Structure of an Atom
• Atoms are composed of 3 subatomic particles:
1. Protons (p+) – Positively charged, found in the
nucleus, number of protons determines the element ,
has a mass of 1 amu
2. Neutrons (n) – Neutral particle, found in the nucleus,
can vary in number, has a mass of 1 amu
3. Electrons (e-) – Negatively charged, orbit around the
nucleus, the farther from the nucleus the higher the
energy level (relate to periods on PT), mass is so small
relative to protons and neutrons we say it has a mass
of 0 amu (it is actually 1/1836 that of a proton)
Charges and the Atom
• A positive and a negative charge cancel each
other out…..
When you have an equal
number of each you have
Protons are positive
a neutral atom, if they are
Electrons are negative
different then there is a
charge
If an atom becomes charged we call it an ion
If an atom has more protons (positives)than
electrons (negatives) it will be positive…. These
are known as CATIONS, usually metals
Ca ion = positive ions
If an atom has more electrons (negatives) than it
has protons (positives) it will be negative…..
These are known as Anions, usually nonmetals
nion = negative ions
Elements and the periodic table
How do we read info from the periodic table?
Atomic number
(equal to the # of protons)
7
N
Element symbol
Abbreviation of element name
Nitrogen
14.0067
Element name
Name of the element
Average atomic mass
Weighted Average mass of 1 mole of all isotopes
based on relative abundance
When things CHANGE in an atom. WHAT HAPPENS?
• If the electrons change  the atom has a positive
or negative charge. AN ION IS FORMED.
• If the neutrons change  the mass of the atom
changes. Version of the same atom. AN ISOTOPE IS
FORMED.
Isotopes:
Isotope Notations
Hyphen Notation
Mass Number
Helium - 3
Name or
Abbreviation
Nuclear Notation
Mass Number
3
He
2
Abbreviation
Atomic Number
Mass Number - vs - Average Atomic Mass
Average Atomic Mass:
Mass Number:
1. Protons + Neutrons 1. Average mass of all
the isotopes
= Mass #
a. Based on relative
2. Is a whole number
abundance
3. Used when talking 2. Has numbers after a
about isotopes
decimal
Using Isotope Notations
• We will use isotope notations to keep track of how many protons
and neutrons we have
• To fill in the table below you will need to use these tricks:
Mass # = Protons + Neutrons
Neutral atoms have equal numbers of protons and
electrons
Hyphen Atomic
Notation Number
Mass
Number
Number
of
protons
Number of
neutrons
N-14
28
80
45
Number of Nuclear
electrons Notation
Relative Abundance
Slide 12
• Relative abundance – refers to the abundance of
naturally occurring isotopes
– An example is Chlorine (Cl), which has 2 isotopes:
1) Chlorine – 35
2) Chlorine – 37
relative abundance = 75.7%
relative abundance = 24.3%
• Calculating the weighted Average atomic mass:
– Multiply the mass # by the relative abundance for
each isotopes, then add them all together
35 x 0.757 = 26.5
Multiply % by mass #
37 x 0.243 = 9.00
Add them all together
Weighted Avg : 26.5 + 9.00 = 35.5 amu
Relative Abundance practice
Slide 13
1. Gallium consists of two naturally occurring isotopes with
masses of 68.926 amu (60.1% abundant) and 70.925 amu
(39.9% abundant). What is the weighted Average atomic
mass?
2. Strontium consists of four isotopes with masses of 84 amu
(abundance 0.50%), 86 amu (abundance of 9.9%), 87 amu
(abundance of 7.0%), and 88 amu (abundance of 82.6%).
Calculate the atomic mass of strontium.
3. Rubidium has two common isotopes, Rb-85 (abundance of
72.5%) and Rb-87 (abundance of 27.8%). What is the average
atomic mass?
4. Argon has 3 naturally occurring isotopes: argon-36, argon-38
and argon-40. Based on argon’s reported atomic mass, which
isotope do you think is the most abundant in nature? Explain.