Download 09/11/03 lecture

From “Microscopic” to
“Macroscopic”
Chem 101 Lecture
09/10/03
What we need to learn…
(the macroscopic world)
• How do we describe/characterize “bulk”
properties of different compounds? (i.e., large
collections of atoms and molecules).
• How do the atomic and molecular structures
related to physical and chemical properties of
“bulk” materials?
Atomic masses revisited...
• We’ve talked about atomic masses in terms of the number
of neutrons and protons present in an atom…but how
much does an atom weigh? What units do we describe the
mass of an atom in?
• The atomic mass unit (amu): defined explicitly in terms of
the 12C atom--the mass of 1 12C atom = 12 amu.
• All other atomic masses are defined relative to the 12C
atom.
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Return to isotopes...
• Some elements can have a number of different possible
atomic structures which differ only in the number of
neutrons in the nucleus--isotopes.
• For example, there are three different atomic structures for
C : 12C, 13C, and 14 C.
• When dealing with bulk quantities of carbon, more useful
to think of an average atomic weight which will take into
account the individual atomic masses and their relative
abundance.
Finding the average mass of carbon
• The relative abundances of the isotopes of carbon are :
–
–
12C
13C
: 98.89%
: 1.11%
• The average mass of carbon = (0.9889)*(12.00 amu) +
(0.0111)*(13.00 amu) = 12.01 amu
• This is the mass of carbon we use when calculating
amounts of the “bulk” carbon.
Going from amu to grams...
• Atomic mass units are still really small…how do we relate
amu to something we can measure easily in the lab (like
grams)?
• Introduce a new unit: the mole
• A mole is equal to the number of carbon atoms in exactly
12 grams of carbon. This number is called Avogadro’s
number and is equal to 6.022 x 10 23.
• In other words, 12 g of carbon contain 1 mole of carbon
atoms (or 6.022 x 10 23 carbon atoms).
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Random facts about moles
• A mole is analogous to a dozen…a unit for counting
things.
• The unit “mole” (abbreviated “mol”) should be
accompanied by the object being counted (just as we say
“one dozen eggs,” we should say “one mole of carbon
atoms”).
So, we can think of molar mass
as well as atomic mass
• Molar mass = mass of a mole of atoms (if its an element) or
molecules (if a compound) of a substance.
• For elements, molar mass is defined in such a way that the
molar mass is the average atomic weight expressed in grams:
Element
Molar mass
Aluminum
Atomic
weight
26.98 amu
Carbon
12.01 amu
12.01 g
26.98 g
Hydrogen
1.01 amu
1.01 g
Potassium
39.10 amu
39.10 g
The molecular weight of a molecule is
the sum of the atomic weight of its
constituent atoms
• What is the molecular weight of methanol CH4O ?
–
–
–
–
Mass of 1 C = 12.01 amu
Mass of 4 H = 4 x 1.008 amu
Mass of 1 O = 16.00 amu
Total mass of CH 4O : 32.04 amu
• What is the molar mass of methanol?
– Molar mass (mass of 1 mole) of CH 4O = 32.04 g/mol
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Key skills
• Convert from # grams to # mol and convert
from # mol to # grams.
• From a molecular formula, calculate mass
percent composition.
• From a mass percent composition, calculate
empirical formula.
• From empirical formula and molar mass,
calculate molecular formula.
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