Quantitative Chemistry
Chapter 3
Objectives
 Learning objective 1.2 The student is able to select and apply
mathematical routines to mass data to identify or infer the composition of
pure substances and/or mixtures.
 Learning objective 1.3 The student is able to select and apply
mathematical relationships to mass data in order to justify a claim
regarding the identity and/or estimated purity of a substance.
 Learning objective 1.4 The student is able to connect the number of
particles, moles, mass, and volume of substances to one another, both
qualitatively and quantitatively.
 Learning objective 1.14 The student is able to use data from mass
spectrometry to identify the elements and the masses of individual atoms
of a specific element
Objectives, Part 2
 Learning objective 3.1 Students can translate among macroscopic
observations of change, chemical equations, and particle views.
 Learning objective 3.3 The student is able to use stoichiometric
calculations to predict the results of performing a reaction in the laboratory
and/or to analyze deviations from the expected results.
 Learning objective 3.4 The student is able to relate quantities (measured
mass of substances, volumes of solutions, or volumes and pressures of
gases) to identify stoichiometric relationships for a reaction, including
situations involving limiting reactants and situations in which the reaction
has not gone to completion.
 Learning objective 3.6 The student is able to use data from synthesis or
decomposition of a compound to confirm the conservation of matter and
the law of definite proportions.
Atomic Mass
 Atoms are so small, it is difficult to discuss how much
they weigh in grams.
 Use atomic mass units.
 An atomic mass unit (amu) is 1/12 the mass of a carbon12 atom. (Adopted in 1961)
 The decimal numbers on the table are atomic masses
in amu.
 Sometimes abbreviated as u.
Decimals??
 Because they are based on averages of atoms and of
isotopes.
 Can figure out the average atomic mass from the mass
of the isotopes and their relative abundance.
 Add up the percent as decimals times the masses of
the isotopes.
Mass Spectrometer
Mass Spec Data
The Mole
 The mole is a number.
 A very large number, but still, just a number.
 6.022 x 1023 of anything is a mole
 The number of atoms in exactly 12 grams of carbon-12.
Molar Mass
 Mass of 1 mole of a substance.
 Often called molecular weight.
 To determine the molar mass of an element, look on
the table.
 To determine the molar mass of a compound, add up
the molar masses of the elements that make it up.
Find the Molar Mass
Why not both?
Percent Composition
 Percent of each element a compound is composed of.
 Find the mass of each element, divide by the total mass,
multiply by a 100.
Find the % Comp.
Working Backwards…
 From percent composition, you can determine the
empirical formula.
 Empirical Formula - the lowest ratio of atoms in a
molecule.
 Based on mole ratios of the constituent elements.
 A sample is 59.53% C, 5.38%H, 10.68%N, and
24.40%O, what is its empirical formula?
Pure O2 in
CO2 is absorbed
Sample is burned
completely to form
CO2 and H2O
H2O is absorbed
Try This!
 A 0.2000 gram sample of a compound (vitamin C)
composed of only C, H, and O is burned completely
with excess O2 . 0.2998 g of CO2 and 0.0819 g of
H2O are produced. What is the empirical formula?
Empirical To Molecular
Formulas
 Empirical is lowest ratio.
 Molecular is actual molecule.
 Ratio of empirical to molar mass will tell you the
molecular formula.
 Must be a whole number because...
Example
 A compound is made of only sulfur and oxygen. It is
69.6% S by mass. Its molar mass is 184 g/mol. What is
its formula?