Naming Compounds
Writing Formulas
and Equations
Naming Compounds
The chemical formula represents the composition of
each molecule.
In writing the chemical formula, in almost all cases
the element farthest to the left of the periodic table is
written first.
So for example the chemical formula of a compound
that contains one sulfur atom and six fluorine atoms
is SF6.
If the two elements are in the same period, the
symbol of the element of that is lower in the group
(i.e. heavier) is written first e.g. IF3.
Naming Ionic Compounds
Ionic compounds are combinations of positive
and negative ions.
In writing the chemical formula the positive ion is
written first, It is then followed by the name of the
negative ion.
Monatomic anions end in ide. Special endings
apply for polyatomic ions
Examples
NaCl Sodium chloride
BaF2 Barium Fluoride
ZnO Zinc Oxide
Names of Polyatomic Ions
with Oxygen
Polyatomic ions
usually contain
oxygen in addition
to another element.
Normally they have
a negative charge.
They end in either
"ate" or "ite"
depending on the
number of oxygen
atoms present.
ClO-
hypochlorite
ClO2ClO3-
chlorite
chlorate
ClO4-
perchlorate
NO2-
Nitrite
NO3-
Nitrate
PO33-
phosphite
PO43-
phosphate
SO32SO42-
sulfite
sulfate
Polyatomic Ion -- Exceptions
Most polyatomic ions contain oxygen
Their names end in “ite” or “ate”.
There are several exceptions
OHCNSCN-
hydroxide
cyanide
thiocyanate
Elements with Multiple Cations
1.
2.
3.
4.
When an element can form more than one cation a
Roman numeral is used to distinguish the oxidation
state of the compound.
Iron, Tin, Lead, Copper, and are common elements with
more than one cation.
Examples
PbSO4 = lead (II) sulfate
This compound is formed
from Pb2+ and SO42Pb(SO4)2 = lead (IV) sulfate This compound is formed
from Pb4+ and SO42Fe(OH)2 = iron (II) hydroxide This compound is
formed from Fe2+ and OHFe(OH)3 = iron (III) hydroxide This compound is
formed from Fe3+ and OH-
Examples of Ionic Compounds
1.
2.
3.
4.
5.
6.
7.
8.
9.
NaCl
ZnF2
KOH
Ca(NO3)2
BaSO3
Al2(SO4) 3
Ca3(PO3)2
NH4Cl
(NH4)2CO3
=
=
=
=
=
=
=
=
=
Sodium chloride
Zinc fluoride
Potassium hydroxide
Calcium nitrate
Barium Sulfite
Aluminum sulfate
Calcium phosphite
Ammonium chloride
Ammonium carbonate
Naming Covalent Compounds
When naming covalent compounds, the
name of the first element in the formula is
unchanged.
The suffix “-ide” is added to the second
element.
Often a prefix to the name of the second
element indicates the number of the element
in the compound
Examples:
SF6 – sulfur hexafluoride
P4O10 – tetraphosphorous decoxide
CO – carbon monoxide
CO2 – carbon dioxide
Covalent molecules with
multiple possibilities
A Roman Numeral is used to indicate the state
of the more positive element
Examples
1. N2O = Nitrogen (I) oxide Since oxygen has a
2- charge, the nitrogen must be 1+ to balance
the charges. Also known as dinitrogen
monoxide
2. N2O3 = Nitrogen (III) oxide Since oxygen has a
2- charge, the nitrogen must be 3+ to balance the
charges Also known as dinitrogen trioxide
Binary compounds of Hydrogen
The binary compounds of hydrogen are
special cases. They were discovered before
a convention was adopted and hence their
original names have stayed.
Water H2O is not called dihydrogen monoxide
Hydrogen forms binary compounds with almost all nonmetals except the noble gases.
Examples
HF - hydrogen fluoride
HCl - hydrogen chloride
H2S - hydrogen sulfide
Acids
When many hydrogen compounds are dissolve in
water they take on the form of an acid. Special rules
apply to acids. The “ite” suffix becomes “ous” and
the “ate” suffix becomes “ic”
HCl
Hydrochloric Acid
Cl-
Chloride
HNO2
Nitrous Acid
NO2-
Nitrite
HNO3
Nitric Acid
NO3-
Nitrate
H2SO3
Sulfurous Acid
SO32-
Sulfite
H2SO4
Sulfuric Acid
SO42-
Sulfate
H3PO3
Phosphorous Acid
PO33-
Phosphite
H3PO4
Phosphoric Acid
PO43-
Phosphate
H2CO3
Carbonic Acid
CO32-
Carbonate
Writing Formulas for Ionic
Compounds
Write the positive ion (cation) first, then the
negative ion.
The positive charges must balance the negative
charges.
Use subscripts to show how many times each
ion must appear in order for the charges to
balance. A subscript is not used if the ion
appears only once
Use parenthesis around polyatomic ions that
appear more than once in the formula
Examples
1.
2.
3.
4.
5.
6.
7.
8.
9.
Na+
Zn2+
K+
Ca2+
Fe2+
Fe3+
Ca2 +
NH4+
NH4+
and Cland Brand OHand OHand SO42and SO42and PO43and Cland CO32-
= NaCl
= ZnBr2
= KOH
= Ca(OH)2
= FeSO4
= Fe2(SO4) 3
= Ca3(PO4)2
= NH4Cl
= (NH4)2CO3
Chemical Reactions
Elements and compounds frequently
undergo chemical reactions to form new
substances
In a chemical reaction, chemical bonds are
frequently broken and new chemical
bonds are formed
Atoms are neither created nor destroyed in
an ordinary chemical change
Chemical Reactions
A balanced chemical reaction is used to
describe the process that occurs in a
chemical change.
For example: Zinc reacts with hydrochloric
acid to produce zinc chloride and
hydrogen gas.
This chemical reaction could be written as
Zn + 2 HCl  ZnCl2 + H2
Reactants and Products
In the chemical reaction
Zn + 2 HCl  ZnCl2 + H2
Reactants
Products
This shorthand way of describing a chemical
reaction is known as a chemical equation
The starting materials are shown on the left and
are known as reactants
The substances formed are shown on the right
and are known as the products
Balancing a Chemical Reaction
A proper chemical reaction must be
balanced
Zn + 2 HCl  ZnCl2 + H2
Reactants
Products
Each element must appear on both sides
of the arrow and equal number of times
Chemical reactions can be balanced by
inserting numbers in front of formulas.
These numbers are called coefficients
Balancing Chemical Reactions
Most simple equations can be balanced by
inspection
Example: Balance the following equation
BaCl2 + K3PO4  Ba3 (PO4)2 + KCl
•
•
•
There are 3 Ba on the right so we need coefficient of 3
in front of BaCl2
There are 2 PO4 on the right so we need a coefficient of
2 in front of K3PO4.
This leaves 6 K on the left so we need a coefficient of 6
in front of the KCl on the right
The balanced equation is
3 BaCl2 + 2 K3PO4  Ba3 (PO4)2 + 6 KCl
Balancing Chemical Reactions
An equation is balanced when there are the same
number and kind of atoms on both sides of the
arrow
3 BaCl2 + 2 K3PO4  Ba3(PO4)2 + 6 KCl
Reactants (Left) Products (Right)
Ba
3
Ba
3
Cl
3x2=6
Cl
6
K
2x3= 6
K
6
P
2
P
2
O
2x4=8
O
2x4 =8
State Symbols
State symbols are often added to chemical equations.
CaCO3 (s) + 2 HCl (aq)  CaCl2 (aq) + CO2 (g) + H2O (l)
Symbols
(s)
Solid
(l)
Liquid
(g)
Gas
(aq)
Aqueous (Water Solution)
Types of Reactions
There are many kinds of chemical
reactions that occur. Some are very
simple while others are very complex
and may occur in multiple steps.
A number of reactions conform to
some relatively simple patterns
Understanding and identifying these
patterns can be helpful in predicting
the products of similar reactions
Direct Combination
In a direct combination, two elements or
compounds combine to form a more
complicated product
Examples
CaO + CO2  CaCO3
2 H2 + O2
 2 H2O
FeCl2 + Cl2  FeCl3
N2 + O2
 2 NO
Decomposition
In a dcecomposition, a single compound
is broken down into two or more simplier
substances
Examples
2 KClO3  2 KCl + 3 O2
ZnCO3  ZnO + CO2
Cu(OH)2  CuO + H2O
Single Replacement
In a single replacement, one substance
(usually an element) takes the place of
another in a compound
Examples
Zn + H2SO4  ZnSO4 + H2
Cl2 + 2 KBr  2 KCl + Br2
Mg + CuCl2  MgCl2 + Cu
Double Replacement
In a double replacement, two substances
exchange places in their respective
compounds
Examples
AgNO3 + NaCl  AgCl + NaNO3
3 CaCl2 + 2 K3PO4  Ca3(PO4)2 + 6KCl
BaCl2 + Na2SO4  BaSO4 + 2NaCl
Diatomic Molecules
Certain elements
exist as diatomic
molecules in nature
H2
Hydrogen
N2
Nitrogen
F2
Fluorine
O2 Oxygen
I2
Iodine
Cl2 Chlorine
Br2 Bromine
Diatomic Molecules
Certain elements
exist as diatomic
molecules in nature
H2
Hydrogen Have
N2
Nitrogen
No
F2
Fluorine
Fear
O2 Oxygen
Of
I2
Ice
Iodine
Cl2 Chlorine
Cold
Br2 Bromine
Beer