Basic Inorganic Nomenclature
Compiled by Josh Schoenly, from:
Mr. Richard Hallowell’s Chemistry lecture notes (1999-2000)
Quick Study Academic. “Chemistry.” Bar Charts Inc: May 2001.
ACDLabs Freeware 5.0; 3D viewer, ChemSketch.
Wilbraham, Antony C., Dennis D. Staley, Candace J. Simpson, Michael S. Matta. Chemistry:
Second Edition. Addison Wesley Publishing: Menlo Park, Calif. 1990.
Introduction – One of the more tedious and less interesting topics about the science of
Chemistry among students would be nomenclature. While for most of the time, a student of
chemistry can simply determine an answer via stoichiometry, VSEPR theory and a general good
logical sense of reactions and electron distribution; the naming of compounds may come as an
un-welcomed guest. Most students will simply memorize names, not even give a second glance
to the logic behind the nomenclature. The skill of naming compounds, however, can be an
important aide when trying to determine the structure of a compound. This skill can also be
learned without the tediousness of creating flashcards and memorizing names of different
compounds. IUPAC, International Union of Pure and Applied Chemistry, has created a logical
set of instructions for the naming of compounds so as to create a universal language among
chemists. The student of chemistry is therefore left with a set of instructions as well as Latin
prefixes and suffixes in order to name compounds fairly easily. The purpose of this paper is to
introduce the new student of chemistry to the art of basic inorganic nomenclature in order to
create a sense of stability and order for the student in his / her journey through world of
chemistry.
I. Binary Compounds – Compounds that consist of only 2 elements
A. Metal and a nonmetal (ionic bond)
1. Molecular Formula  Name
a. Name the metal
b. Name the nonmetal
c. Add the prefix (-ide) to the end of the nonmetal
d. Examples:
(1). NaCl – Sodium Chloride
(2). AlP – Aluminum Phosphide
(3). BaBr2 – Barium Bromide
(4). Li2O – Lithium Oxide
(5). Rb2Se – Rubidium Selenide
2. Name  Molecular Formula
a. Determine the charge on the metal using a Periodic Table
b. Determine the charge on the nonmetal using a Periodic Table
c. If charges are equal, write down the two elements as is.
d. If the charges are unequal, balance using subscripts.
e. Examples:
(1). Lithium Fluoride  Li and F
Li is an alkali metal and therefore has a charge of +1  Li+1
F is a nonmetal, Halide (Halogen anion) and
has a charge of –1  F-1
Both charges are equal so: LiF
(2). Cesium Selenide  Cs and Se
Cs is an alkali metal and has a charge of +1  Cs+1
Se is a group 7 metal and has a charge of –2  Se-2
Charges are not equal so Cs must have a subscript of 2:
Cs2Se : Since subscripts are multiplied by the charge to get a
balance.
(3). Aluminum Oxide  Al and O
Al is a group 3 metal and has a charge of +3  Al+3
O is a group 6 nonmetal and has a charge of –2  O-2
Charges are not equal so Al must have a subscript of 2 and O a
subscript of 3 to balance the molecular equation: Al2O3
B. Two Nonmetals (covalent bond)
1. Molecular Formula  Name
a. Use a prefix on the first element if the first element has more than 1
atom
b. Name the first element
c. Always use a prefix for the second element
d. Name the second element
e. Add the suffix (-ide) to the end of the second element
Prefix
Number of atoms
Mono1
Di2
Tri3
Tetra4
Penta5
Hexa6
Hepta7
Octa8
Nona9
Deca10
Undeca11
Dodeca12
f. Examples:
(1). PCl3 – Phosphorus trichloride
(2). N2O5 – Dinitrogen pentoxide
(3). N2O – Dinitrogen monoxide
(4). CCl4 – Carbon tetrachloride
2. Name  Molecular Formula
a. Determine the symbols for the two elements
b. Use prefixes to put in subscripts to denote number of atoms
c. Examples
(1). Carbon dioxide – CO2
(2). Sulfur hexachloride – SCl6
C. Naming compounds with transition metals
1. Molecular Formula  Name (classical system)
a. Determine the charge on the metal
(1). Trick – Look at the charge on the anion
(2). Trick – Look at the subscripts
b. Name the metal using the classical system (below table)
(1). Metals of lower charge will have the suffix (-ous)
(2). Metals of a higher charge will have the suffix (-ic)
c. Name the nonmetal as seen in rule IA.
Element (English)
Charge
Latin root Stock
Classical
Iron (Fe)
+2
Ferrum
Iron (II)
Ferrous
Iron (Fe)
+3
Ferrum
Iron (III)
Ferric
Tin (Sn)
+2
Stannum
Tin (II)
Stannous
Tin (Sn)
+4
Stannum
Tin (IV)
Stannic
Lead (Pb)
+2
Plumbum
Lead (II)
Plumbous
Lead (Pb)
+4
Plumbum
Lead (IV)
Plumbic
Copper (Cu)
+1
Cuprum
Copper (I)
Cuprous
Copper (Cu)
+2
Cuprum
Copper (II)
Cupric
+2
Mercury (Hg2 )
+1
Mercurius
Mercury (I)
Mercurous
Mercury (Hg+2)
+2
Mercurius
Mercury (II)
Mercuric
Chromium (Cr)
+2
Chromium (II)
Chromous
Chromium (Cr)
+3
Chromium (III)
Chromic
Manganese (Mn)
+2
Mangnes
Manganese (II)
Manganous
Manganese (Mn)
+3
Mangnes
Manganese (III)
Manganic
Cobalt (Co)
+2
Cobalt (II)
Cobaltous
Cobalt (Co)
+3
Cobalt (III)
Cobaltic
Zinc (Zn)
+2
Silver (Ag)
+1
d. Examples
(1). FeO – Ferrous Oxide
(2). PbCl4 – Plumbic Chloride
(3). Hg2Se – Mercurous Selenide
(4). AgCl – Silver Chloride
2. Molecular Formula  Name (stock system)
a. Follow the classical system rules for determining charge on metal
b. Use the “stock” column of the table above to name cation
c. Name the nonmetal as seen in Rule IA.
d. Examples
(1). FeO – Iron (II) Oxide
(2). PbCl4 – Lead (IV) Chloride
(3). Hg2Se – Mercury (I) Selenide
3. Name (classical and stock system)
a. Determine charge on metal using the table above
(1). Charge in classical needs to be memorized
(2). Charge in stock gives the charge in Roman numerals
b. Balance using subscripts
II. Polyatomic ions – A group of atoms that act as an ion.
A. Tables of polyatomic ions.
1. Polyatomic anions: -1
Formula
CH3COONO2NO3ClOClO2ClO3ClO4CNOHSCNHCO3HSO3HSO4H2PO4MnO4-
Name
Acetate
Nitrite
Nitrate
Hypochlorite
Chlorite
Chlorate
Perchlorate
Cyanide
Hydroxide
Thiocyanate
Hydrogen carbonate, bicarbonate
Hydrogen sulfite, bisulfite
Hydrogen sulfate, bisulfate
Dihydrogen phosphate
Permanganate
2. Polyatomic anions: -2
Formula
CO32SO42SO32CrO42O22HPO42Cr2O72S2O32S22SiO32-
Name
Carbonate
Sulfate
Sulfite
Chromate
Peroxide
Hydrogen phosphate, biphosphate
Dichromate
Thiosulfate
Disulfide
Silicate
3. Polyatomic anions: -3
Formula
PO43PO33-
Name
Phosphate
Phosphite
4. Polyatomic anions: -4
Formula
SiO445. Polyatomic cations: +1
Formula
NH4+
H3O+
Name
Silicate
Name
Ammonium
Hydronium
B. Molecular Formula  Name
1. Locate the common groups of polyatomic ions in the formula
2. Name the polyatomic ions
3. Name the remaining cation/anion if there was only one polyatomic ion
4. Examples
a. Na3PO4 - Looking at the table, the only polyatomic ion in this
formula is phoshphate. It acts as the anion. The metal in this
equation is sodium, therefore, the compound is called sodium
phosphate.
b. Ba(NO2)2 – Looking at the table, the only polyatomic ion in this
formula is nitrite. It acts as the anion. The metal in this equation is
Barium, therefore, the compound is called barium nitrite.
c. (NH4)2SO4 – Looking at the tables, there are two polyatomic ions in
this formula. Ammonium acts as the cation while sulfate acts as the
anion. The compounds name is ammonium sulfate.
C. Name  Molecular Formula
1. Write down the formula for any polyatomic ion using the tables on polyatomic
ions as well as the charge.
2. Write down any none polyatomic ions given in the name of the formula with
their charge as well.
3. Balance the equation using subscripts so that the total charge of the compound
is 0.
4. Examples
a. Potassium dichromate – The polyatomic ion in the equation is
dichromate, which is given as: Cr2O72-. It is acting as the anion.
Potassium, or K+, is acting as the cation and is monatomic.
Balancing using subscripts gives K2Cr2O7.
b. Ammonium chloride – The polyatomic ion in the equation is
ammonium, which is given as: NH4+. It is acting as the cation.
Chloride, or Cl-, is acting as the anion and is monatomic. Balancing
using subscripts gives NH4Cl.
c. Ammonium phosphite – The polyatomic cation in the equation is
ammonium, which is given as: NH4+. The polyatomic anion in the
equation is phosphite, which is given as: PO33-. Balancing using
subscripts gives (NH4)3PO3.
III. Acids – more specifically, naming bronsted-lowry acids. Dissolved in water.
A. Binary acids – Acids with only two elements.
1. Add prefix hydro2. Name anion (since H+ is acting as the cation)
3. Add suffix –ic
4. Add word acid
5. Examples
a. HCl – Hydrochloric acid
b. H2S – Hydrosulfuric acid
c. H3P – Hydrophosphoric acid
d. HBr – Hydrobromic acid
B. Acids with anions that end in –ite.
1. Name the anion (since H+ is acting as the cation)
2. Replace the suffix –ite with –ous
3. Add word acid
4. Examples
a. H2SO3 – Sulfurous acid
b. H3PO3 – Phophorous acid
c. HNO2 – Nitrous acid
C. Acids with anions that end in –ate.
1. Name the anion (since H+ is acting as the cation)
2. Replace the suffix –ate with –ic
3. Add the word acid
4. Examples
a. H2SO4 – Sulfuric acid
b. H3PO4 – Phophoric acid
c. HNO3 – Nitric acid
d. H2CO3 – Carbonic acid
IV. Some Common Names
Formula
H2O
CaO
Ca(OH)2
NaOH
K2CO3
NaHCO3
NaCl
HCl
N2O
Na2S2O3
Chemical
Dihydrogen monoxide, hydric acid
Calcium oxide
Calcium Hydroxide
Sodium hydroxide
Potassium carbonate
Sodium bicarbonate
Sodium chloride
Hydrochloric acid
Dinitrogen monoxide
Sodium thiosulfate
Common Name
Water
Lime
Slaked lime
Lye
Potash
Baking soda, bicarbonate soda
Table salt
Muriatic acid
Laughing gas
Hypo