Notes – Isotopes and Average Atomic Mass

Notes – Isotopes and Average Atomic Mass
How do you weigh an atom?
Atoms have masses far too small to be
weighed, so we use the ______________________ to
express the mass of atoms in
___________________________________ (________)
Atomic Mass is RELATIVE
 Atomic masses are
___________________________ where the
mass of one atom is compared to
 It is NOT an absolute scale of measure
like grams.
Think “Weighted Average”
 The Mass listed in the periodic table
for each element is the
________________________ of all
________________________ for that element.
 Remember: ______________________ of the
same element have _____________________
because they have
Example, PLEASE!!
If you analyzed a lump of pure carbon from
the planet earth, you would find that 98.9%
of all carbon atoms on earth are C-12 atoms
and 1.11% of all carbons atoms on earth are
C-13 atoms.
Look at the PT Mass Again!
What is AMU?
 Atomic Mass Unit
_______________________________and has the
symbol u or amu.
 1 amu is _____________________ to the
mass of a single _________ or ____________.
It’s All About Carbon
 The ____________________ is the basis for
the relationship between all other
 The mass of a carbon atom was
divided into 12 parts (6 protons, 6
neutrons) and the mass of each part is
known as an _____________________ (_____).
 1 amu = _________________ the mass of a
Carbon-12 atom
So what’s “12.011” mean
 We define the atomic mass of C-12 as
exactly ____________________.
…but that is NOT the number you see
in the Periodic Table!
Which Isotope counts the most?
 The mass in the Periodic Table will be
closest to that of the
____________________________________ of that
 The weighted average factors in the
_________________________ of each isotope
as it occurs in nature.
One more example…
The natural abundance of Cu-63 is 69.09%
and for Cu-65 is 30.19%.
If the atomic weight of Cu-63 is 62.93 amu
and Cu-65 64.93 amu, what is the average
atomic weight for natural copper?