1. Lesson 8
Chemical Bonding & Nomenclature
You should view all lectures as a slide show. In the toolbar
above, click “Slide Show” and “From Beginning” to start. Hit
the space bar to forward to next slide or item.
1
*Read: Watch the Lecture as a slideshow, hit the space bar to
move forward or just left click.
1
Topics:
Chemical Bonds
Lewis Dot Structures
Molecular Geometry
Polar Bonds
Naming Compounds
Writing Chemical Formulas
Naming Molecular Compounds
Lesson 8: Chemical Bonding and Nomenclature
2

2. 2
Chemical Bonds
Sodium (Na) is a silver-colored metal that reacts so
violently with water that flames are produced when sodium
gets wet.
Chlorine (Cl) is a greenish-colored gas that is so poisonous that
it was used as a weapon in World War I.
When the electrons of sodium metal and chlorine gas interact,
the compound sodium chloride (NaCl) is formed, which is table
salt.
Since the body is over 90% water, ingesting sodium metal
would literally set someone on fire!
If we make a compound out of Na, we can make something
totally different such as the table salt used to season our food.
3

3. Chemical Bonds
Atoms can interact with each other to form new substances
called
compounds.
Compounds are formed when electrons in an atom’s last energy
level interact and form chemical bonds. A chemical bond is an
attractive force between atoms that holds them together.
An atom’s outermost energy level of electrons is called the
valence shell (or valence level) and the electrons in the valence
shell are called valence electrons.
4

4. Na Atom
Cl Atom
Sodium has 1 valence electron in the valence shell.
Chlorine has 7 valence electron in the valence shell.
Chemical Bonds
In ionic bonds, metals always lose electrons to nonmetals and
become positive (or cations).
In ionic bonds, nonmetals always attract electrons from metals
and become negative (anions).
Ionic compounds are neutral compounds made up of cations and
anions.
Covalent bonds are formed between nonmetals and electrons are
shared so no ions are formed.
Cl
nonmetal
Na
metal
O
nonmetal
O
nonmetal
Metal + Nonmetal = Ionic Compound
Nonmetal + Nonmetal = Covalent Compound

5. Chemical Bonds: Ionic
The octet rule states that atoms will gain, lose, or share
valence electrons in a way that will give each atom eight
electrons in their valence shell.
Na has 1 electron in its valence shell and Cl has 7 electrons in
its valence shell. Cl needs 1 valence electron to have 8 in its
outer shell.
If Na transfers its 1 valence electron to Cl, Na’s second energy
level becomes the valence shell, which already has 8 electrons.
Now, Na has a positive 1 charge and Cl has a negative 1 charge.
But, the charge on NaCl is zero (the charges cancel each other
out).
6
Na Atom
Cl Atom
Ionic Bond

6. Chemical Bonds: Ionic
An Ionic bond is formed when there is a transfer of electrons
from a metal to a nonmetal. Compounds formed by ionic bonds
are called ionic compounds or formula units.
When Mg bonds with O, Mg transfers its 2 valence electrons to
Oxygen. Oxygen then goes from 6 valence electrons in its
outermost shell to 8 electrons.
Now, both atoms have 8 electrons in its outermost shell.
Mg will have a positive 2 charge and O will have a negative 2
charge. But the compound is neutral because a +2 plus -2 equals
zero.
7
Mg Atom
O Atom
Ionic Bond
2
2

7. Chemical Bonds: Covalent
A Covalent bond is formed when there is a sharing of electrons
between two nonmetals. Compounds formed by covalent bonds
are called covalent compounds, molecular compounds or
molecules.
All atoms want to have 8 valence electrons (octet rule) and
oxygen has 6 valence electrons and needs to gain 2 more to have
8 electrons.
When 2 oxygen atoms bond, the valence shell of both oxygen
atoms overlap so that the two atoms are sharing electrons
instead of giving up or transferring electrons.
Now, both atoms “feel like” they have 8 electrons in their
outermost shell.
8
O Atom
O Atom
Covalent Bond
2 nonmetals form covalent bonds

8. Sharing 4 electrons, 2 from each atom
Lewis Dot Structures
Chemical bonds between atoms involve valence electrons.
Remember, valence electrons are the electrons in the last energy
level of an atom.
There are two ways to determine the number of valence
electrons an atom has.
First, draw the Bohr Model and count the number electrons in
the last shell (level).
Second, the group number is also the number of valence
electrons.
For example, atoms in Group 1A have 1 valence electron, Group
2A atoms have 2 valence electrons and Group 3A atoms have 3
valence electrons, and so on.
Chlorine is in Group 7A and has 7 valence electrons.
There are exceptions to these rules, but we will not learn those
in this course.
9
Valence Electrons
Cl is in Group 7A, so Cl has 7 valence electrons.
Chlorine atom
Lewis Dot Structures
A Lewis dot structure is a drawing that shows the structure of a
compound and/or the position of valence electrons around the

9. nucleus of an atom.
Lewis dot structures also show how atoms are bonded together.
Let’s start with Lewis dot structures of atoms.
The periodic table on the right shows Lewis dot structure of
various atoms.
Notice that the number of red dots (which represent valence
electrons) around each chemical symbol is equal to the group
number.
Notice that Cl has 3 paired electrons and 1 unpaired electron.
Paired electrons are 2 electrons that appear side-by-side (2 dots
paired) in Lewis structures and the unpaired electrons are
electrons appear single (one dot) in Lewis structures.
Now, let’s practice drawing Lewis structures for atoms.
10
Lewis dot structures of atoms
Chlorine atom
Lewis Dot Structures
To draw a Lewis dot structure:
Write down the chemical symbol
Determine the number of valence electrons for the atom
Place the valence electrons around the chemical symbol of the
element in a clockwise motion.
11
Na
Mg
Al
Cl
P
S

10. Lewis structures for atoms
Si
Ar

11. 1 unpaired e-
2 unpaired e-
3 unpaired e-
4 unpaired e-
3 unpaired e-
1 pair of e-
2 unpaired e-
2 pairs of e-
1 unpaired e-
3 pairs of e-
4 pairs of e-
Group 1A =
1 valence e-
Group 2A =
2 valence e-
Group 3A =
3 valence e-
Group 4A =
4 valence e-
Group 5A =
5 valence e-
Group 6A =
6 valence e-
Group 7A =
7 valence e-
Group 8A =
8 valence e-
C = Group 4A = 4e- = 4e-
H = Group 1A = 1e- x 4 = 4e-

12. Lewis Dot Structures
To draw a Lewis dot structure for a molecule:
Write down the chemical formula.
The chemical formula is a formula that shows the type and
number of atoms in a compound.
Determine Total Valence Electrons.
Add up number of valence electrons from each atom.
This number is how many total electrons in the final Lewis
structure.
Draw a skeletal structure (starting structure).
Start by making the first atom in chemical formula, the central
atom and connect the other atoms around the central atoms
using single lines.
A single line between 2 atoms represent 1 chemical bond and
every single line represents 2 electrons.
1 single line = 1 chemical bond = 2 electrons
The first atom in the chemical formula is usually the center
atom (except H).
Determine the missing electrons.
Subtract the total number of valence electrons in the
skeletal structure from the total valence electrons (found in step
2). If the answer is zero and all atoms have 8e- around them
(except H), the structure is complete.
Lewis structures for Molecules
CH4
Chemical formula
Total Valence e- = 8e-
(there are 4 H atoms, so multiply by 4)
C
H
H

13. H
H
2 e-
2 e-
2 e-
2 e-
Step 4: 8 (from step 2) – 8 = zero
2 e-
2 e-
2 e-
N = Group 5A = 5e- = 5e-
H = Group 1A = 1e- x 3 = 3e-
Lewis Dot Structures
To draw a Lewis dot structure for a molecule:
Write down the chemical formula.
Determine Total Valence Electrons.
Draw a skeletal structure (starting structure).
Determine the missing electrons:
Subtract the total number of valence electrons in the skeletal
structure from the total valence electrons (found in step 2). If
the
answer is zero and all atoms have 8e- around them (except H),
the
structure is complete.
Place the missing electrons in the Lewis structure.
Place the missing electrons on the outermost atoms first until
each atom has 8 e- around it. H can only have 2 e-, so go to next
step, Step 5b.
If you cannot put the missing e- on the outermost atoms …,
place them on the central atom.
Count the total valence electrons in the structure, if it is equal

14. to Step 2 (8e-) and all atoms have 8e- around them, the structure
is complete.
Lewis structures for Molecules
NH3
Chemical formula
Total Valence e- = 8e-
N
H
H
H
Step 5a: e- cannot go on the H’s because H can only have 2 e-
around it. Each line is 2 e-.
Step 4: 8 (from step 2) – 6 = 2 e-
2e-
4e-
6e-
8e-
Step 5b: place the 2e- on the central atom.
2 e-
2 e-
O = Group 6A = 6e- = 6e-
H = Group 1A = 1e- x 2 = 2e-
Lewis Dot Structures
To draw a Lewis dot structure for a molecule:
Write down the chemical formula.
Determine Total Valence Electrons.

15. Draw a skeletal structure (starting structure).
Determine the missing electrons:
Subtract the total number of valence electrons in the skeletal
structure from the total valence electrons (found in step 2). If
the
answer is zero and all atoms have 8e- around them, the structure
is complete..
Place the missing electrons in the Lewis structure.
Place the missing electrons on the outermost atoms first until
each atom has 8 e- around it. H can only have 2 e-, so go to next
step, Step 5b.
If you cannot put the missing e- on the outermost atoms …,
place them on the central atom. Place the 4e- on O.
Count the total valence electrons in the structure, if it is equal
to Step 2 (8e-) and all atoms have 8e- around them, the structure
is complete.
Lewis structures for Molecules
H2O
Chemical formula
Total Valence e- = 8e-
O
H
H
Step 4: 8 (from step 2) – 4 = 4 e-
Step 5a: e- cannot go on the H’s because H can only have 2 e-
around it. Each line is 2 e-.
Step 5b: place the 4e- on the central atom.
2e-
4e-
6e-

16. 8e-
2 e-
2 e-
2 e-
N = Group 5A = 5e- = 5e-
F = Group 7A = 7e- x 3 = 21e-
Lewis Dot Structures
To draw a Lewis dot structure for a molecule:
Write down the chemical formula.
Determine Total Valence Electrons.
Draw a skeletal structure (starting structure).
Determine the missing electrons:
Place the missing electrons in the Lewis structure.
Place the missing electrons on the outermost atoms first until
each atom has 8 e- around it. There are 20 missing e-. You can
only put 6 electrons around the 3 F atoms (18 e-), which means
there are 2 more e- to place.
….when you have finished placing e- around outermost atoms
and there are still e- to place), place them on the central atom.
Count the total valence electrons in the structure, if it is equal
to Step 2 (26e-) and all atoms have 8e- around them, the
structure is complete.
Lewis structures for Molecules
NF3
Chemical formula
Total Valence e- = 26e-
N
F

17. F
F
Step 5a: Place missing e- on outermost atoms (F) up to 8e-
(each F already has 2 e- around it).
Step 4: 26 (from step 2) – 6 = 20 e- missing
Step 5b: place the 2e- on the central atom.
2e-
4e-
6e-
8e-
C = Group 4A = 4e- x 2 = 8e-
H = Group 1A = 1e- x 6 = 6e-

18. Lewis Dot Structures
To draw a Lewis dot structure for a molecule:
Write down the chemical formula.
Determine Total Valence Electrons.
Draw a skeletal structure (starting structure).
If there is more than one C, make them both central atoms
and place the other atoms around the C’s.
Determine the missing electrons.
Subtract the total number of valence electrons in the
skeletal structure from the total valence electrons (found in step
2). If the answer is zero and all atoms have 8e- around them,
the structure is complete.
Lewis structures for Molecules
C2H6
Chemical formula
Total Valence e- = 14e-
Step 4: 14 (from step 2) – 14 = 0 e- missing
C
H
H
H
2 e-
2 e-
2 e-
2 e-
C
H
H

19. H
2 e-
2 e-
2 e-
C = Group 4A = 4e- x 2 = 8e-
H = Group 1A = 1e- x 4 = 4e-
Lewis Dot Structures
To draw a Lewis dot structure for a molecule:
Write down the chemical formula.
Determine Total Valence Electrons.
Draw a skeletal structure (starting structure).
Determine the missing electrons.
Place the missing electrons in the Lewis structure.
Place the missing electrons on the outermost atoms first until
each atom has 8 e- around it. H can only have 2 e-, so go to next
step, Step 5b.
If you cannot put the missing e- on the outermost atoms …,
place them the central atoms. Place the 2e- on one of the
carbons.
Count the total valence electrons in the structure, if it is equal
to Step 2 (12e-) and if all atoms have 8e- around them, the
structure is complete. It is not complete because one carbon has
only 6 e- around it.
Lewis structures for Molecules
C2H4
Chemical formula
Total Valence e- = 12e-
Step 5a: cannot place missing e- on H
Step 4: 12 (from step 2) – 10 = 2 e- missing
C

20. H
H
2 e-
2 e-
2 e-
C
H
H
2 e-
2 e-
Step 5b: place the 2e- on the a central atom.
Step 5c: one carbon has 8e- and the other one only has 6e-
around it. So, not complete.
Lewis Dot Structures
continued from previous slide…
Write down the chemical formula.
Determine Total Valence Electrons.
Draw a skeletal structure (starting structure).
Determine the missing electrons.
Place the missing electrons in the Lewis structure.
If all atoms do not have electrons, shift electrons from one atom
between atoms that are deficient in order to make double (2
lines) or triple bonds (3 lines). Shift the 2e- on the carbon
between the 2 central carbons. Count the total valence electrons
again and make sure all atoms have 8e- around them.
C = Group 4A = 4e- x 2 = 8e-

21. H = Group 1A = 1e- x 4 = 4e-
Lewis structures for Molecules
C2H4
Chemical formula
Total Valence e- = 12e-
C
H
H
2 e-
2 e-
2 e-
C
H
H
2 e-
2 e-
Step 5d: shift the lone pair on the carbon between the two
carbons to form a double bond. Now recount, they both have 8e-
.
C
C
H
H
2 e-

22. 2 e-
2 e-
H
H
2 e-
2 e-
2 e-
Complete (12 e- in the structure)
One double bond formed
C = Group 4A = 4e- x 2 = 8e-
H = Group 1A = 1e- x 2 = 2e-
Lewis Dot Structures
To draw a Lewis dot structure for a molecule:
Write down the chemical formula.
Determine Total Valence Electrons.
Draw a skeletal structure (starting structure).
Determine the missing electrons.
Place the missing electrons in the Lewis structure.
Place the missing electrons on the outermost atoms first until
each atom has 8 e- around it. H can only have 2 e-, so go to next
step, Step 5b.
If you cannot put the missing e- on the outermost atoms …,
place them the central atoms. Place the 4e- on the carbons.
Count the total valence electrons in the structure, if it is equal
to Step 2 (10e-) and if all atoms have 8e- around them, the

23. structure is complete. It is not complete because both carbons
have only 6 e- around it.
Lewis structures for Molecules
C2H2
Chemical formula
Total Valence e- = 10e-
Step 5a: cannot place missing e- on H
Step 4: 10 (from step 2) – 6 = 4 e- missing
C
H
2 e-
2 e-
C
H
2 e-
Step 5b: place the 4e- on the a central atom.
Step 5c: both carbons have 6e- . So, not complete.
C
H
2 e-
H
2 e-

24. Lewis Dot Structures
continued from previous slide…
Write down the chemical formula.
Determine Total Valence Electrons.
Draw a skeletal structure (starting structure).
Determine the missing electrons.
Place the missing electrons in the Lewis structure.
If all atoms do not have electrons, shift electrons from one atom
between atoms that are deficient in order to make double (2
lines) or triple bonds (3 lines). Shift the 2e- on the carbon
between the 2 central carbons. Count the total valence electrons
again and make sure all atoms have 8e- around them.
C
C = Group 4A = 4e- x 2 = 8e-
H = Group 1A = 1e- x 2 = 2e-
Lewis structures for Molecules
C2H2
Chemical formula
Total Valence e- = 10e-
C
H
2 e-
2 e-
C
H
2 e-
Step 5d: shift the lone pairs on the carbon between the two
carbons to form a triple bond. Now recount, they both have 8e-.

25. 2 e-
Complete (10 e- in the structure)
2 e-
2 e-
Lewis Dot Structures
21
Summary Lewis structures for Molecules
Shared pair of electrons = pairs that are bonded
N
F
F
F

26. Lone pair of electrons = pairs not bonded
Cl
Lewis Dot Structures
Write down the chemical formula.
Draw the Lewis structure for each individual atom.
Transfer the valence e- from the metal to the nonmetal.
Place the charges on the individual atoms. That’s it.
(we will not focus on Lewis structures for ionic compounds)
Lewis structures for Formula Units (Ionic Compounds)

27. Chemical formula
MgCl2
Cl
Mg
Cl
Cl
Mg
2+
−
−

28. The is the final Lewis dot structure for this ionic compound.
Polar Bonds
A nonmetal can attract electrons away from a metal because a
nonmetal has a greater electronegativity than a metal.
Electronegativity is a measure of an atom’s attraction for
another atom’s bonding electrons.
Electronegativity increases as you move from left to right on
the periodic table and it also increases as you move up a group.
Electronegativity
Electronegativity increases
Electronegativity increases
Polar Bonds
If H and F formed a bond, the bond would be polar because F
has a higher electronegativity than H.
F pulls the electrons in the bond towards itself so that the F side
of the H-F molecule has more electrons and becomes slightly
negative and the other side becomes slightly positive. This is a
polar covalent bond (not an ionic bond).
Electronegativity
Electronegativity increases
Electronegativity increases
F

29. H
α+
α−
Partially positive and partially negative symbols.
Polar Bonds
In general, if 2 different atoms are bonded together, the bond is
polar.
In general, if identical atoms are bonded together, the bond is
nonpolar.
Review the examples below. The dots are not included and the
single line represents chemical bonds.
Cl
H
C
C
H
H
S
C
Polar bond
Nonpolar bond
Nonpolar bond

30. Polar bond
α+
α−
α+
α−
Molecular Geometry
The molecular geometry of a molecule is the shape of the
molecule.
Below, you will find the 5 main molecular shapes and examples
of molecules that would assume that shape.

31. 1. Linear shape
BeF2, CO2, BeCl2, BH2, HCN, C2H2, N2
3. Tetrahedral Molecule shape
CH4, CCl44, SO42-, C2H6, SiF4
5. Trigonal pyramidal shape
NH3, NF3, PCl3
4. Bent molecule shape
H2O, SO2, ClO2, OF2
2. Trigonal planar shape
BF3, BH3, C2H4
Molecular Geometry

32. To determine the shape of a molecule, first draw the Lewis dot
structure and then determine the number of bonded atoms and
lone pairs of electrons.
1. Linear
If a molecule has 2 bonded atoms or 2 atoms bonded to a central

33. atom, the molecule is linear.
3. Tetrahedral
If the molecule has 4 bonded atoms to a central atom, the
molecule is tetrahedral.
2. Trigonal planar
If the molecule has 3 bonded atoms to a central atom, the
molecule is trigonal planar. Boron is an exception to the octet
rule and has 6 e- around it instead of 8e-.
4. Trigonal pyramidal
If the molecule has 3 bonded atoms and 1 lone pair, the
molecule is trigonal pyramidal.
5. Bent
If the molecule has 2 bonded atoms and 2 lone pairs, the
molecule is bent.
Naming Compounds
Ionic vs Molecular
It is important to learn how to name compounds and how to
write chemical formulas.
You will learn how to name Ionic compounds and how to name
molecular compounds – the rules are different.
Ionic compounds consists of metals and nonmetals (ions)
Molecular compounds consist of nonmetals (no ions, no metals).

34. 28
H2O
Molecular
compound
CO2
Molecular
compound
MgCl2
Ionic
compound
Ca(PO4)2
Ionic
compound
Naming Ionic Compounds
First, you must be able to distinguish between atomic and
polyatomic ions.
An atomic ion is an ion formed by one atom.
A polyatomic ion is an ion formed by more than one atom
grouped together.
29
Atomic vs Polyatomic
Chloride ion
atomic
PO4

35. 3−
SO4
2−
ClO3
2−
P
3−
S
2−
Cl
−
Phophide ion
atomic
Sulfide ion
atomic
Chlorate ion
polyatomic
Phosphate ion
polyatomic
Sulfate ion
polyatomic
Naming Ionic Compounds
Atomic Ions
The charges on atomic atoms represent the number of electrons
the atom gains or loses to form an ion.
The charge on an atomic ion can be determined by the Group
number. Review the picture.
Atoms in Group 1A, 2A and 3A form 1+, 2+, and 3+ charges
respectively. These atoms lose 1, 2, and 3 e-.
Atoms in groups 5A, 6A, and 7A form 3-, 2-, and 1- charges
respectively. These atoms are gaining 3, 2 and 1 e- respectively

36. 30
Chloride ion
Gains 1e-
Cl
−
P
3−
Phosphide ion
Gains 3e-
S
2−
Sulfide ion
Gains 2e-
Al
3+
Aluminum ion
Loses 3e-
Na
+
Sodium ion
loses 1e-
Mg
2+
Magnesium ion
Loses 2e-
Naming Ionic Compounds

37. Atomic Ions
Some atomic ions (especially transitional metals) form more
than one charge. Use Roman numerals to distinguish between
charges.
For example, Copper can be Cu+ or Cu2+ and Iron can be Fe 2+
or Fe 3+. So, these ions are written with Roman numerals
Copper (I) for Cu+, Copper (II) for Cu2+ , Iron (II) for Fe 2+,
and Iron (III) for Fe 3+.
You will not have to memorize atomic ions with multiple
charges.
31
Fe
2+
Iron (II) ion
Loses 2e-
Cu
+
Copper (I) ion
loses 1e-
Cu
2+
Copper (II) ion
Loses 2e-
Fe
3+
Iron (III) ion
Loses 3e-

38. Naming Ionic Compounds
Atomic ions: Notice that some atoms can from more than one
charge. Use Roman numerals to name these ions. You will not
have to memorize ions with multiple charges.
Naming Ionic Compounds
The table below provides a list of polyatomic ions.
You will not need to memorize these ions. You will just learn
how to use them.
Polyatomic ions
IonNameIonNameNH4+Ammonium ionCN−Cyanide
ionC2H3O2−Acetate ionNO3−Nitrate ionCO32−Carbonate
ionNO2−Nitrite ionHCO3−Hydrogen carbonate
ionSO42−Sulfate ionClO3−Chlorite ionSO32−Sulfite
ionClO4−Chlorate ionPO33−Phosphate ionOH−Hydroxide
ionPO23−Phosphite ion
Naming Ionic Compounds
To name an atomic ionic compound, do the following:
Name the cation first
The name of the cation is simply the name of the element + the
word ion
The cation is always written first in the chemical formula
Name the anion
To name the anion, write the name of the element and change

39. it’s ending to “ide” then add the word ion.
Name the ionic compound
Combine the name of the cation and anion without the word ion.
34
Rules for naming atomic ionic compounds
MgCl2
cation
anion
Magnesium ion
Chlorine
Chloride ion
Magnesium
Magnesium Chloride
drop the word ion
atomic ionic compound will be the term used for ionic
compounds made up of atomic ions (no polyatomic ions)
Naming Ionic Compounds
Name the cation first
The name of the cation is simply the name of the element + the
word ion
The cation is always written first in the chemical formula
Name the anion
To name the anion, write the name of the element and change
it’s ending to “ide” then add the word ion.
Name the ionic compound
Combine the name of the cation and anion without the word ion.

40. 35
Rules for naming atomic ionic compounds
Ca3N2
cation
anion
Calcium ion
Nitrogen
Nitride ion
Calcium
Calcium Nitride
drop the word ion
Naming Ionic Compounds
Polyatomic ions are placed in parenthesis if there is more than
one in the chemical formula. There are two PO43- in
Ca3(PO4)2, so parenthesis must be used to show that there are
two of these groups in the compound.
To name a polyatomic ionic compound, follow the same rules
for atomic ionic compounds:
Name the cation first
If polyatomic, look up the name in your notes.
Name the anion
If polyatomic, look up the name in your notes.
Name the ionic compound
Combine the name of the cation and anion without the word ion.
36

41. Rules for naming polyatomic ionic compounds
Ca3(PO4)2
cation
polyatomic anion
Calcium ion
Phosphate ion
Calcium
Calcium Phosphate
drop the word ion
polyatomic ionic compound will be the term used for ionic
compounds made up of one or more polyatomic ions
Naming Ionic Compounds
Notice that charges do not appear in the chemical formulas.
To name a polyatomic ionic compound, follow the same rules
for atomic ionic compounds:
Name the cation first
If polyatomic, look up the name in your notes.
Name the anion
If polyatomic, look up the name in your notes.
Name the ionic compound
Combine the name of the cation and anion without the word ion.
37
Rules for naming polyatomic ionic compounds
(NH4)2CO3
polyatomic cation
polyatomic anion

42. Ammonium ion
Carbonate ion
Ammonium Carbonate
drop the word ion
Naming Ionic Compounds
To name a polyatomic ionic compound, follow the same rules
for atomic ionic compounds:
Name the cation first
If polyatomic, look up the name in your notes.
Name the anion
If polyatomic, look up the name in your notes.
Name the ionic compound
Combine the name of the cation and anion without the word ion.
The original charge on the cation in the chemical formula is the
subscript on the anion and the charge on the anion is the
subscript on the cation. This is how we can determine the Fe
has a 3+ charge. This will be clear later in the lesson.
38
Rules for naming polyatomic ionic compounds
Fe2(SO4)3
polyatomic cation
polyatomic anion
Iron (III) ion
Sulfate ion
Iron (III) Sulfate
drop the word ion

43. Writing Chemical Formulas: Ionic
Now, let’s work backwards. You are given the name and must
write the chemical formula.
To write the chemical formula for ionic compounds:
Write the chemical symbols for the cation and anion and include
the charges.
For atomic ions, determine the charge by the group number or
by the Roman number given
For polyatomic ions, look up the charge.
Crisscross the charges
The charge on the cation will become the subscript on the anion
and the charge on the anion will become the subscript on the
cation.
Use parenthesis for more than one Polyatomic ion and no
parenthesis for atomic ions.
Remove the charges
Charges are not written in the chemical formula
No subscripts are written if charges are the same on the cation
and the anion.
39
Rules for writing chemical formulas: ionic
Fe2(SO4)3
polyatomic cation
atomic anion
Iron (III) ion
Sulfate ion
Iron (III) Sulfate

44. Use parenthesis for sulfate
Fe
SO4
3+
2−
Writing Chemical Formulas: Ionic
To write the chemical formula for ionic compounds:
Write the chemical symbols for the cation and anion and include
the charges.
For atomic ions, determine the charge by the group number or
by the Roman number given
For polyatomic ions, look up the charge.
Crisscross the Charges
The charge on the cation will become the subscript on the anion
and the charge on the anion will become the subscript on the
cation.
Use Parenthesis for more than one Polyatomic ion
Remove the charges
Charges are not written in the chemical formula
No subscripts are written if charges are the same on the cation
and the anion.
40
Rules for writing chemical formulas: ionic
Ca3N2
atomic cation
atomic anion
Calcium ion
Nitride ion
Calcium Nitride

45. No parenthesis for atomic ions
Ca
N
2+
3−
Writing Chemical Formulas: Ionic
To write the chemical formula for ionic compounds:
Write the chemical symbols for the cation and anion and include
the charges.
For atomic ions, determine the charge by the group number or
by the Roman number given
For polyatomic ions, look up the charge.
Crisscross the Charges
The charge on the cation will become the subscript on the anion
and the charge on the anion will become the subscript on the
cation.
Use Parenthesis for more than one Polyatomic ion
Remove the charges
Charges are not written in the chemical formula
No subscripts are written if charges are the same on the cation
and the anion.
41
Rules for writing chemical formulas: ionic
CaS
atomic cation
atomic anion

46. Calcium ion
Sulfide ion
Calcium Sulfide
No subscripts written in formula if charges on ions are the same
Ca
S
2+
2−
Naming Molecular Compounds
Always use Greek prefixes to name covalent compounds
To write the name of a molecular compound from the chemical
formula, do the following:
Determine the appropriate Greek prefix to use for each element
based on the number of atoms in the chemical formula.
Write the name of the compound.
To name the first element, simply use element’s name and add
the appropriate prefix. Tetra means 4, so use tetra in front of
phosphorus.
To name of the second element, use the element’s root name
with the “ide” ending and add the appropriate prefix. Deca
means 10, so use deca in front of oxide. Drop the vowel on the
prefix when the name yields two vowels together (decoxide and
not decaoxide)
If there is only one of the first atom, you do no need to use
“mono”, if there is one of the second atom, you will use mono
as a prefix.
42
Rules for naming molecular compounds
P4O10

47. 4 Phosphorus
atoms = tetra
10 oxygen
atoms = deca
Tetraphosphorus
=
decoxide
Naming Molecular Compounds
To write the name of a molecular compound from the chemical
formula, do the following:
Determine the appropriate Greek prefix to use for each element
based on the number of atoms in the chemical formula.
Write the name of the compound.
To name the first element, simply use element’s name and add
the appropriate prefix. No prefix needed if there is just one of
the first element.
To name of the second element, use the element’s root name
with the “ide” ending and add the appropriate prefix. Mono
means 1, so use mono in front of oxide. Drop the vowel on the
prefix when the name yields two vowels together (monoxide and
not monooxide)
If there is only one of the first atom, you do no need to use
“mono”, if there is one of the second atom, you will use mono
as a prefix.
43
Rules for naming molecular compounds
HBr
1 hydrogen atom = no prefix
1 bromine atom = mono

48. Hydrogen
=
monoxide
Writing Chemical Formulas: Molecular
To write the chemical formula of a molecular compound from
the name, do the following:
Write the chemical symbols of the first and second element
based on the name given.
Add subscripts to the first and second element in the chemical
formula based on the appropriate Greek prefixes given in the
name.
Never write the number 1 as a subscript.
Tetra means 4, so there will be 4 chlorine atoms
44
Rules for writing chemical formulas: molecular
C
1 Carbon
4 chlorine atoms
Carbon Tetrachloride
=
Cl
4
Writing Chemical Formulas: Molecular

49. To write the chemical formula of a molecular compound from
the name, do the following:
Write the chemical symbols of the first and second element
based on the name given.
Add subscripts to the first and second element in the chemical
formula based on the appropriate Greek prefixes given in the
name.
Di means 2, so there will be 2 hydrogen and 2 oxygen atoms.
45
Rules for writing chemical formulas: molecular
H
2 Hydrogen atoms
2 Oxygen
atoms
Dihydrogen Dioxide
=
O
2
2
The End
46