I. Chemical Formulas
A. A chemical formula is a combination of symbols that represents the composition of a
compound.
1. Chemical formulas show the following information:
a) The elements present in the compound.
b) The relative number of atoms of each element present in the compound.
2. Example:
a) H2O
1. Contains both hydrogen and oxygen.
2. Two atoms of hydrogen to one atom of oxygen.
II. Binary Compounds
A. Compounds that contain only two elements.
B. Molecular (covalent) versus Ionic Compounds
1) Binary molecular compounds almost always consist of two nonmetallic elements.
a. The elements in a molecular compound are electrically neutral (no charge), and they
are held together by covalent bonding.
b. Molecular compounds do not conduct electricity when they are dissolved or melted.
c. Examples:
H2O, CO, CO2, NO
O2, H2, N2, F2, Cl2  known as diatomic molecules
2) Binary ionic compounds usually consist of a metallic and a nonmetallic element.
a. The metallic ion (cation) carries a positive charge and it is held together with the
nonmetallic negative ion (anion) by an ionic bond.
b. Ionic compounds dissociate and conduct electricity when they are dissolved or
melted.
c. Examples:
NaCl, KCl, MgBr2, BaF2
III. Binary Molecular (Covalent) Compounds
A. Rules for Writing the Names of Molecular Compounds:
1. Name the elements in the order that they appear in the formula.
2. Drop the last syllable in the name of the second element and add the –ide ending.
3. Use Greek prefixes (Table 2.4, Pg. 56) with each element to indicate the number of atoms.
a) The mono- prefix should only be used with the second element in the name of the
compound.
B. Examples:
1. SiO2 
2. H2S 

3. NO
4. S4N2 
C. Writing Formulas for Molecular Compounds:
1. Represent each kind of element in a compound with the correct symbol for that element.
a) Write the symbol for the more metallic element first.
2. Use subscripts to indicate the number of atoms of each element in the compound. If there is
only one atom of a particular element, no subscript is used.
D. Examples:
1. chlorine dioxide 
2. diphosphorous tetroxide 
3. carbon monoxide 
IV. Binary Ionic Compounds
A. Determining Ionic Charge from the Periodic Table:
1. Group 1A (IA) elements form ions with a 1 charge.
2. Group 2A (IIA) elements form ions with a 2 charge.
3. Group 3A (IIIA) elements usually form ions with a 3 charge.
4. Elements in the other Groups form ions with various charges. However, the following
always holds true:
a) Nitride ion always carries a 3 charge.
b) Sulfide ion always carries a 2 charge.
c) Oxide ion always carries a 2 charge.
d) Halide ions (halogens) always carry a 1 charge.
B. Writing Formulas for Ionic Compounds:
1. Write the symbols for each element in the same order that they appear in the name.
a) The positive ion (cation) is always listed first.
2. Use subscripts for each element so that the net charge on the compound is zero.
a) No subscript is used when there is only one atom of a particular element in
the compound.
3. Examples:
a) calcium chloride  CaCl2
Net charge on the compound:
2
calcium
ion

(1  2)

0
chloride
ions
b) potassium oxide 
c) aluminum sulfide 
d) lithium fluoride 
V.
Ionic Compounds Containing Polyatomic Ions
A. A polyatomic ion is a group of atoms held together by covalent bonding (similar to molecular
compounds) that carries an overall charge.
B. Common Polyatomic Ions and their Charges  Table 2.5, Pg. 66
C. Writing Formulas for Ionic Compounds that contain Polyatomic Ions
1. Same rules as with binary ionic compounds.
a) Parentheses must be used around the polyatomic ion when a subscript other
than 1 is required.
2. Examples:
a) Magnesium nitrate  Mg(NO3)2
2
Net charge on the compound:
Magnesium
ion

(1  2)

0
nitrate
ions
b) sodium sulfate 
c) aluminum hydroxide 
VI. Ionic Compound Nomenclature
A. Rules for Writing the Names of Ionic Compounds:
1. Name the elements in the same order that they are listed in the formula.
2. Change the ending of the last element (negative ion) by adding –ide.
3. If the compound contains a polyatomic ion, use the name of the polyatomic ion as it is
given in Table 2.5, Pg. 66.
4. Ions that may have multiple charges are named by using Roman numerals to indicate the
proper charge.
Note: Several of the common ions that carry multiple charges are listed on the back of
your periodic table handout.
B. Examples:
1. Na2O

2. Al(NO3)3

3. (NH4)2SO4 
4. Cu2S

5. CuS
