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									          The Mole
            Kelley Kuhn
Center for Creative Arts
The Mole!

When you look at the periodic table, the
 average atomic mass below each element
 represents the average mass of one atom
 in amu. (Review: what is an amu?) This
 is an incredibly small mass and not really
 reasonable to work with in a lab setting.
 SO, if we take the average atomic mass
 and treat it as a quantity of grams, we call
 this quantity a mole.
Someone else explaining the mole

The Mole!
For example:
Look at carbon on the periodic table:
1 atom of carbon has the average mass of 12.01
1 mole of carbon has the mass of 12.01 grams.

 Question:
  How many grams in one mole of lead?
  How many grams in 4.5 moles of lead?
A few more details about the mole

***A guy named Avogadro calculated how
 many atoms would be in one mole of an
 element and determined that it would
 equal 6.022 x 1023 atoms of that element.
 In this sense, we use the term mole in the
 same way we use the term "1 dozen". A
 dozen always equals 12 of something; a
 mole always represents 6.022 x 1023 of
Moles and compounds
 ***The mole is also used to describe quantities of
  compounds and molecules. The "molar mass" (or grams
  in a mole) of any substance can be determined by
  adding up the individual masses of the atoms in the
 For example:
 Water, H2O, contains two hydrogen atoms and one
  oxygen atom.
 H = 1.01 gram per mole
 H = 1.01 gram per mole
 O= 16.00 gram per mole
 H2O = 18.02 grams per mole
How many grams would be in 1 mole of
How many grams would be in 1 mole of
 ammonia, NH3?

This quantity (grams in 1 mole) is called:
molar mass or molecular mass or
 molecular weight
Percent Composition

 Percent composition is a way to describe what a
  substance is made of by listing the percent of each
  element in the compound.
 Percent = part/whole x 100
 To determine the percent composition of one or all
  elements in a compound:
 1. Determine the molar mass of the compound by
  adding up the individual masses of all the atoms in the
 2. Divide the total individual mass of each element by
  the overall molar mass and multiply by 100.
Percent Composition example:
   Example:
   H3PO4
   3 H = 3(1.01) = 3.03
   P = 30.97       = 30.97
   4 O = 4 (16.00) = 64.00
                     98.00
   H    3.03/98.00 x 100 =              %H
   P    30.97/98.00 x 100 =             %P
   O 64.00/98.00 x 100 =                %O
   Practice: what is the percent composition of glucose,
Empirical Formula

 So, we have gone from the chemical formula to
  the percent composition. Now, we learn to go
  from percent composition back to the chemical
  formula, in a way. From these percents, we can
  determine the empirical formula, which is the
  simplest whole number ratio of atoms in a
  molecule. The empirical formula may or may not
  be the same as the molecular formula which is
  the actual ratio of atoms in a molecule.
Empirical formulas cont.

Molecular formula of glucose: C6H12O6
Empirical formula of glucose: CH2O
Determining empirical formula:
1. Treat the percentages as grams.
2. Convert the grams to moles.
3. Divide each mole quantity by the smallest mole
  quantity to achieve an initial ratio.
4. If this ratio is extremely close to whole numbers,
  STOP. This is your ratio for your empirical
5. If this ratio is not close to whole numbers,
  multiply by a whole number to achieve a whole
  number ratio. (For example, if the ratio is 1:1.5,
  you will need to multiply by 2 to get to 2:3.)

Example: A compound is made of 23.29%
 Mg, 30.72% S, and 45.99% O. What is its
 empirical formula?

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