Download isotopes - rmansell

Recall from grade 10:
# of protons = # of electrons = atomic #
# of neutrons = atomic mass – atomic #
Example: carbon: atomic = # 6
atomic mass = 12.01
# of protons = 6
# of electrons =
6
# of neutrons = 12- 6 = 6 *always round off
Different variations of atoms of the same
element occur in nature. These variations are
called isotopes. The average mass of the
isotopes for each element is a characteristic of
that element.
Isotopes are atoms of the same element (same #
of protons) with different # of neutrons. They
have identical atomic #’s but different mass #’s
(# of protons and neutrons).
A = mass #
X = symbol
Z = atomic #
Isotopes are usually represented in several ways.
24Na
sodium – 24 or
The atomic mass unit is defined as 1/12th the
mass of a carbon-12 atom. The magnitude of
the amu is arbitrary. In fact, 1/24th the mass of a
carbon atom or 1/10th the mass of the iron atom
could have been selected just as easily.
3 Reasons for using the 1/12th the mass of a C-12 isotope:
Carbon is a very common element
2. Results in nearly whole-number atomic masses for
most elements.
3. The lightest, H2, has a mass of ~ 1 amu.
1.
Calculating Average Atomic Mass
Example:
Isotope
Cl- 35
Cl – 37
Relavtive Abundance
75%
25%
total: 26.25 + 9.25 = 35.50
Mass
35x.75 = 26.25
37 x .25 = 9.25