CHEMICAL FORMULAS AND CHEMICAL COMPOUNDS CHAPTER 7

CHEMICAL FORMULAS AND CHEMICAL COMPOUNDS
CHAPTER 7
I. Chemical Names and Formulas
A. Significance of a Chemical Formula
1. hydrocarbons:
2. The element’s names were given because of their discoverers, properties, place of
discovery, mineral where it can be found, among others.
3. Chemical symbols are a shorthand representation of the elements. They always
have a capital letter, and maybe, a lower case letter.
B. Monatomic Ions: ions formed from a single atom
1. Naming Monatomic Ions
a. Cations: identified by element’s name most times; add Roman numeral when
more than one charge for the same atom
b. Anions: change ending of name to -ide
C. Binary Ionic Compounds: compounds composed of two elements; metal - nonmetal
1. The charges from the ions are cross over to create the formula. The cation
(metal) is always written first
2. It has to be in lowest terms: empirical formula.
***Examples:
sodium chloride
sodium oxide
tin(II) oxide
D. Naming Binary Ionic Compounds
1. nomenclature:
magnesium chloride
iron(II) oxide
lead (IV) chloride
2. Cation is written first; anion is written last, ending in -ide.
3. The Stock System of Nomenclature
a. To name the metal, use Roman numerals if needed, written by the side of the
metal’s name.
*** Examples:
MgO AlCl 3 Al 2 O 3
Fe 2 O 3 SnO 2 CuO
4. Compounds Containing Polyatomic Ions (ternary compounds)
a. oxyanion:
b. Follows same rules for naming as the binary compounds. Use the name of

the polyatomic ion when present.
*** Examples:
Sodium nitrite
Magnesium nitrate
Fe 2 (SO 4 ) 3 FeSO 4
Ammonium chloride
Ammonium sulfide
(NH 4 ) 3 N
NH 4 CLO
Ammonium chromate
Ammonium phosphate
E. Naming Binary Molecular Compounds (two nonmetals)
1. Use Greek prefixes that give the number of atoms, not the charge.
mono 1 tetra 4 hepta 7 deca 10
di 2 penta 5 octa 8
tri 3 hexa 6 nona 9
2. Element with smaller group number is written first or, if in the same group, the one
with the highest period number. Only assign prefix if it’s more than one atom.
3. The second element ends in ide and always uses a prefix.
*** Examples:
CO
carbon dioxide
SF 6
phosphorus trichloride
PCl 5
dinitrogen monoxide
NO
dinitrogen trioxide
NO 2
dinitrogen pentoxide
F. Covalent-Network Compounds
G. Acids and Salts
1. Binary acids: consists of two elements (begins with hydrogen); solutions in water
a. H-nonmetal: hydro _____ ic acid
*** Examples:
HCl
hydroiodic acid
hydrobromic acid
H 2 S
2. Ternary acids (oxyacids): contain hydrogen, oxygen and another element
a- H- polyatomic ion
b- ate ending is changed to ic; ite ending is changed to ous.
c- DO NOT USE HYDRO!!! The word acid implies the presence of hydrogen.
*** Examples:
Perchloric acid HClO 3

Chlorous acid
HNO 4
Nitrous acid
H 2 SO 5
HClO
Nitric acid
HNO
Sulfuric acid
Sulfurous acid H 2 SO 2
3. salt:
II. Oxidation number:
1. oxidation number / oxidation state: used to indicate the distribution of electrons among the
atoms in a compound; charge on a monatomic ion.
A. Assigning Oxidation Numbers
1. Pure elements are always ______.
2. Binary molecular compounds: more electronegative element= negative charge as an
anion / less electronegative element= positive charge as a cation.
3. Fluorine= ____ in all compounds
4. Oxygen= ____ in most compounds; exception: peroxides (__) and compounds with
F (__)
5. Hydrogen= ___ with more electronegative elements; ___ with metals
6. Sum of the oxidation numbers for a neutral compound is ______.
7. Sum of the oxidation numbers for an ion _______ the charge of the ion.
B. Using Oxidation Numbers for Formulas and Names
1. We can use the Stock system instead of prefixes to name compounds.
III. Using Chemical Formulas
A. Formula Masses: _____________________________________________ (formula
unit, molecule or ion).
*** Examples:
SO 3 C 12 H 22 O 11
NaCl (NH 4 ) 2 CO 3
B. Molar Masses: mass of one mole of a compound
C. Molar Mass as a Conversion Factor
*** Examples
a) Calculate the grams in 2.00 moles of magnesium hydroxide.
b) Calculate the number of moles of iron (II) oxide in 1.204 x 10 21 formula units.

c) Calculate the number of particles in 32.0 grams of methane.
d) Calculate the number of oxygen atoms in 3 moles of carbon dioxide.
e) Calculate the number of cations in 3.55 grams of sodium oxide.
D. Percentage Composition:
*** Examples:
1. Calculate the percent hydrogen and percent oxygen in water.
2. Calculate the percent nitrogen, percent hydrogen, percent phosphorus, and percent
oxygen in ammonium phosphate.
3. Calculate the percent copper (II) sulfate and the percent water in copper (II) sulfate
pentahydrate.
*** Hydrates: substances that absorb water from the air. Written like: CuSO 4 .5H 2 O
4. Calculate the grams of phosphorus in 351 grams of calcium phosphate.
IV. Determining Chemical Formulas
1. empirical formula:
A. Calculation of Empirical Formulas
% ---> grams -----> moles -----> mole ratio ------> empirical formula
*** Examples:
1. Calculate the empirical formula for a compound that is 88.8% copper and 11.2% oxygen.
2. Calculate the empirical formula for a compound that is 38.67% calcium, 19.98% phospho
rus, and 41.25% oxygen.
3. Calculate the empirical formula for a compound that is 72.3% iron and 27.7% oxygen.
4. Calculate the empirical formula for a compound that is 83.9% Sb and 16.1% N.
5. Calculate the empirical formula for the hydrate BaCl 2 .?H 2 O with the following data: 2.66
grams of barium chloride and 0.44 grams of water.

B. Calculation of Molecular Formulas
- molecular formulaempirical
formula ---> molar mass of e.f. ----> divide weight from problem by molar
mass of empirical --->multiply this answer by empirical formula to get molecular for
mula
*** Example
1. Calculate the molecular formula for a compound that is 92.3% carbon and 7.7% hydrogen
with a molecular weight of 78 grams.
*** The molecular formula may or may not be the same as the empirical formula.