GENERAL CHEMISTRY
EMPIRICAL FORMULAS OF IONIC COMPOUNDS
1.
COMMON CATIONS:
Note – the cation is always listed first in both the chemical formula and the name of a salt.
When naming the cation in a salt, the name of its element is used, unchanged.
A) Group 1 Elements: Always +1 as cations, i.e., hydrogen H+, lithium Li+, sodium Na+,
potassium K+, rubidium Rb+, cesium Cs+
B) Group 2 Elements: Always +2 as cations, i.e., beryllium Be2+, magnesium Mg2+,
calcium Ca2+, strontium Sr2+, barium Ba2+
C) Group 3 Elements: Always +3 as cations, i.e., boron B3+, aluminum Al 3+,
gallium Ga3+, etc.
D) Transition Metals: Many can occur with more than one charge.
a)
Ones with only one charge-type: (learn their charges!) silver Ag+, zinc Zn2+
b) Ones with multiple charge-types, (only the most important ones are listed):
[Note – you do not need to memorize their charges. As discussed later in this
handout, the charges can always be determined either from the name of the salt
they are contained in or by its empirical formula]
copper(I) Cu+, copper(II) Cu 2+
iron(II) Fe 2+, iron(III) Fe 3+
lead(II) Pb 2+, lead(IV) Pb 4+
2+
mercury(I) Hg 2
2 , mercury(II) Hg
tin(II) Sn 2+, tin(IV) Sn 4+
cobalt(II) Co2+, cobalt(III) Co3+
chromium(III) Cr 3+
nickel(II) Ni 2+
manganese(II) Mn 2+, manganese(IV) Mn 4+
E) Polyatomic Cations:
ammonium NH 4
hydronium H 3 O 
+
mercury(I) Hg 2
2 , (which involves 2 Hg ions that are bound together).
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2.
COMMON ANIONS:
A) Monatomic Anions: Always have -ide suffix, except in acids (see below).
a)
Group 7A (17) Elements: Always –1 as anions, i.e., hydride H  , fluoride F  ,
chloride, Cl–-, etc.
b) Group 6A (16) Elements: Always –2 as anions, i.e., oxide O 2 , sulfide S 2 , etc.
c)
Group 5A (15) Elements: Always –3 as anions, i.e., nitride N 3 , phosphide P 3 ,
etc.
B) Polyatomic Anions, (other than oxoanions): (Learn their names and formulas!)
hydroxide OH 
cyanide CN 
acetate CH3COO 
C) Polyatomic Oxoanions: Usually can have more than one form, with different
numbers of O atoms bound to a central atom. The most common one is most often
named with the -ate suffix added to the central atom’s name. Learn the ones in the list
below, then use the rules in part b to determine the names and forms for the other
possibilities.
a)
Most Common Forms: (learn all of these – both names and forms and charges!)
nitrate NO 3
chlorate ClO 3 , bromate BrO 3 , iodate IO 3
permanganate MnO 4
carbonate CO 32
sulfate SO 42 , hydrogen sulfate HSO 4
chromate CrO 42 , dichromate Cr2 O 72
phosphate PO 43 , monohydrogen phosphate HPO 42 ,
dihydrogen phosphate H 2 PO 4
b) Other Oxoanions Related to the Ones Given Above:
The names and forms are derived from the following simple rule: The most
common form for the anion is given the –ate suffix. Then the other forms are
named according to whether they have more O atoms or fewer O atoms than the
-ate form, as shown in the diagram below. Note that the charge on the ion stays
the same, regardless of how many O atoms are contained in the anion.
–O
–O
+O
hypo____ite  ______ite  ______ate  per______ate
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I.e., the -ite form has 1 less O atom than the -ate form, the hypo- -ite form has 2
less than the -ate, and the per- -ate form has 1 more than the -ate form.
Note that if you learn the formula, including the charge, of the –ate version of an
anion, you can then derive the formulas for all related versions.
Examples: The most common form for the ion, (i.e., the –ate ion), is listed first,
followed by the names and formulas of related ions.
chlorate ClO 3 , chlorite ClO 2 , hypochlorite ClO  , perchlorate ClO 4
nitrate NO 3 , nitrite NO 2
sulfate SO 4 2 , sulfite SO 3 2 , hyposulfite SO22
phosphate PO43 , phosphite PO33 , hypophosphite PO23
3.
WRITING EMPIRICAL FORMULAS:
The salt or acid which is formed when we combine cations with anions must not have an
overall net charge. [We cannot have a beaker full of + charges or full of – charges.] Start
by using the numerical charge on the cation as the subscript used for the anion, and vice
versa:
C c  A a  CaAc .
If both subscripts are evenly divisible by a common integer, then divide it through. If the
cation or anion is polyatomic and there is more than one of that ion in the chemical
formula, surround the symbol for it with parentheses. [For example, the formula
Fe3 (PO4 )2 tells us that there are 3 Fe2+ ions for every 2 PO34 ions. I.e., each unit of Fe3 (PO4 )2
contains 3 Fe atoms, 2 P atoms and 8 O atoms. If we left out the parentheses, the formula would
look like Fe3PO4 2 , would mean something very different. It would mean that each unit contains 3
Fe atoms, 1 P atom and 42 O atoms. For this reason, we must use parentheses when there is more
than one polyatomic ion in the formula!] No parentheses are needed when there is only 1 of an ion,
as shown with MgSO3, below.
Examples:
Al 3   S 2   Al 2 S 3 aluminum sulfide
Mg 2  SO32  Mg 2 ( SO3 )2  MgSO3 magnesum sulfite
Fe2  PO 43  Fe3 (PO 4 ) 2 iron (II) phosphate
Pb 4   CO 32   Pb 2 (CO 3 ) 4  Pb(CO 3 )2 lead (IV) carbonate
Na   HSO 4  NaHSO 4 sodium hydrogen sulfate
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4.
NAMES AND FORMULAS OF ACIDS: The simplest acids involve any of the anions
listed, (except hydroxide ion), with H+ as the cation. Naming the acid then depends on the
suffix used for the anion
a)
-ide anions  hydro- -ic acid
H  Cl   HCl hydrochloric acid
H   S 2   H 2 S hydrosulfuric acid
H  CN  HCN hydrocyanic acid
b) -ate anions  -ic acid
H   NO 3  HNO 3 nitric acid
H   MnO 4  HMnO 4 permanganic acid
H   SO 42   H 2 SO 4 sulfuric acid
H   PO 34   H 3 PO 4 phosphoric acid
c)
-ite anions  -ous acid
H   NO2  HNO2 nitrous acid
H   SO 32   H 2 SO 3 sulfurous acid
H   ClO 2  HClO 2 chlorous acid
H   ClO   HClO hypochlorous acid
d) Note that the oxo-acids follow the same rules for the prefixes as those shown earlier
for oxo-anions, but with –ous acid, instead or –ite, and with –ic acid, instead of –ate.
The diagram below shows the progression for both salts and acids, as the number of O
atoms is increased or decreased relative to the number present in the –ate form:
–O
–O
+O
hypo____ite
 ______ite
 ______ate
 per______ate
hypo____ous acid  ______ous acid  ______ic acid  per______ic acid
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NAMING BINARY MOLECULAR COMPOUNDS [I.e., non-metal – non-metal compounds.]
Binary compounds involve non-metal atoms from one element bound to non-metal atoms
from another element. Unlike ionic species, which do not form individual molecular units,
these compounds do form individual molecules. There can be many different possible
combinations that can occur between different pairs of elements. For example, the following
are all compounds of nitrogen and oxygen: NO, NO2, N2O, N2O4, … [The most familiar pair
are CO and CO2 – carbon monoxide and carbon dioxide.] In order to distinguish between
different possible combinations, we use a naming convention that essentially lists how many
of each type of atom occurs in the molecule.
1.
Prefixes: (Learn these!)
# Atoms
1
2
3
4
5
6
7
8
9
10
Prefix
Mono- (or none)
DiTriTetraPentaHexaHeptaOctaNonaDeca-
Note – the “a” ending on the prefix is omitted when oxides are involved– see examples
below. Also, the mono- prefix is almost always omitted for the first element listed and is
usually omitted for the second element also. The most familiar case when it is used is for
CO, which is called carbon monoxide. Note, however, that SiC is called silicon carbide, not
silicon monocarbide, (and OH– is the hydroxide ion, not the hydrogen monoxide ion).
2.
Rules For Naming Binary Molecular Compounds:
If the formula for a compound has the form of AnBm, then its name is given as follows:
(prefix for n)(name of element A) (prefix for m)(name of element B, using the –ide suffix).
The order of the elements is usually according to the order of the groups to which they
belong. Thus N and C precede O, etc.
Examples:
PCl5 – phosphorus pentachloride
NO2 – nitrogen dioxide
N2O4 – dinitrogen tetroxide
For binary acids, (those involving hydrogen with only one other element), we name them
using different forms, depending on whether we are talking about the acid in isolation, (as
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a gas or solid), or about them in water solution. If they are in water solution, we use the
rules given earlier for acids. If they are not in water solution, we use the rules given here
for binary compounds.
Examples:
HBr as a gas – hydrogen bromide
HBr dissolved in water – hydrobromic acid
3.
Exceptions to the Above Rules:
Some of the common hydrides, (compounds containing hydrogen), have common names
that are used. The ones you should be familiar with, (especially the first 3), are given in the
table below:
Formula
Common Name
H2O
Water
NH3
Ammonia
CH4
Methane
PH3
Phosphine
B2H6
Diborane
Occasionally, water is referred to as dihydrogen oxide, in accordance with the rules given
earlier. However, the other compounds in the table above are always referred to using
their common names – those given in the table. [I.e., methane is never referred to as
tetrahydrogen carbide, (or as carbon tetrahydride).]
Caution: Note the distinction between the molecular compound ammonia, NH3, and the
cation ammonium, NH4+. Do not confuse either the formulas or the names of these two
forms.
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