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Name: ____________ Date: ______________
Unit 3 Notes Activity #5: Quantitative Chemistry, Counting Atoms and Molecules
Scaling up from Microscopic to Macroscopic Measurements!
Earlier in the year we discussed the relative masses of subatomic particles in atomic mass units (amu), the
existence of isotopes, and how the atomic weight of an element is a weighted average of the masses of each
atom’s isotopes.
To bridge the gap between the masses of individual atoms, to the sum of the masses of many atoms
measured in the laboratory complete chemical reactions, chemists developed the concept of the mole, which
is just a number that represents a specific number of atoms of an element or molecules of a compound.
One dozen atoms = _________________
One pair of atoms = _________________
Neither of the quantities above could be counted, let alone their masses measured on a laboratory scale!!
One mole of atoms = ___________________ (in scientific notation), the mass of this quantity of atoms can be
measured on a laboratory scale!
The Molar Mass of Elements:
Microscopic Scale
Macroscopic Scale
Atomic Weight of an Atom in amu
Mass of 1 mole of Atoms in grams
1 Hydrogen atom
1 mole of Hydrogen atoms
1 Carbon atom
1 mole of Carbon atoms
1 O atom
1 mole of O atoms
1 Al atom
1 mole of Al atoms
1 Br atom
1 mole of Br atoms
Mole: The number of atoms in 12.00 grams of the Carbon-12 isotope (6 p+, 6 n0, 6e-) is:
1 mole = 12.00 g = 6.02 x 1023 atoms.
1.992 x 10-23 g/atom
1 mole of atoms (or molecules, or ions) = _________________ atoms (or molecules, or ions)
So, the mass in grams equivalent to the atomic weight of any element on the periodic table provides a means
to measure 1 mole of atoms of the element, or 6.02 x 1023 atoms.
This number, 6.02 x 1023, is called Avogadro’s number, and the mass of this number of atoms of an element
(or molecules of a compound) is also called the Molar Mass of the element!
Number of Moles
Molar Mass
(g/mol)
Measured Mass
on lab scale
Number of atoms
1 mole of Hydrogen atoms
1 mole of Carbon atoms
