300 Chemistry
Atomic Structure Notes
Key questions:
What is matter made up of?
How tiny are the particles?
How do you discover and describe tiny particles that you can’t see?
All matter is made up of ATOMS
Atom = smallest particle of an element that maintains the properties of that element in a
chemical reaction
Theories about how to describe atoms have evolved over time… here are some of the
important ideas
Scientist(s)
Democritus
(Greek
philosopher)
Dalton
(English
chemist and
teacher)
Time
frame
~400
B.C.
Key Discoveries or Ideas (items in bold are still known to be
true today)
Atoms are indivisible and indestructible
~1800
Dalton’s Atomic Theory (from experiments & scientific
method):
1. Elements are made up of tiny, indivisible atoms
2. All atoms of the same element are identical
3. All atoms of one element are different from those of
any other element
4. Atoms can mix together physically or chemically
combine in simple, whole number ratios to form
compounds
5. Chemical reactions occur when atoms are separated,
joined, or rearranged
6. Atoms of one element are never changed into atoms
of another element in a chemical reaction
 Did experiments with cathode rays (colored rays of
light) and discovered that they were made of electrons
(negatively charged particle in the atom) (rays would
bend toward a positive plate)
 Described atoms using the “plum pudding model”
(positive atom with negative electrons sprinkled
throughout)
 Did famous “gold foil experiment”
 Atom is mostly empty space
 Atom has a small, dense region in the center
(nucleus)
 Nucleus contains protons
J.J. Thomson
(English
physicist)
~late
1800s
Ernest
Rutherford
(Thomson’s
student)
~1911
Niels Bohr
(Danish
physicist and
Rutherford’s
student)
Various
scientists
ideas
combined
~1913




Electrons surround nucleus and are spread out
No explanation of chemical properties of the elements
Electrons are found in specific paths (circular orbits)
around the nucleus
They have fixed energy (energy levels, like rungs on a
ladder)
~1920s Quantum mechanical model:
 Electrons are not located in specific, fixed, circular
orbits
 Electrons have certain allowed energies and one can
determine how likely it is to find an electron with
that particular amount of energy
 Looks at probabilities of locating an electron
For now- we will focus on the basic structure of the atom… then focus on the nucleus
and later in the year, electrons and the quantum mechanical model
Basic Structure of the Atom
In the nucleus: protons and neutrons
Outside the nucleus: electrons
Protons = positively charged (+1 relative charge), mass of 1 atomic mass unit (1 amu)
Neutrons = neutral (no charge), mass of 1 atomic mass unit (1 amu)
Electrons = negatively charged (-1 relative charge), negligible mass (very tiny)
How do we determine how many protons, neutrons, and electrons an atom has?
Atomic number = Whole number on the Periodic Table, indicates the # of protons
In a NEUTRAL atom (atom with no charge), the # of protons must = the # of
electrons…. So for neutral atoms, atomic number = # of electrons
Atomic mass or average atomic mass = Decimal number on the Periodic Table
Mass number = Atomic mass from Periodic Table rounded to a whole number
Since only protons and neutrons have significant mass…
Mass number = # of protons + # of neutrons
So, to summarize, for a neutral atom:
# of protons = # of electrons = atomic number
# of neutrons = mass number - # of protons = mass number – atomic number
Variations on Atoms
2 types of variations can occur: ions and isotopes
Note: With these variations, the number of protons NEVER changes- it gives the
element its identity
Ion = an atom with the usual number of protons and a different number of electrons,
resulting in a net charge
Isotope = an atom with the usual number of protons and a different number of neutrons,
resulting in a different mass
Ions
Ions are formed when an atom loses or gains electrons
If an atom loses 1 electron, it has a +1 charge; loses 2 electrons, a +2 charge, etc.
If an atom gains 1 electron, it has a -1 charge; gains 2 electrons, a -2 charge, etc.
(graphs to show why charge is +/-)
So, you will have to use the charge to determine the number of electrons (protons
and neutrons will remain THE SAME)
Ex: Na1+ has a +1 charge and has lost 1 electron (so it has 11 protons, 12
neutrons, and 10 electrons)
Ex: N3- has a -3 charge and has gained 3 electrons (so it has 7 protons, 7 neutrons,
and 10 electrons)
Cation = positively charged ion (formed by losing electrons)
Anion = negatively charged ion (formed by gaining electrons)
Isotopes
Isotopes exist when an atom has a different number of neutrons and its mass changes
The number of protons and electrons will remain THE SAME as always
You might see isotopes written like this:
23
Na which means a sodium isotope with a mass number of 23, since Na always has 11
protons, it must have 23-11 or 12 neutrons
24
Na would have 11 protons, 11 electrons, and 24-11 or 13 neutrons
22
Na would have 11 protons, 11 electrons, and 22-11 or 11 neutrons
Note: Sometimes, they will include the atomic number underneath the mass number…
that just saves you the time of having to look it up on the Periodic Table
Isotopes and Average Atomic Mass
Masses on the Periodic Table are a WEIGHTED AVERAGE of all isotopes found in
nature, based on their percent abundance (how much they occur in nature)
Average atomic mass = weighted average of the mass of isotopes in a naturally occurring
sample of an element
Example:
Zinc has 5 isotopes that occur in nature as shown in the table…. Find its average atomic
mass.
Isotope
Mass number
Zinc-64
Zinc-66
Zinc-67
Zinc-68
Zinc-70
64
66
67
68
70
Mass of isotope
(amu) {Sometimes
you won’t have this
level of detail and
will use the mass
number}
63.929
65.927
66.927
67.925
69.925
Percent abundance
(%)
48.89
27.81
4.11
18.57
0.62
Average atomic mass of Zn = 0.4889(63.929) + 0.2781(65.927) + 0.0411(66.927) +
0.1857(67.925) + 0.0062(69.925) = 65.39 = mass on the Periodic Table! Voila!