Stoichiometry
Dalton’s Model (1803)
• All matter is ultimately composed of atoms.
• Atoms can be neither subdivided nor changed
into one another.
• Atoms can be neither created or destroyed.
• Atoms of a particular substance are identical.
• Chemical change is the union or separation of
atoms.
Mole concept
• Basis of stoichiometry
• A mole is a count – 6.02 x 1023 particles
(Avogadro’s number)
• Why this strange number? If you take the
molar mass of any compound in grams,
this quantity will contain Avogadro’s
number of particles, 1 mole.
Formulae
• Show the number of atoms of each
element in a covalent molecule or relative
number of units in an ionic compound.
• Take the molar mass in grams. You get
one mole of molecules.
• Chemical substances react by count, not
mass. However, in lab, we measure
quantities by mass. The mole concept
gives us a way of relating count to mass.
Example
• Which of the following weighs most:
50 g of iron, 5 moles of nitrogen gas, 0.10
mole of silver, 1 x 1023 atoms of radium ?
Example
• A compound was analyzed to give the
following percentage composition:
potassium (26.57%), chromium (35.36%),
and oxygen (38.07%). What is the formula
for the compound?
Example
• Compound Q contains 3 atoms of L and 2
atoms of M for every atom of X. In a given
reaction starting with 3.6 x 1022 atoms of X
and 0.12 mole of L, 3.6 g of M is required.
What is the atomic mass of M?
Chemical Equations
• Atoms are neither destroyed or created in
a chemical reaction. We need the same
number of atoms of each element on both
sides of the equation.
• Chemical equations show counts of atoms
involved, not masses. Need to convert
moles to masses and vice versa.
Types of Chemical Reactions
Combustion
C2H6O + 3 O2  2 CO2 + 3 H2O
Replacement (displacement)
Zn + CuSO4  Cu + ZnSO4
Double displacement
AgNO3 + NaCl  NaNO3 + AgCl (s)
Types of Chemical Reactions
Acid-base neutralization
HCl + NaOH  NaCl + H2O
Combination (single product)
2 SO2 + O2  2 SO3
Decomposition
2 HgO  2 Hg + O2
Example
Balance FeS2 + O2  Fe2O3 + SO2
Example
Balance C7H6O2 + O2  CO2 + H2O
example
Balance
Ba(NO3)2 + Na3PO4  Ba3(PO4)2 + NaNO3
Example
• Before the public became serious about
pollution, it was common to improve the
performance of gasoline by the addition of lead
compounds. A particular 100-octane aviation
gasoline used 1.00 cm3 of tetraethyl lead,
(C2H5)4Pb, per liter of product. The pure
tetraethyl lead is made as follows:
4 C2H5Cl + 4 NaPb  (C2H5)4Pb + 4 NaCl + 3 Pb
How many grams of ethyl chloride, C2H5Cl, are
used to make enough tetraethyl lead for 1 liter of
gasoline?
Example
• A solution containing 2.00 g of mercury (II)
nitrate was added to a solution containing
2.00 g of sodium sulfide. Calculate the
mass of mercury (II) sulfide that was
formed.
Reactions and Energy
• Reactions will either produce energy or
require energy.
• Exothermic – reactions liberate energy.
Products have lower energy content than
reactants
• Endothermic – reactions consume energy.
Products have higher energy content than
reactants.
Activation Energy
• Energy in endothermic reactions is NOT
the same as activation energy. Activation
energy relates only to the rate of reactions.
Catalysts provide an alternative reaction
path to avoid need for high activation
energy. Even though activation energy
may well be required, a reaction can still
be very exothermic.