Unit 1 Week 3 Thursday
IONIC AND COVALENT BONDING AND
NOMENCLATURE
IONIC NOMENCLATURE
NOMENCLATURE

a branch of taxonomy concerned with the
application of scientific names to taxa, based
on a particular classification scheme and in
accordance with agreed international rules and
conventions
IUPAC NOMENCLATURE

is a system of naming chemical
compounds and for describing the science
of chemistry in general. It is maintained by
the International Union of Pure and Applied
Chemistry.
IONIC BINARY COMPOUNDS

Metal non-metal -ide
IONIC BINARY COMPOUNDS

A binary compound is one that only contains
two elements in the compound. It may have
multiple atoms of each element but can only
have two elements.

An ionic binary compound is a compound
contains one metal and one non-metal. Either
element may have multiple atoms but there
can only be the two elements involved.
IONIC BINARY COMPOUNDS: IUPAC NAMING
Consists of two types of monoatomic ions
1. The metal ion is always written first and
retains its whole name
2. The non-metal is written second and has a
slight change, the ending (suffix) is changed to
–ide
 Do not write ones (Ex Na1Cl1) and if both
elements have the same number reduce to
lowest terms (Ca2O2 = CaO)

EXAMPLE
Example:
 Na+ Cl- use the cross over method NaCl
 IUPAC name: sodium chloride
 The metal name is written in full and the nonmetal has the –ide suffix added to it.

Sodium chloride

Binary compounds can be made up of more
than two ions, provided that there are only two
types of elements. Example: Al2O3

STUDY TIP: All metals in group 1 and 2 follow
periodic law. Check all the others metals when
naming.
PRACTICE

Work on the problems given on the sheet.
IONIC MULTIVALENT BINARY COMPOUNDS

Metal (charge) non-metal-ide
IONIC MULTIVALENT BINARY COMPOUNDS

A multivalent compound is one that may have
varied numbers of electrons in its valence shell.
This occurs with elements that fall outside of
the representative elements. The transition
metals are elements that commonly have
multiple valence shell electrons.

This means that they can form compounds in
various proportions.
Example:
Copper + Oxygen
 Copper and oxygen could have two different
formulas with two completely different
properties.
 CuO and Cu2O
 In order to differentiate the two compounds we
must use a different method to name them to
avoid confusion.

IONIC MULTIVALENT BINARY
COMPOUNDS: IUPAC NAMING

Same as Ionic Binary but it indicates the
metals charge

List the metal name first

After the metal name indicate the ion charge in
brackets using roman numerals.

The non-metal has -ide suffix added.

Do not write 1’s and reduce when possible

ONLY SHOW ROMAN NUMERALS FOR
MULTIVALENT COMPOUNDS

Not all transition metals are multivalent and
thus do not have roman numerals
EXAMPLE

SnO2  Sn2+O -  Sn4+O2-  tin (IV) oxide

SnO  Sn+O -
 Sn2+O2-
 tin (II) oxide

Work on the practice questions,

They are homework
POLYATOMIC IONS
MONATOMIC IONS
Ions that are composed of more than one atom
are called monatomic ions.
 Monoatomic Ion- an ion composed of only one
atom.
 We have only looked at these so far

POLYATOMIC IONS

Polyatomic Ions are ions that are composed of
more than one atom. The entire molecule
carries a charge to it.
POLYATOMIC IONS

Example-
NO3
-
SO4
O
2-
PO4
2-
O
3-
O S O
O P O
O
O
3-
BONDING

Ionic Bonding with polyatomic ions occurs in
the same manner as it does with binary atomic
molecules.
 Use
the crossover method
 Be sure that the charge that is crossed over
applies to the whole polyatomic ion. If the charge is
greater than 1, use brackets around the polyatomic
ion to indicate the number applies to the whole ion.
EXAMPLE 1
Step 1
Na + NO3-
Step 2
1+
Na
Step 3
1- 1+
NO3 Na
1NO3
Step 4
NaNO3
EXAMPLE 2
Step 1
Mg + NO3-
Step 2
2+
Mg
1NO3
Step 3
2+
Mg
1NO3
Step 4
Mg (NO3)2
PRACTICE
Mg + ClO3- =
Na + SiO32- =
Fr + HS- =
NH+ + Cl =
Li + Cl =
Be + O22- =
Sr + PO43- =
Al + BO33- =
Al + SO42- =
Na + OH- =
NAMING IONIC POLYATOMIC COMPOUNDS:
IUPAC
Multivalent: Metal (charge) polyatomic ion
 Monovalent: Metal polyatomic ion


Tertiary ionic compounds are comprised of a
metal ion and a polyatomic ion. Write the
polyatomic ions in the same way as monatomic
ions.
EXAMPLES

NaOH =
 sodium

Cu(ClO4)2 =
 copper

hydroxide
(II) perchlorate
Tin (IV) chlorate =
 Sn
(ClO3)4
OXYANIONS
Oxyanions = a polyatomic ion that includes
oxygen.
 Their name depends on how many Oxygen’s it
has in the poly atomic ion.

OXYANIONS





Example
ClO- hypochlorite ion
ClO2- chlorite ion
ClO3- chlorate ion
ClO4- perchlorate ion
hypo-_________-ite
___________-ite
__________-ate
per-_______-ate


This also applies to Br and I
PRACTICE

Do the questions on the note.
COVALENT BONDS (P36-39)
COVALENT BOND

A bond that arises when two atoms share one
or more pairs of electrons between them. The
shared electron pairs are attracted to the
nuclei of both atoms.
COVALENT BONDING
Lone pair
 A pair of valence electrons that is not involved
in bonding.
COVALENT BONDS
Unlike ionic compounds, covalent compounds
chemical formulas cannot be reduced to their
simplest forms.
 Ex C2H2 is a very different compound from
C6H6.

H
H
H
H
H
H
H
H
H
H
H
H
H
H
COVALENT BONDS
Classification of covalent bond molecules
 Molecules are classified by the number of
atoms they contain.
 -Diatomic molecules only contain two
molecules
 Ex. Carbon monoxide CO
COVALENT BONDS
-Polyatomic molecules contain more than two
molecules
 Ex Ammonia NH3

A molecule with covalent bonds can also have
ionic bonds (ammonium carbonate)

NH4+ + CO32
ELEMENTS AS MOLECULES
Some elements exist as molecules.
 diatomic H2 and O2
 polyatomic S8 and P4

In ionic bonding atoms lost or gained electrons
to form a stable octet for all of the molecules
involved. Ex
 Mg2+ Cl- = MgCl2
 Na+ Cl- = NaCl

For covalent bonds, atoms can share valence
electrons to obtain a stable octet in all of the
atoms.
FORMATION OF COVALENT BONDS

Ex: H has one valence electron. To form a
stable octet it needs two electrons. If two
Hydrogen atoms were to share their valence
electrons they would each have the required
two electrons for them to be stable.

Ex: Two Cl atoms. Each Cl atom has 7 electrons.
In order for each atom to obtain 8 electrons to
fill its outer shell it will have to share an
electron with the other Cl atom which will also
share an electron.
 This
sharing allows each Cl to have 8 electrons in
the valence shell. The other electrons are known as
known as lone pairs.
Octet Rule
 Elements will form bonds, so that in total they
all have a full valence shell like a noble gas.
BONDING CAPACITY
The number of electrons lost, gained or shared
by an atom when it bonds chemically. Nitrogen
has a bonding capacity of three (think back to
its charge if it was ionic bonding)
 Examples:

C=4
N=3
Halogen = 1
Hydrogen = 1
O=2
BONDING ARRANGEMENT
How do we decide which atom will go in the
middle of the molecule?
 The element with the highest bonding capacity
is usually in the middle.
 Ex CO2
BONDING ARRANGEMENT

If there is a choice between central atoms pick
the one with the lowest Electronegativity as the
central atom. Hydrogen is never a central atom.
Halides and oxygen are rarely central atoms but
there are exceptions.
QUESTIONS

Page 39 # 1-4
COORDINATE COVALENT BONDS

A covalent bond in which both of the shared
electrons come from the same element. (sort of
like a combination between an ionic bond and
a covalent bond)

Once the H+ is bonded all four H are equivalent.
The whole molecule shares the charge.
H
+
H N H
H
n

Ex ammonium chloride NH4Cl . It dissolves
rapidly in water to form (polyatomic ion) NH4+
and Cl- . It has many properties of a ionic
compound but it does not contain any metals.
H
+
+
H N
H N H
H
H
H
n
n
Cl
n
STRENGTH

They are strong; it requires a lot of energy to
separate them. Therefore they are relatively
stable at high temperatures. The strength of
the bond increases as the number of electrons
shared is added.

Triple bonds > double bonds > single.
DRAWING LEWIS STRUCTURES
1. Arrange the symbols of the elements of the
compounds as you would expect the atoms to
be arranged in the compound. Generally the
element with the highest bonding capacity is in
the central position. Draw SO3. (put the S in
the middle with three O’s around it)
O
S
O
O
2. Add up the number of valence electrons in
each atom.
 If it is a polyatomic ion, add one electron for
each unit of negative charge, or subtract one
for each unit of positive charge.
 Then add up the number of electrons needed
for all of the elements to have a stable octet.
Subtract the total number need by the total
number and divide by two. This gives you the
number of bonds needed.

What we need
 4(8)

What we have
 3(6)

= 32
O + 6 S = 24
Number of bonds
 (32-24)

/ 2 = 8/2 = 4 bonds
It is divided by two because there are two
electrons needed for each bond.
3. Place one pair of electrons between each
adjacent pair of elements.
 Fill
in the electrons around SO3 in part 1.
O
S
O
O
5. Fill in the remaining valence electrons. Make sure that there is
the right number of electrons in total and that there are the
right number of bonds.
O
S
O
O
4. Check to see that there are enough electrons
to form the right number of bonds.
 In
this case add an extra set of electrons between
one of the Oxygen’s and the Sulfur.
O
S
O
O
6. Make the shared electrons into dashes and
remove the lone pairs.
O
S
O
O
7. If representing a polyatomic ion place square
brackets around the entire structure and write
the charge outside the brackets.
STRUCTURAL FORMULA
The lone pairs are not indicated and thus only
the bonds sharing electrons are shown.
 Atoms that only need one valence electron will
generally form single bonds where atoms that
need two electrons will generally form two
bonds.
 O2 as an example. Draw the full Lewis
structure.
 N2 with its triple bonds.


More examples on the board
HOME WORK
Ionic Nomenclature practice
 Ionic polyatomic nomenclature
 Lewis structures
 Covalent nomenclature –opps, next class
 Take home quiz – hand in


REMINDER – Hand in Line spectra Lab