Chemical Formulas/Naming
Chapter 7
Binary Compounds
 Ionic Compounds
 (binary) with 2 monatomic ions – ions formed from a single
atom.
 Writing a Formula
1.
2.
3.
4.
Recognize positive and negative ion
Write the ions showing their charges
Adjust the number of each ion, using subscripts, so that the
total of the charges equal zero
Show the proper chemical formula without any charges
noted
Examples
 Ca+2
Cl-1
CaCl2 calcium chloride
 Al+3
S-2
Al2S3 aluminum sulfide
 Ba+2 O-2
BaO
barium oxide
 H+1
H3P
hydrogen phosphide
P-3
Naming a Compound
 Chemical nomenclature-collective term for the rules of
naming chemical compounds
 Binary compounds-use –ide ending on the anion
 Examples




BeI2 – beryllium iodide
Ba3N2 - barium nitride
NaBr – sodium bromide
Li2O – lithium oxide
 Notice the subscripts have no affect on the names!!
The criss cross method will work each time.
Notice:
+3
Al
and
Al2S3
-2
S
But if you use the crisscross method, don’t forget
to reduce if possible! (Save time, don’t write “1”)
+3
Al
and
-3
N
Al3N3 = AlN
Always write the metal first!
Polyatomic ions are groups of elements (poly =
many) that are covalently bonded together and
have an overall ionic charge.
When forming a compound, polyatomic ions work
as a group. So after criss-crossing the charges, it
will be necessary to use parentheses if the subscript
is bigger than 1.
Examples: Lithium and Arsenate
Li+ and AsO4 -3
Li3AsO4
Beryllium and Arsenate
Be+2 and AsO4 -3
Be3(AsO4)2
The 4 is not changed!
The 4 is not changed!
When naming ionic compounds with polyatomic
ions, the polyatomic ion already has a special name
that never changes.
Examples: LiC2H3O2 = lithium acetate
NH4OH = ammonium hydroxide
If the ending is a normal nonmetal,
still change the ending to “ide”!
Examples
Oxyanion – a polyatomic ion that contains oxygen





AgNO3
K2SO3
Ca3(PO4)2
Mg(ClO3)2
Ammonium sulfate
 (NH4)2SO4
 Cupric bicarbonate
 Cu(HCO3)2
 Beryllium hydroxide
 Be(OH)2
Polyvalent Metals in Compounds
 Stock Names
1.
Based on International Union of Pure and Applied
Chemistry (IUPAC)
2. Utilizes a Roman numeral system to identify the
charge of the metal




PbCO3
lead (II) carbonate
Pb(CO3)2
lead (IV) carbonate
Iron (II) bromide
FeBr2
Iron (III) bromide
FeBr3
Lead (II) carbonate
Polyvalent Metals in Compounds
 Classical Names
1.
Based on traditional (usually Latin) names for the
metals
2. Consider if it is the low state (-ous) or the high state
(-ic) for that metal
 SnBr4
 Fe2S3
tin in high state
iron in high state
stannic bromide
ferric sulfide
 Plumbous oxide Pb+2 O-2 low state PbO
 Chromous nitrate Cr+2 NO3_1 low state Cr(NO3)2
Lead (IV) oxide, PbO2 &
Lead (II) oxide, PbO
Molecular Compounds
 Writing a formula
1. Identify the “apparent” charge that the nonmetal
has based on electronegativity values
2. Traditional (prefixes) and Stock (Roman numeral)
names are both used
3. Prefixes define quantity
4. Mono-is not used for the first part of the compound
Examples
 Traditional
 Carbon tetrachloride
CCl4
 Nitrogen monoxide
NO
 Stock examples
 Nitrogen (IV) oxide
NO2
 Phosphorus (V) oxide
P2O5
Naming Molecular Compounds
1. Identify negative charge, “apparent” positive
charge, and name by either method if appropriate
 Examples
 CO O=-2, C++2
carbon monoxide or carbon (II) oxide
 PCl3 Cl=-1, P=+3 phosphorus trichloride or
phosphorus (III) chloride
Practice
 Name the following binary molecular compounds:
 SO3
 Sulfur trioxide or sulfur (VI) oxide
 ICl3
 Iodine trichloride or iodine (III) chloride
 PBr5
 Phosphorus pentabromide or phosphorus (V) bromide
More Practice
 Write formulas for the following:
 Carbon tetraiodide
 CI4
 Phosphorus (III) chloride
 PCl3
 Dinitrogen trioxide
 N2O3
Acid Names
 Binary Acids
 Consist of 2 elements, hydrogen and another nonmetal.
 Names include hydro- prefix, with –ic suffix as acid
 Examples:
 HCl
 HI
 H2S
hydrochloric acid
hydroiodic acid
hydrosulfuric acid
Acid Names
 Oxyacids
 Consist of hydrogen, oxygen, and a third nonmetal
element.
 -ite
becomes
–ous
 -ate
becomes
-ic
 No hydro prefix is used
 Examples
 HClO4 (hydrogen perchlorate) becomes perchloric
acid
 HNO2 (hydrogen nitrite) becomes nitrous acid
Naming Acids
Anion Ending
Example
Acid Name
Example
-ide
Clchloride
Hydro-(stem)-ic
acid
Hydrochloric acid
-ite
SO2-3
sulfite
(stem)-ous acid
Sulfurous acid
-ate
NO-3
(stem)-ic acid
Nitric acid
Salts
 We’ve talked about salts many times before….
 A salt is an ionic compound composed of a cation
from a base and an anion from an acid.
 Examples
CaCl2 from the base, Ca(OH)2 and the acid, HCl
Na2CO3 from the base, NaOH, and the acid, H2CO3
Practice







Name the following acids:
HF
hydrofluoric acid
HNO3
nitric acid
H2SO4
sulfuric acid
Problem of the Day







Write molecular formulas for these acids:
Chloric acid
HClO3
Hypochlorous acid
HClO
Acetic acid
HC2H3O2
Mass Spectrometry
 Identifying
unknown
molecules
can be done
using mass
spectrometry
Using Chemical Formulas
 The formula mass of any molecule, formula unit, or ion is the
sum of the average atomic masses of all atoms represented in
its formula.
 The molar mass is the formula mass with assigned g/mol units.
Also called gram formula mass.
Molar Mass Examples
1. sodium
fluoride - NaF
1 atom Na x 23.0
1 atom F x 19.0
= 23.0
= 19.0
42.0 g/mol
2. potassium carbonate - K2CO3
2 atoms K x 39.1
1 atom C x 12.0
3 atoms O x 16.0
= 78.2
= 12.0
= 48.0
138.2 g/mol
3. ammonium sulfite - (NH4)2SO3
2 atoms N x 14.0
8 atoms H x 1.0
1 atom S x 32.1
3 atoms O x 16.0
=
=
=
=
28.0
8.0
32.1
48.0
116.1 g/mol
Mole Conversions
 Uses the molar mass for each compound based on the exact chemical formula.

1. Convert 52.8 grams of calcium nitrate to moles.
a. Ca(NO3)2
b. molar mass = 164.1 g/mol
c. convert: 52.8 g x 1 mol = 0.322 moles Ca(NO3)2
164.1 g
 2. Convert 1.75 moles of cupric sulfate to grams.
a. CuSO4
b. molar mass = 159.6 g/mol
c. convert: 1.75 mol x 159.6 g
1 mol
= 279.3 grams CuSO4
Practice w/ Mole Conversions
 Convert 15 g of aluminum oxide to moles

Answer - .147 mols Al2O3
 Convert 0.06 moles of zinc hydroxide to grams

Answer – 5.96 grams Zn(OH)2
Percent Composition
 The percentage by mass of each element in a
compound is known as the percentage composition
of the compound.
 Mass of element in sample of compound
mass of sample of compound
X 100
 This gives us the percent of each element in the
compound
Example
Find the percentage composition of stannous fluoride.
a. SnF2
b. molar mass = 156.7 g/mole
c. % Sn = 118.7 g/mole = 76% Sn
156.7 g/mol
d. therefore % F = 24 %
Example 2
Find the percentage composition of ammonium dichromate.
a. (NH4)2Cr2O7
b. molar mass = 252.0 g/mol
c. % N = 28.0 g/mol = 11% N
252.0 g/mole
% H = 8.0 g/mole = 3% H
252.0 g/mole
% Cr = 104.0 g/mole = 41% Cr
252.0 g/mole
% O = 112.0 g/mole = 45% O
252.0 g/mole
Percent Composition Practice
 Find the percentage composition of calcium sulfate.
Determining Chemical
Formulas
 Empirical formula – consists of the symbols for the
elements combined in a compound, with subscripts
showing the smallest whole-number mole ratio of the
different atoms in the compound.
 For ionic compounds, the formula unit = empirical formula
 For molecular compounds, the empirical formula does not
necessarily indicate the actual numbers of atoms present
in each molecule.
 Ex. BH3 and B2H6
Steps for Determining Empirical
Formula
 If given percent data, you must assume 100 grams.
 Ex. 92.3% = 92.3 grams if we have 100 grams of the sample.
 Convert grams to moles of each element.
 Divide each element’s number of moles by the least
number of moles. This gives a ratio of elements in the
compound.
 Write the empirical formula.
 To determine a molecular formula the molar mass of the
compound must be known. (Molecular is a whole number
multiplier of empirical)
Examples
Example 2
Example 3
Empirical Formula Practice
 A compound is 10 g Li, 20 g N, and 70 g O. What is the
empirical formula?