PPSlidesStoich - Barrington Public Schools

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
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
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
• Sample Exercise 3.5 page 80
Percentage Composition from
• 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
Avagadro’s number.
• Thus 1 mole of carbon atoms =
6.0221421 x 1023 carbon atoms
Molar Mass
• mass in grams of 1 mole of a
• 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
• 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 &
• Can be used to determine
empirical formula or percent
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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