Atom – nucleus (protons and neutrons)
electron cloud (electrons)
Atomic Number – equal to the number of protons
Mass Number – protons + neutrons
Charge – when # of electrons ≠ # of protons
Negatively Charged Ion (anion) – has more electrons
than protons
Positively Charged Ion (cation) – has more protons
than electrons
Neutral Atom - # of electrons = # of protons
All of this information may be written as part of a
chemical symbol:
mass number
15
7
atomic number
N 3
charge
Isotopes
So far, you know that the mass number is equal to
the number of neutrons plus the number of protons.
This number is important in our discussion of
elements (and their atoms), but why? Let’s
investigate…
In nature, elements are made up of atoms that are
not exactly the same…some are heavier than
others…but why??
Consider the relative masses of the sub-atomic
particles:
Protons – 1.00728 atomic mass units (amu)
Neutrons – 1.00867 amu
Electrons – 0.00055 amu
Since the number of protons NEVER changes, they
are not responsible for the changes in mass.
Since electrons weigh pretty much NOTHING
compared to the other particles, they are not
responsible for the changes in mass.
Therefore, it is the change in the number of
neutrons that is responsible for the changes in mass.
In nature, elements are made up of atoms that have
slightly different numbers of neutrons. This results
in the different “varieties” of atoms – called
isotopes. Thus, it is important to understand how to
interpret isotopic symbols; it allows us to figure out
WHICH “version” of the element (atom) we are
dealing with.
So…For an element, the atomic number (number of
protons) NEVER changes, but the mass number will
change depending on the isotope.
Some additional info…
These different versions of atoms are taken into
account when determining the atomic mass. On the
periodic table, the atomic mass is the average of all
the masses of the isotopes.
Example: C-12 and C-14
Lewis (Electron) Dot Diagrams
Atoms tend to gain or lose electrons in an attempt to
satisfy the octet rule. That is, to become like the
noble gases, which are the most stable elements.
Atoms can share or give/take electrons. Electrons
are being used to form bonds.
In order to draw atoms and their electrons in the
most concise and efficient manner, Lewis suggested
that we only include the valence electrons, since they
are the only ones that are involved in the bonding.
Examples:
Na
Be
Al
Sn
N
S
F
Ar
Give/Take Diagrams – used to show the transfer of
electrons (ionic bonding…)
Examples:
Na
Li
Cl
NaCl (sodium chloride)
O
Li2O (lithium oxide)
Li
Be
F
F
BeF2 (beryllium fluoride)
Question: What do you notice about how the
electrons move?
 Electrons are given from the metals
 Electrons are taken by the non-metals
 Electrons are always transferred from the atom
with the least electrons to the ones with more
(it is easier)
How to Draw Lewis Structures
Lewis structures are a way to write chemical compounds where all the atoms and
electrons are shown.
Steps to follow:
1. Count the total valence electrons for the molecule: To do this, find the
number of valence electrons for each atom in the molecule, and add them
up.
2. Figure out how many octet electrons the molecule should have, using the
octet rule: The octet rule tells us that all atoms want eight valence
electrons (except for hydrogen which only wants 2), so they can be like
the nearest inert gas. Use the octet rule to figure out how many
electrons each atom in the molecule should have, and add them up.
3. Subtract the valence electrons from octet electrons: Or, in other words,
subtract the number you found in step 1 from the number you found in
step 2 above. The answer you get will be equal to the number of bonding
electrons in the molecule.
4. Divide the number of bonding electrons by two: Remember, because every
bond has two electrons, the number of bonds in the molecule will be equal
to the number of bonding electrons divided by two.
5. Draw an arrangement of the atoms for the molecule that contains the
number of bonds you found in step 4 above: Some handy rules are:
- hydrogen always has one bond, never more.
- the halogens usually have one bond.
- the family that oxygen is in usually makes two bonds.
- the family that nitrogen is in usually makes three bonds.
- the family that carbon is in always makes four bonds.
6. Find the number of lone pair (nonbonding) electrons by subtracting the
bonding electrons (step 3) from the valence electrons (step 1). Arrange
these around the atoms until all of them satisfy the octet rule:
Remember, ALL elements EXCEPT hydrogen want eight electrons around
them, total. Hydrogen only wants two electrons.
CHEMICAL BONDING
- chemical bonds are attractive forces.
- valence electrons are the only electrons that take
part in chemical bonds. They are the outer
energy level electrons so they are available.
There are two different types of chemical bonds
that occur between atoms: ionic bonding and
covalent bonding.
Note:
i) chemical bonds involve a competition for
bonding electrons in unfilled orbital.
ii) bonds continue to form until each atom
has filled its valence orbital (octet rule).
IONIC BONDING
- ions that are formed when neutral atoms either
gain or lose valence electrons.
ion - an atom that has gained or lost electrons.
Cation  electron  neutral  electron  Anion
+
loss
atom
gain
-
- the driving force behind the formation of ions is
the tendency of atoms to reach a more stable
electron configuration (to have a stable octet).
This is done by acquiring either a completely full
or a completely empty valence level. This gives
the ion the same stable electron configuration as
a noble gas.
Example.
1. S + 2 e- = [ S ]2-
2. Na - 1 e- = [ Na ]+
- the ions with opposite charges will attract one
another by electrostatic forces. These forces
of attraction that bind oppositely charge ions
are called ionic bonds.
ATOMS
IONS
IONIC COMPOUND
1.
Na + Cl 
[ Na ]+ + [ Cl ]-
 NaCl
2.
Ca + 2 Cl 
[ Ca ]2+ + 2 [ Cl ]-
 CaCl2
3.
2 Al + 3 O  2 [ Al ]3+ + 3 [ O ]2-  Al2O3
Note:
1. The number of ions in the formula is adjusted
so that each compound is neutral overall.
2. Each element in the compound has achieved a
noble gas configuration.
Use electron dot diagrams to show the ionic bonding
process with these elements and write the ionic
equation to show the compound formed.
1. Potassium
+
Oxygen
2. Magnesium
+
Nitrogen
3. Aluminum
+
Bromine
4. Potassium
+
Iodine
5. Calcium
+
Sulfur
6. Sodium
+
Phosphorus
7. Aluminum
+
Sulfur
8. Sodium
+
Oxygen
COVALENT BONDING
- some atoms share electrons in order to attain
stable octets.
By sharing electrons, two fluorine atoms can
complete their octets and become more stable.
This process can be shown with electron dot
diagrams. When one pair of electrons is shared
between two atoms, it is called a single covalent
bond.
F+F= F F
- atoms can share more than one pair of electrons.
Double covalent bonds involve two shared pairs.
O+O= O O
- nitrogen atoms combine with a triple covalent bond.
N+N=
N N
Diatomic Molecules: molecules that contain two of
the same atoms.
Example: Iodine, bromine, chlorine, fluorine,
oxygen, nitrogen, hydrogen
POLYATOMIC IONS



tightly bound groups of atoms that behave as a
unit and carry a charge.
there are many of these ions. Since we cannot
predict what their charges will be, you will be
given a list to work with.
when there is a compound formed with
polyatomic ions, both ionic and covalent bonding
takes place.
Properties of Ionic and Covalent Bonding
Properties of Ionic Compounds:
1. They have relatively high melting points due
to strong attraction in bonds
2. They conduct electricity when molten or
dissolved in water (ions can move freely)
3. Those in solid state are not electrical
conductors (ions cannot move)
4. Ionic compounds dissolved in water form
electrolytic solutions. Electrolyte – a
substance that dissolves in water to produce
a solution that conducts electricity.
Properties of Covalent Compounds:
1. They have relatively low melting points due
to weaker attraction in bonds
2. They tend not to conduct electricity when in
solid or liquid state or dissolved in water (do
not form ions; non-electrolytes)