CHAPTER 2
CHEMICAL FORMULAS &
COMPOSITION
STOICHIOMETRY
CHEMICAL FORMULAS
show the ratio of the elements present in the
molecule or compound
 He, Au, Na - monatomic
 O2, H2, Cl2 - diatomic
 O3, P4, S8 - more complex elements
 H2O, C12H22O11 – compounds ~ contains 2 or
more elements


Allotropes ~ different forms of the same element
in the same physical state. Ex. Both O2 and O3
are gases
Compound

HCl

H2O

NH3

C4H10
1 Molecule Contains
1 H atom & 1 Cl atom
ratio of elements: 1/1
 2 H atoms & 1 O atom
ratio of elements: 2/1
 1 N atom & 3 H atoms
ratio of elements: 1/3
 4 C atoms & 10 H atoms
ratio of elements: 4/10 =
2/5

IONS & IONIC COMPOUNDS
ions are atoms or groups of atoms with an
electrical charge
 two basic types of ions


positive ions or cations (+)


negative ions or anions (-)


one or more electrons less than neutral, ex. Na+
one or more electrons more than neutral, ex. Cl-
cations + anions must give neutral charge
KOH
 CaSO4
 Sr3N2
 Al(OH)3

potassium hydroxide (+1 & -1)
calcium sulfate
(+2 & -2)
strontium nitride
((2x3)&(-3x2))
aluminum hydroxide (+3 & (-1x3))
IONS & IONIC COMPOUNDS
Sodium chloride - table salt is an ionic compound
 When a soluble ionic compound is dissolved in water, it
dissociates (breaks apart) into its ions. Ex. Salt in water,
Na+ and Cl
ATOMIC WEIGHTS
Scientists made a scale for comparing the masses
of all elements. The units are arbitrary ~ called
atomic mass units (amu)
 How? Weighted average of the masses of the
constituent isotopes, its the lower number on
periodic chart
 1 amu = 1/12 mass of Carbon-12
 On the periodic table,

atomic weight of H = 1.0079 amu
 atomic weight of Ca = 40.078 amu
 Ca atom has 40 times more mass than the H atom

THE MOLE
Atoms are very small, difficult to weigh and
count, so how can we measure them accurately?
 Tha MOLE!!!
 strictly a convenience unit



amount that is large enough to see and handle in lab
mole = number of things
dozen = 12 things
 mole = 6.022 x 1023 things


Avogadro’s number = 6.022 x 1023

The atomic weight (amu) = mass of 1 mol (g), also
called molar mass - mass in grams equal to the atomic
weight

Units of atomic weight are amu/atoms or g/mol

He exists as atoms:
 4.0026 g of He atoms = 1 mol He = 6.022 x 1023
atoms He

H2 exists as molecules:
 2.0158 g of H2 molecules= 1 mol H2 = 6.022 x 1023
molecules H2 = 2 (6.022 x 1023)atoms of H = 1.204
x 1024 atoms of H
 Ex.
1) Calculate the mass of a single K atom
in grams to 4 significant figures.
 Ex.
2) Calculate the number of atoms in onemillionth of a gram of K to 4 significant
figures.
 Ex.
3) What element do you have if 3.00
moles of it weighs 80.9 g?
 Ex.
4) How many moles of Mercury are
contained in 9.84mL of Hg? How many
atoms? Specific gravity of Hg = 13.5939
FORMULA AND MOLECULAR WEIGHTS




Formula weight (for ionic compounds) formula units,
molecular weight (for covalent compounds) molecules
– sum of masses of elements in a compound
Ex. 5) The molar mass of calcium nitrate is:
Ex. 6) Calculate the formula weight of hydrated
Al2(SO4)3.18H2O. Hydrates have water attached.
Ex. 7) How many a) moles, b) molecules, c) oxygen
atoms are contained in 70.0 g of ozone, O3?
 Ex.
8) Calculate the number of Na atoms
in 36.50 g of Na2CO3.
 Ex.
9) What mass of ammonium
phosphate, (NH4)3PO4, would contain 15.0
g of N? B) If you start with 15.0 g of H,
how much (NH4)3PO4 would you make?
 Ex.
10) Calculate the number of mmol in
0.23 g of oxalic acid, (COOH)2.
PERCENT COMPOSITION



mass of each element divided by the mass of the
whole compound x 100%
Ex. 11) Find the % by mass of the elements in
hydrated FeSO4 .7H2O. Assume you have 1 mole
of the compound.
All samples of hydrated iron(II) sulfate have this
composition b/c of Law of Definite Proportions
which states that different samples of a pure
compound contains the same elements in the
same proportions.
EMPIRICAL
& MOLECULAR FORMULAS
Many times in the lab, a chemist will synthesize
a compound. To help prove what was made the
compound is sent for elemental analysis. From
this the simplest formula is found.
 empirical formula - simplest molecular formula,
shows ratios of elements but not actual numbers
of elements
 molecular formula - actual numbers of atoms of
each element in the compound
 determine empirical & molecular formulas of a
compound from percent composition


percent composition is determined experimentally
EMPIRICAL

& MOLECULAR FORMULAS
Ex. of molecular and empirical formulas

Molecular
C6H6
P4 O10
SO2
Empirical
CH
P2O5
SO2
(same)

Ex. 12) A compound is found to contain 85.63%
C and 14.37% H by mass, what is it’s empirical
formula? In another experiment its molar
mass is found to be 56.1 g/mol. What is its
molecular formula?
PURITY OF SAMPLES

The percent purity of a sample of a substance is
always represented as
% purity = mass of pure substance x 100%
mass of sample
~ mass of sample includes impurities (works just
like percentages)

Ex. 13) A bottle of sodium phosphate, Na3PO4, is
92.3% pure Na3PO4. What are the masses of
Na3PO4 and impurities in 250. g of this sample of
Na3PO4?

In 1986, Bednorz and Muller succeeded in
making the first of a series of chemical
compounds that were superconducting at
relatively high temperatures. This first
compound was La2CuO4 which superconducts at
35K. In their initial experiments, Bednorz and
Muller made only a few mg of this material. How
many La atoms are present in 3.56 mg of
La2CuO4?

Within a year after Bednorz and Muller’s initial
discovery of high temperature superconductors,
Wu and Chu had discovered a new compound,
YBa2Cu3O7, that began to superconduct at 100 K.
If we wished to make 1.00 pound of YBa2Cu3O7,
how many grams of yttrium must we buy?