds are formed when two or
more nonmetal atoms share electrons
(Terminology note:
The terms molecular
or covalent may be
used. Molecular is
the more current term
and the one we will
use throughout the
semester.)
Focus on Ionic Compounds
•The remainder of this unit will focus on the
ionic compounds. Molecular compounds will
be covered in the next unit.
•An abbreviated chart is provided on the next
slide that will help focus on the elements of
interest in our early discussion.
Abbreviated Periodic Chart
•The abbreviated periodic table here leaves out the
transition and inner transition elements
•The red stair-stepped line separates metals and
nonmetals
•The gray elements are metals
•The green elements
1
2
13
14
15
16
are nonmetals
H
•In reality, there is
degree of being
metal or nonmetal –
for now we aren’t
concerned with that.
17
18
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Ga
Ge
As
Se
Br
Kr
Rb
Sr
In
Sn
Sb
Te
I
Xe
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
Fr
Ra
←In the throes of discovery→
The Octet Rule and the Formation of Ions
•The octet rule states that atoms tend to lose, gain, or
share electrons to attain eight valence electrons, i.e. the
same number as the noble gases (Group 18)
•Group 1 metals have one extra electron compared to
their nearest noble
1
2
13
14
15
16
17
18
gas. Thus, they will
H
He
tend to lose one
Li
Be
B
C
N
O
F
Ne
electron to have the
Na
Mg
Al
Si
P
S
Cl
Ar
same number as the
K
Ca
Ga
Ge
As
Se
Br
Kr
nearest noble gas.
Rb
Sr
In
Sn
Sb
Te
I
Xe
•Losing one electron
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
leads to the formation
Fr
Ra
←In the throes of discovery→
of a 1+ cation.
The Octet Rule and the Formation of Ions
(continued)
•Similarly, Group 2 metals will tend to lose 2
electrons to become 2+ cations.
•Group 13-16 metals will tend to also lose valence
electrons (remember
the number of valence
1
2
13
14
15
16
17
18
electrons is the last
H
He
digit in the group
Li
Be
B
C
N
O
F
Ne
number) to become
cations. These are a Na Mg Al Si P S Cl Ar
K
Ca
Ga
Ge
As
Se
Br
Kr
little less of a sure
thing than the Groups Rb Sr In Sn Sb Te I Xe
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
1 and 2
Fr
Ra
←In the throes of discovery→
The Octet Rule and the Formation of Ions
(continued)
•Consider the Group 17 nonmetals. They are one
electron short of having the same number as their
nearest noble gas.
•As a result, they will
1
2
13
14
15
16
17
18
tend to form 1- anions. H
He
•The same approach Li Be B C N O F Ne
Na
Mg
Al
Si
P
S
Cl
Ar
is used to consider
Groups 13-16, though K Ca Ga Ge As Se Br Kr
Rb
Sr
In
Sn
Sb
Te
I
Xe
things become a little Cs Ba Tl Pb Bi Po At Rn
less absolute in
Fr
Ra
←In the throes of discovery→
Groups 13-14.
Summary of Ionic Charges
•Some of the charges suggested on the previous
slides are virtually certain things – some are a little
more flexible. The table here summarizes the
charges that you can be certain of for these elements
when combining metals and nonmetals.
1
2
13
Li+
Al3+
14
15
16
17
N3-
O2-
F-
P3-
S2-
Cl-
Na+
Mg2+
K+
Ca2+
Br-
Rb+
Sr2+
I-
Cs+
Ba2+
Fr+
Ra2+
18
Lewis Symbols (4.2)
•Lewis symbols are a method of writing the symbol for an
element and representing its valence electrons. One dot
is placed around the symbol for each valence electron.
•Examples:
Group 1(1 valence electrons): Li·
Na·
Group 2(2 valence electrons): ·Mg· ·Ca·
••
Group 16(6 valence electrons):
O
••
Group 17(7 valence electrons):
••
:F
••
K· Rb·
·Sr· ·Ba·
Cs·
••
S
••
••
:Cl
••
••
:Br
••
••
:I
••
•The actual location around the symbol is not important – once
there are more than four electrons we usually start pairing them.
Writing Formulas for Ionic Compounds (4.4)
•Consider the reaction of sodium and chlorine using Lewis symbols. Overall:
Na + 

Cl : 

Na
+
+
:
:Cl
•Things to note:
•The sodium (Na) begins with one valence electron but donates it to the chlorine. Thus
the sodium has the same number of electrons as the noble gas neon and chlorine has
the same number as argon.
•When the sodium gives up its electron it becomes a 1+ cation, Na+
•When the chlorine accepts the electron it becomes a 1- anion, Cl•Since the sodium gives up one electron and the chlorine takes on one electron, they
react in a 1:1 ratio – all electrons are accounted for.
•The formula for the compound becomes NaCl, indicating that one Na reacts with one Cl
•All of the Group 1 elements react with the Group 17 elements in the same ratio.
Writing Formulas for Ionic Compounds (4.4)
•What if it’s not a 1:1 ratio?
•Consider the reaction of barium and fluorine using Lewis symbols. Overall:

F:
Ba  +


Ba
2+
+2
:
:F
F:
•Things to note:
•The barium begins with two valence electrons but donates them to two fluorine atoms.
Thus the barium has the same number of electrons as the noble gas xenon and fluorine
has the same number as neon.
•When the barium gives up its electrons it becomes a 2+ cation, Ba2+
•When the fluorine accepts the electron it becomes a 1- anion, F•Since the barium gives up two electrons and the each fluorine takes one electron, the
atoms react in a barium:fluorine ratio of 1:2 – all electrons are accounted for.
•The formula for the compound becomes BaF2 indicating that one Ba reacts with two F
•All of the Group 2 elements react with the Group 17 elements in the same ratio.
Writing Formulas for Ionic Compounds (4.4)
•The key to formula writing for ionic compounds is to recognize that all
electrons must be accounted for. Every electron given up by the metal
must be accepted by the nonmetal.
•A relatively easy way of doing this is to consider a crossover approach.
Write the element with its ionic charge and place the charge on one as the
subscript on the other and vice versa
•Examples:
•Sr and Br
as ions
Sr2+ and Br- crossover charges Sr2+ and Br- gives
SrBr2
•Al and O
as ions
Al3+ and O2- crossover charges
Al3+ and O2- gives
Al2O3
•K and S
as ions
K+ and S2-
K+
K2O
crossover charges
and S2- gives
Nomenclature of Binary Ionic Compounds(4.5)
•Naming binary ionic compounds involves two key rules
•The metal retains its name
•The nonmetal ending changes to –ide
•For example:
BaO → barium oxide
K2S → potassium sulfide
SrCl2 → strontium chloride
Al2S3 → aluminum sulfide
Ca3N2 → calcium nitride
Metals with Multiple Possible Charges
•Besides the charges illustrated in the abbreviated periodic table below, several elements can
take more than one charge in ionic form. For example, iron (Fe) exists both as Fe2+ and Fe3+.
•This does not change the formula writing approach, but does complicate nomenclature a
little.
•For example, iron could form both FeCl2 and FeCl3. The name iron chloride would not
distinguish between them so another approach, outlined on the next slide, is necessary.
1
2
13
Li+
Al3+
14
15
16
17
N3-
O2-
F-
P3-
S2-
Cl-
Na+
Mg2+
K+
Ca2+
Br-
Rb+
Sr2+
I-
Cs+
Ba2+
Fr+
Ra2+
18
Naming Compounds Containing Metals with
Multiple Possible Charges
• Some of the key elements that show this trait include:
Fe2+
Fe3+
Cu+
Cu2+
Pb2+
Pb4+
Au+
Au3+
•The simplest approach to naming compounds with these ions in them is to write the name of
the metal followed by its charge in Roman numerals in parentheses and then the nonmetal
ending in –ide. The names of the ions above would become:
Fe2+
iron (II)
Cu+
copper(I)
Pb2+
lead(II)
Au+
gold(I)
Fe3+
iron (III)
Cu2+
copper(II)
Pb4+
lead(IV)
Au3+
gold(III)
•The compounds below would be named as indicated:
FeCl2
iron (II) chloride
FeCl3
iron (III) chloride
Cu3N2
copper(II) nitride
PbO
lead (II) oxide
PbO2
lead (IV) oxide
Au2O3
gold (III) oxide
•Notice in the bottom rows that the charge on the metal had to be figured out by working
backwards. For PbO, O is 2- so Pb must be 2+ for the sum of the charges to be zero.