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Ch. 8 Formula
Stoichiometry
Stoichiometry – studies the
mathematical relationships involving
chemical formulas and equations.
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The mass of an atom is very very small:
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oxygen atom 2.66 x 10-23 grams
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hydrogen atom 1.67 x 10-24 grams
Atomic Mass
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Although it is possible to
measure the actual mass
of an atom, it is more
convenient to use relative masses.
The relative mass of an atom is a
comparison of the atom’s mass to the mass
of a carbon-12 atom.
Carbon-12 was chosen as the standard for
measuring relative atomic masses.
Atomic Mass Unit
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The mass of a carbon-12 atom = exactly 12 u
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One atomic mass unit (amu or u) =
1/12 the mass of a carbon-12 atom.
Atomic masses for each
element are listed on the
periodic table.
Some atomic masses:
• Hydrogen 1.008 u
• Oxygen 16.00 u
• Iron
55.85 u
What is the atomic mass of:
Nitrogen
14.01 u
Chlorine
35.45 u
Magnesium
24.31 u
Formula Mass
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The sum of the atomic masses
of all the atoms in a chemical
formula.
Formula mass can refer to an
element or an ionic or
molecular compound.
EX: Formula mass of CO2 is:
12.01 + 2 x 16.00 = 44.01 u
SUBSTANCE FORMULA MASS
Fe
55.85 u
H2
2 x 1.008 = 2.016 u
H2O
2 x 1.008 + 16.00 = 18.02 u
What is the formula mass of Mg(NO2)2?
• 1 Mg + 2 N + 4 O
• 24.31 u + (2 x 14.01 u) + (4 x 16.00 u) = 116.3 u
•Calculate the formula mass of:
1) NaCl 2) NO2 3) CaCO3
4) Ba(NO3)2 5) (NH4)3PO4
Molecular Mass
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molecular mass is the formula mass of a
molecular substance. It is the mass of a
molecule of that substance.
EX: the molecular mass of C6H12O6 =
6 X 12.01
+ 12 X 1.008
+ 6 X 16.00
= 180.2 u
The Mole: Basic Concepts
Measuring Matter
• Chemists need a convenient method
for counting accurately the number of
atoms, molecules, or formula units in
a sample of a substance.
• As you know, atoms and molecules are
extremely small. There are so many of
them in even the smallest sample that
it’s impossible to actually count them.
• That’s why chemists created their own
counting unit called the mole.
• The mole, commonly abbreviated mol, is the
SI base unit used to measure the amount of a
substance.
• 1 mole (mol) = 6.02 x 1023 particles
• “particles” may be atoms, molecules, or formula
units
The number 6.02 x 1023 is called
Avogadro’s number (in honor of
Italian Amedeo Avogadro who
determined the volume of one
mole of a gas.)
• If you write out Avogadro’s number, it looks
like this: 602 000 000 000 000 000 000 000
• A representative particle is any kind of particle
such as atoms, molecules, formula units,
electrons, or ions.
• We will now look at one-mole quantities of
three substances, each with a different
representative particle.
• The representative particle
in a mole of water is the
water molecule.
• 1 mol H2O =
6.02 x 1023 H2O molecules
• The representative
particle in a mole of
copper is the copper
atom.
• 1 mol Cu =
6.02 x 1023 Cu atoms
• The representative
particle in a mole of
sodium chloride is the
formula unit.
• 1 mol NaCl =
6.02 x 1023 NaCl
formula units
• The term “formula unit”
is used to represent one
“unit” of an ionic
compound.
Mole  Particle Conversions
Avagadro's number is used as a conversion
factor to convert between # of moles and
# of particles (atoms, molecules, ions, etc.)
Converting Moles to Particles
• Suppose you want to determine how
many particles of sucrose are in 3.50
moles of sucrose. You know that one
mole contains 6.02 x 1023
representative particles.
• Therefore, you can write a conversion
factor, Avogadro’s number, that relates
representative particles to moles of a
substance.
• There are 2.11 x 1024 molecules of
sucrose in 3.50 moles of sucrose.
EX: How many iron atoms are in 0.025 mol of Fe?
Converting Particles to Moles
• Now, suppose you want to find out how many
moles are represented by a certain number of
representative particles.
• You can use the inverse of Avogadro’s
number as a conversion factor.
• Zinc is used as a corrosionresistant coating on iron
and steel. It is also an
essential trace
element in your diet.
• Calculate the number of moles that
contain 4.50 x 1024 atoms of zinc (Zn).
The Mass of a Mole
• Technically, the mole is defined as the
number of carbon-12 atoms in exactly 12 g of
pure carbon-12. (This has experimentally
been shown to be 6.02 x 1023).
Thus, the mass of one mole of carbon-12
atoms is 12 g.
What about other elements?
(Play Video)
(Whether you are considering a single atom or
Avogadro’s number of atoms (a mole), the
masses of all atoms are established relative to
the mass of carbon-12. )
Mass of one mole of: Sulfur 32.06 g
Aluminum 26.98 g
H2O 18.02 g
Sodium 22.99 g
O2 32.00 g
CO 28.01 g
Molar mass – the mass of one mole of a
substance. The molar mass of a substance
is equal to its formula mass expressed
in grams. (units for molar mass are g/mol)
(molar mass is sometimes called “gram formula mass”
or “gram molecular weight”, etc.)
What is the molar mass of:
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iron?
phosphorus?
nitrogen gas (N2)?
sodium chloride?
sulfur dioxide?
Remember: Molar masses are the same as
formula mass, except the units are grams; you
get molar masses from the periodic table.
Mass(grams)   Moles Conversions
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Molar mass is used as a conversion factor
to convert between grams (mass) of a
substance and # of moles.
Converting Moles to Mass
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to convert moles to grams, use the following
set up:
molar mass
given moles X
1 mol
= grams
EX: Calculate the mass of 0.625 moles of
calcium.
Converting Mass to Moles
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to convert grams to moles, use the following
set up:
1 mol
given grams X molar mass = moles
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EX: Convert 848 g of copper to moles.
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Two-Step Problems
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converting: # of particles to grams;
or grams to # of particles.
1 mol
X molar mass
GRAMS
(mass)
molar mass
X
1 mol
MOLES
6.02 x 1023
X
1 mol
PARTICLES
1 mol
X 6.02 x 1023
PerkinElmer 2400 Series II CHNS/O
Analyzer
Percentage Composition
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the percentage composition of a compound is
the percentage (by mass) of each element in
that compound.
To calculate the percentage of an element in a
compound from the compound’s formula:
divide the total atomic masses of all the atoms
of that element in the formula BY the formula
mass of the formula (then multiply by 100 to
get a percentage).
Another type of % composition
problem:
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Sometimes they just tell you how many grams of
each element are in a sample of a compound.
To calculate the percentage of an element in a
compound from actual masses:
divide the mass of that element by the total
mass of the compound (then multiply by 100 to
get a percentage).
Determining Empirical Formulas
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In one type of chemical analysis, a compound is
decomposed and the masses of the elements
that made up the compound are measured. We
then use the masses of the elements in the
compound to calculate the empirical formula.
Steps for Determining Empirical
Formula:
1.
2.
3.
Change the mass of each element to moles.
(If you are given the percentage of each
element, just change the “%” sign to “g”,then
convert to moles.)
Divide each answer in step 1 by the smallest
answer to obtain the ratio of atoms in the
compound.
If step 2 does not give you the simplest
WHOLE number ratio, find a number that you
can multiply each answer in step 2 by in order
to achieve a whole number ratio.
Mass Spectrometer – determines molecular mass
Determining a molecular formula
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1.
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You can determine the molecular formula from an
empirical formula (the molecular mass must be
known).
The molecular formula is always a whole number
multiple of the empirical formula.
Steps for determining molecular formula:
Determine empirical formula.
Find formula mass of empirical formula.
Divide the given molecular mass by the
empirical formula mass.
Multiply each subscript in the empirical
formula by the answer in step 3 to give you
the molecular formula.