# PPSlidesStoich - Barrington Public Schools

```Stoichiometry
Quantitative nature of chemical formulas
and chemical reactions
Chapter 3
(Sections 3.3 - 3.7)
Atomic Mass Scale
• Atomic mass units (amu) are
convenient units to use when dealing
with extremely small masses of
individual atoms
• 1 amu = 1.66054 x 10-24 g
• 1 g = 6.02214 x 1023 amu
• By definition, the mass of C-12 is
exactly 12 amu
Average Atomic Mass
•
•
•
•
•
(Atomic Weight)
We average the masses of isotopes using their
masses and relative abundances to give the
average atomic mass of an element.
Naturally occurring C consists of 98.892%
C-12 (12.00 amu) and 1.108% C-13 (13.00335
amu)
The average mass of C is
(0.98892)(12.00 amu) + (0.01108)(13.00335)
= 12.01 amu
Atomic weights are listed on the periodic
table
Calculating % Abundance
• Chlorine is made up of two isotopes,
Cl-35 atomic mass = 34.969 amu and
Cl- 37 atomic mass = 36.966 amu.
Given chlorine’s atomic weight of
35.453, what is the % abundance of
each isotope?
34.969(x) + (36.966) (1-x) = 35.453
34.969x + 36.966 – 36.966x = 35.453
x = .7576
75.76% Cl-35 and 24.24% Cl-37
Formula & Molecular Weights
• Formula Weight is the sum of atomic
weights for the atoms present in the
chemical formula
• Molecular Weight is the sum of
atomic weights of the atoms in a
molecule as shown in the molecular
formula
• Sample Exercise 3.5 page 80
Percentage Composition from
Formulas
• Obtained by dividing the mass
contributed by each element
(number of atoms times atomic
weight) by the formula weight of
the compound and multiplying
by 100
• Sample Exercise 3.6 page 80
The Mass Spectrometer
page 81
• Mass spectrometers are pieces of equipment
designed to measure atomic and molecular
masses accurately.
• The sample is converted to positive ions by
collisions with a stream of high-energy
electrons upon entering the spectrometer.
• The charged sample is accelerated using an
applied voltage.
• The ions are then passed into an evacuated
tube through a magnetic field.
• The magnetic field causes the ions to be
deflected by different amounts depending on
their mass – more mass, less deflection.
• The ions are then detected.
The Mole
• The mole is a convenient
measure of chemical quantities.
• 1 mole of something is 6.0221421
x 1023 of that thing.
• This number is called
• Thus 1 mole of carbon atoms =
6.0221421 x 1023 carbon atoms
Molar Mass
• mass in grams of 1 mole of a
substance
• expressed in units of g/mol
• formula weights are numerically
equal to the molar mass
• Sample Exercise 3.8 page 84
Interconverting Masses, Moles,
& Number of Particles
• To convert between grams and
moles, we use the molar mass
• To convert between moles and
particles (atoms, molecules, or ions)
we use Avogadro’s number
• Sample Exercises 3.7, 3.9, 3.10, 3.11
pages 82-86
Empirical Formula
• Gives the relative number of
atoms of each element in the
substance
• Can be calculated from mass
percent data
• Sample Exercise 3.12 page 87
Molecular Formula
• Actual number of atoms of
each element in one molecule
of the substance
• Whole number multiple of
empirical formula
• Sample Exercise 3.13 page 88
Combustion Analysis
• A sample containing C, H,
and O is combusted in excess
oxygen to produce CO2 &
H2O
• Can be used to determine
empirical formula or percent
composition
Quantitative Information
from Balanced Equations
• Coefficients can be interpreted
as the relative numbers of
molecules or formula units in the
reaction as well as the relative
number of moles
• See Figure 3.14 page 92
• Sample Exercises 3.14, 13.15
pages 92 & 93
Limiting Reactant(s)
• The reactant(s) that is completely consumed in
a reaction
• Limits or determines the amount of product
that will be formed
• The other reactant(s) that is left over is called
the excess reactant
• Sample Exercise 3.16 and 3.17 pages 95-96
• How much of the excess reactant is left over?
Percent Yield
(actual yield / theoretical yield) X 100
• Actual is the amount of product
recovered in the lab
• Theoretical is the amount predicted
from stoichiometry
• Sample Exercise 3.18 page 97
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