Chapter 5

Chapter 5
Elaine Teto
The Mole
• the formula mass in
grams of a substance
contains one mole of
• A mole is also equivalent
to Avogadro’s number
(6.022 x 1023.)
• The formula would be:
mole = n = mass (gr) /
molar mass (gr/mole)
• Returning back to what
was previously stated a
mole of anything has
Avogadro’s number of
objects in it
• For example, a mole of
water (H2O) has 6.022 x
1023 water molecules.
• a mole of 12C atoms will
weigh 12 grams, a mole
of 197Au atoms will
weigh 197 grams.
Atomic mass, formula weight and
molar mass
• atomic mass unit (amu) is
defined by setting the
mass of one C atom
equal to 12 amu
• We can determine the
amu experimentally to be
1.66 x 10-24 grams.
• These weights can be
derived from the periodic
• the formula weight, which
is the sum of the atomic
weight of each atom in a
chemical formula.
• For example, the formula
weight for NaOH is (22.99
+ 15.99 + 1.0079) 39.988
• When chemical formulas
equal the molecular
formula than the formula
weight is also known as
the molecular weight.
• the molar mass, the
mass in grams of 1
mole of a substance,
can be determined.
• The molar mass is
numerically equal to
the formula weight.
For example one
mole of water (FW =
18 amu) has a molar
mass of 18 grams.
• Stoichiometry is the math behind chemistry.
• Given enough information, one can use
stoichiometry to calculate masses, moles, and
percents within a chemical equation.
• when considering balancing equations one must
remember the Law of Conservation of Mass,
which states that matter can neither be created
nor destroyed
Balancing Equations
• Example 1: Al + Fe3O4 
Al2O3 + Fe
• When balancing equations it is
important to take one atoms at
a time and work with each
individually. Trying to balance
an entire equation at once can
be difficult.
8 Al + 3 Fe3O4 
4 Al2O3 + 9 Fe
• Example 2: CH4 + 2O2 
CO2 + H2O + Heat
CH4 +2O2  CO2
+ 2H2O + Heat
Empirical and Molecular formulas
• The molecular formula is the
form of the term as it would
appear in a chemical equation.
The empirical formula and the
molecular formula can be the
same, or the molecular formula
can be any positive integer
multiple of the empirical
• Examples of empirical
formulas: AgBr, Na2S,
Examples of molecular
formulas: P2, C2O4,
C6H14S2, H2, C3H9
• One can calculate the
empirical formula from the
masses or percentage
composition of any compound.
• Example: Calculate the
empirical formula for a
compound that has 43.7 g P
(phosphorus) and 56.3 grams
of oxygen.
• First we convert to moles:
• Next we divide the moles to try
to get an even ratio.
• When we divide, we did not
get whole numbers so we must
multiply by two (2).
• The answer=P2O5
• After one has calculated
the empirical formula one
can calculate the
molecular formula. By
dividing the molecular
mass of the compound by
the mass of the empirical
formula, one would attain
the molecular formula
• Example: The empirical formula of a compound
is HCN, 2.016 grams of hydrogen are necessary
to make the compound, what is the molecular
In the empirical formula hydrogen weighs
1.008 grams. Dividing 2.016 by 1.008 we see
that the amount of hydrogen needed is twice as
much. Therefore the empirical formula needs to
be increased by a factor of two (2). The answer
Concentrations of
• The concentration of a
solution is typically given
in molarity. Molarity is
defined as the number of
moles of solute (what is
actually dissolved in the
solution) divided by the
liters of solution (the total
volume of what is
dissolved and what it has
been dissolved in).
• Molarity = moles of solute
liters of solution
• Example: If 5.00 grams of NaOH are
dissolved in 5000 mL of water, what is the
molarity of the solution?
• the molarity (M) of the solution is 0.025
• Molality is another
common measurement of
concentration. Molality is
defined as moles of
solute divided by
kilograms of solvent (the
substance in which it is
dissolved, like water).
• It is possible to convert
between molarity and
molality. The only
information needed is
Chemical Reactions
• one of the best solvents is
water, because it is a polar
• Electrolytes are substances
whose aqueous solutions
conduct electricity
• When an ionic substance is
broken apart into the solvated
cations and anions, the
process is properly called
dissociation. However, when
the ions are formed by the
reaction of the substance with
water, and didn't exist in the
original substance, then the
process is properly called
• Some examples of substances which are
classified as strong electrolytes are:
• * strong inorganic acids (the common ones are:
hydrochloric acid, hydrobromic acid, hydroiodic
acid, sulfuric acid, nitric acid, chloric acid, and
perchloric acid
• * strong inorganic bases, e.g. NaOH.
• Dissociation reaction: NaCl + H2O ---> Na+ (aq)
+ Cl- (aq)
Ionization Reaction: HCl(g) + H2O(l) --->
H3O+ (aq) + Cl- (aq)
• Some examples substances
which behave as weak
electrolytes are:
• * Weak inorganic acids
(inorganic acids which are not
listed above), e.g. carbonic
acid, nitrous acid,
hypochlorous acid, etc.
• * Organic acids (carboxylic
acids, e.g. acetic acid)
• * Organic amines, e.g.
equations for chemical
• Molarity equation
(Molarity) x (Volume [in liters]) = moles of
• Dilution equation
Oxidation states
• Oxidation is the loss of electrons and reduction
is the gain of electrons
• The method used for keeping track of the
electrons is the oxidation number.
• The only time that an oxidation number is a real
charge is for monatomic ions, like Cl- (oxidation
number -1) or aluminum ion, Al3+ (oxidation
number +3). In hydrogen chloride, the oxidation
numbers of +1 for the hydrogen and -1 for the
chlorine are clearly not real; HCl is a covalent
compound which doesn't contain ions.
• Chapter five contained
many significant topics in
the study of chemistry.
• It is important to
remember that, starting
from the simplest form of
a concept, such as the
mole, and than working
toward a more elaborate
concept, such as
stoichiometry, is the best
way to master the topics.