GENERAL CHEMISTRY
EMPIRICAL FORMULAS OF IONIC COMPOUNDS
1.
COMMON CATIONS:
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., berillium 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: Most occur with more than one charge.
a)
Ones with only one charge-type: silver Ag+, zinc Zn2+
b) Ones with multiple charge-types, (only the most important ones are listed):
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(IV) Co4+
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 
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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  , etc.
b) Group 6A (16) Elements: Always –2 as anions, i.e., oxide O 2 , sulfide S 2 , etc.
Group 5A (15) Elements: Always –3 as anions, i.e., nitride N 3 , phosphide P 3 ,
etc.
B) Polyatomic Anions, (other than oxoanions): hydroxide OH  , cyanide CN  ,
acetate CH3COO 
c)
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 of the other
possibilities.
a)
Most Common Forms:
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 this
form, as shown in the diagram below. Note that the charge on the ions stays the
same, regardless of how many O atoms are contained in the anion.
–O
–O
+O
hypo____ite  ______ite  ______ate  per______ate
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.
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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 ,etc.
3.
WRITING EMPIRICAL FORMULAS:
The salt or acid which is formed by combining cations with anions cannot have a net
charge. Start by using the numerical charge on the cation as the subscript for the anion,
and vice versa: C  c  A  a  C a A c . 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 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 ion. I.e., each unit of
Fe3 (PO4 )2 contains 3 Fe atoms, 2 P atoms and 8 O atoms. Writing it as 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, the parentheses are very important, when more than one polyatomic ion is
present in the formula!]
Examples:
Al 3   S 2   Al 2 S 3 aluminum sulfide
Mg 2  SO32  Mg 2 ( SO3 )2  MgSO3 magnesum sulfite
Fe 2  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
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 of 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
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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 similar rules to those shown earlier for oxo-anions:
–O
–O
+O
hypo____ite
 ______ite
 ______ate
 per______ate
hypo____ous acid  ______ous acid  ______ic acid  per______ic acid
NAMING BINARY MOLECULAR COMPOUNDS [Those containing non-metal atoms only.]
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 for oxides – 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).
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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 with –ide suffix).
Examples:
PCl5 – phosphorus pentachloride
NO2 – nitrogen dioxide
N2O4 – dinitrogen tetroxide
For binary acids, (those involving hydrogen with only one other element), we name them
differently depending on whether we are talking about them in isolation or 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 for binary compounds.
Examples:
HBr as a gas – hydrogen bromide
HBr in water – hydrobromic acid
3.
Exceptions to the Above Rules:
Many of the common hydrides, (compounds containing hydrogen), have common names
that are used instead of using the rules above. The ones you should be familiar with,
(especially the first 4), are given in the table below:
Formula
H2O
NH3
CH4
H2S
PH3
B2H6
Common Name
Water
Ammonia
Methane
Hydrogen sulfide
Phosphine
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 almost never referred to as
tetrahydrogen carbide.]
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