Unit 5 – Bonding &
Nomenclature
Ionic Bonds, Covalent Bonds, and
Writing Formulas
In our last unit…
•We talked all about the ATOM
– Development of the atomic model, subatomic
particles
•HOWEVER, we were only talking about atoms
as one individual unit
In this unit we will be looking at how atoms
combine (or BOND) with other atoms to
form compounds!
Electrons are Arranged in Shells
– As we move down a family, the elements gain an
electron shell.
• The first shell can hold up to 2 electrons.
(Row/Period 1)
• The second and third shells can hold up to 8
electrons each. (Periods 2 & 3)
• The fourth and fifth shells can hold up to 18
electrons each. (Periods 4 & 5)
• The sixth and seventh shell can hold up to 32
electrons each. (Periods 6 & 7)
The Valence Electrons Give Atoms Their
Properties
• The outer most shell is called the valence shell.
• The electrons in the valence shell are called the
valence electrons.
• Atoms tend to gain or lose electrons to end up
with a full valence shell. (8 e-)
– Except H2 and He which want a duet (2 e-)
• The number of valence electrons correspond to
the digit in the ones place of the group number
for elements in groups 1, 2, 13-18.
Electron Dot Diagrams Show Valence
Electrons
• An electron dot diagram depicts an atom as its
symbol and its valence electrons.
– Ex: Carbon
Carbon has four electrons in its valence shell
(carbon is in group 14), so we place four dots
representing those four valence electrons
around the symbol for carbon.
Electron Dot Diagrams Show Valence
Electrons
• Electrons are placed one at a time in a
clockwise manner around the symbol in the
north, east, south and west positions, only
doubling up if there are five or more valence
electrons.
– Example: Chlorine (7 valence electrons b/c it is in
group 17)
Paired and Unpaired Electrons
• As we can see from the chlorine example, there are
six electrons that are paired up and one that is
unpaired.
• When it comes to bonding, atoms tend to pair up
unpaired electrons.
– They do this in two ways that we will discuss in the
chapter.
1. A bond that forms when one atom gives an unpaired
electron to another atom is called an ionic bond.
2. A bond that forms when atoms share unpaired
electrons between each other is called a covalent
bond.
Ionic Bonding
• Ionic Bond
– Bond between metal and nonmetal due
to “electrostatic interactions”
– Electrons are transferred from metal to
nonmetal forming a positively charged
metal ion (cation) and negatively charged
nonmetal ion (anion)
6.1 Ions and Ion Formation
• If one or more valence electrons are lost or gained,
the balance between positive and negative
particles in the atom is upset.
– When an atom gains electrons, there are more negative
charges in the atom than positive charges, so it becomes
negative
• Example – Oxygen gains two electrons so it gains a charge of -2.
– When an atom loses electrons, there are more positive
charges in the atom than negative charges, so it
becomes positive.
• Example – Magnesium loses two electrons so it gains a charge
of +2.
Atoms Do What is Easiest For Them!
Atoms tend to do the least to gain a filled valence
shell.
• Elements in groups 1, 2 and 13 tend to lose
electrons to wind up with a filled valence shell
because it only involves the removal of 1, 2, or 3
electrons, not the addition of 7,6, or 5 electrons
respectively.
• Likewise for elements in groups 15,16, and 17, only
in their case, they tend to gain 3,2, or 1 electrons,
not lose 5,6, or 7 electrons respectively.
Examples
• Sulfur
• Sodium
Get Out Your Periodic Tables!
• What are the charges for groups 1, 2, 13, 15,
16, 17, and 18?
– Group 1:
– Group 2:
– Group 13:
– Group 15:
– Group 16:
– Group 17:
– Group 18:
Atoms That Gain Or Lose Electrons Are
Called Ions
(we saw this last chapter!)
• An ion is formed when the number of
electrons differs from the number of protons
in an atom.
– This is accomplished by gaining or losing
electrons, not by adjusting the number of
protons!
– The charge of an ion is denoted as a superscript in
the atom’s isotope notation.
• Ex: A Nitrogen ion:
6.2 Ionic Compounds
• When an atom that tends to lose electrons
comes in contact with an atom that tends to
gain them, an electron transfer occurs.
– As a result, two oppositely charged ions are
formed.
– The two oppositely charged ions are attracted to
each other via the electrostatic attractions.
– These attractions create an IONIC BOND
Ionic bonds Result from a Transfer of
Valence Electrons
+
-
Ionic Bonds Result from a Transfer of
Valence Electrons
• Ionic Compounds - All chemical compounds
containing ions (one cation, one anion)
• Ionic compounds typically are formed from
elements on opposite sides of the periodic
table.
• For all ionic compounds, positive and negative
charges must balance out.
– Formula must be neutral overall!
Properties of Ionic Compounds
• Positive and negative ions come together to
create an organized pattern called a lattice
structure.
• An example of a
lattice structure
for sodium
chloride
Properties of Ionic Compounds
•
•
•
•
Ionic compounds tend to be brittle
They have high melting and boiling points
They tend to be soluble in water.
They can conduct electricity when dissolved in
H2O
Many of the physical and chemical properties of
the components of the compound are changed
when they bond.
6.3 Writing Ionic Formulas
+3
+2
+1
-2
-1
+3
+2
+1
-2
-1
-2
-1
6.3 Writing Ionic Formulas
+3
-2
+3
-2
-2
Subscripts Show the Number of Atoms of Each
Element in a Compound
• A subscript notes how many atoms of each
element are in a subunit of a compound.
• 1 is NEVER used as a subscript.
• Ex: NaCl  1 Sodium atom : 1 Chlorine atom
Al2O3  2 Aluminum atoms : 3 Oxygen atoms.
• The subscripts determine how many of each
atom is required to balance out the positive
and negative charges from the ions.
Writing Formulas for Ionic Compounds
• We can use the charges of the ions in a
compound to easily determine the formula of
the compound.
• All you do is crisscross the charges to make them
subscripts (Reduce the subscripts if you can!!)
– Ex: magnesium nitride
+2
-3
Mg3 N2
Don’t forget to drop the positive/negative signs
once they’re subscripts!!
PRACTICE, PRACTICE, PRACTICE!!
Write formulas for the following ionic compounds:
Magnesium chloride
Calcium bromide
Sodium oxide
Beryllium oxide
Potassium iodide
Polyatomic Ions
• There are some ions that are made of more
than one element
– These ions are called POLYATOMIC IONS
• “poly” meaning more than 1, “atomic” meaning atoms 
more than 1 type of atom in the ion
• Examples: (see chart in your notes for more)
– Sulfate
SO42– Hydroxide
OH– Nitrate
NO3-
Writing Formulas w/ Polyatomic Ions
1. Write both the ions with their accompanying charges.
2. Crisscross the charges to make them subscripts.
– Use parenthesis around the polyatomic ion if there is more
than one of the ion!
3. Drop the positive and negative signs once they are
subscripts.
4. Reduce subscripts if possible.
Example:
Barium phosphate
PRACTICE WITH POLYATOMICS!
Write formulas:
Don’t forget to reduce subscripts that
come from the charges if you can!
calcium phosphate
beryllium bicarbonate
lithium carbonate
Magnesium sulfate
Barium hydroxide
Writing Formulas w/ Transition Metals
• Transition metals – all elements in groups 3-12
and elements in groups 13, 14, and 15 below the
stair step line
– Can have more than one charge in ion form!
• Except Ag+ and Zn2+
• When writing formulas for compounds involving
transition metals, you will get the charge on the
metal cation from the Roman numeral in
parenthesis after the metal’s name
Writing Formulas w/ Transition Metals
Writing Formulas of Compounds w/ Transition Metals:
1. Write the metal’s atomic symbol with the value of the
Roman numeral as a positive charge
2. Write the value of the anion (negatively charged ion)
with its charge
3. Crisscross the charges to get your formula
Example:
Lead (IV) hydroxide
Writing Formulas w/ Transition Metals
Examples:
Copper (II) chloride
Tin (IV) sulfate
Gold (I) nitrate
Don’t forget to reduce subscripts that
come from the charges if you can!
On a separate piece of paper:
Please write formulas for the following compounds & turn them in to
me on a separate piece of paper so I can check them.
1. Sodium fluoride
2. Magnesium phosphate
3. Copper (II) sulfate
4. Tin (IV) phosphide
5. Hydrogen cyanide
6.4 Naming Ionic Compounds
• When naming ionic compounds, we take the
name of the metal followed by the name of
the non-metal, only we drop the last syllable
and add –ide as an ending.
– Examples: NaCl: Sodium Chloride
Al2O3: Aluminum Oxide
PRACTICE, PRACTICE, PRACTICE!
Write names for the given ionic compounds:
LiBr
AlCl3
BeO
Ra3N2
Naming Ionic Compounds with
Polyatomic Ions
• The rules for naming polyatomic ions are the
same rules for naming normal ionic
compounds except you use the polyatomic
ion name for the second word
Example:
MgSO4 – magnesium sulfate
LiNO3 – lithium nitrate
Ba(OH)2 – barium hydroxide
PRACTICE WITH POLYATOMICS!
Name the compounds:
RbNO3
NaHCO3
Mg(OH)2
Naming Formulas w/ Transition Metals
• Transition metals – all elements in groups 3-12
and elements in group 14 beneath the stair step
line
– Can have more than one charge in ion form!
• Except Ag+ and Zn2+
• When naming compounds involving transition
metals, you need to include a roman numeral in
parenthesis to indicate the charge of the ion
Naming Compounds w/ Transition Metals
Naming Compounds Containing Transition Metals:
1. Name the metal from its symbol
2. Determine the charge on the metal by multiplying the
subscript on the anion by the charge the anion
normally has (see box below)
3. Write the charge as a roman numeral in parenthesis
4. Write the name of the anion either as an element with
an “–ide” ending or as the unchanged name of the
polyatomic ion
Charge of metal = |charge of anion × anion subscript|
Naming Compounds w/ Transition Metals
Examples:
CrCl6
Fe(OH)2
NiO
Covalent Bonding
Covalent Bonds
–Bonds in which e- are shared
between two atoms
–Most common type
6.5 Covalent Bonds Result from a
Sharing of Valence Electrons
• Elements that have a partially filled valence
shell can complete them by sharing electrons
with another element
• The mutual attraction for shared electrons is
called a covalent bond (co- signifies sharing,
-valent signifies valence electrons)
Shared Electrons Complete Shells
F
F
6.5 Covalent Bonds Result from a
Sharing of Valence Electrons
• A substance made up of atoms which are held
together by covalent bonds is a covalent
compound.
–They are also called molecules.
Drawing Electron Dot Diagrams for
Molecules
• Chemists usually denote a shared pair of
electrons as a straight line.
F F
• Sometimes the nonbonding pair of electrons
are left off of the electron dot diagram for a
molecule
Nonmetals Tend to Form
Covalent Bonds
• A covalent bond is formed when two atoms
that tend to gain electrons come into contact
with each other.
• Hydrogen tends to form covalent bonds
because it has a fairly strong attraction for an
additional electron.
• The number of covalent bonds an atom can
form is equal to the number of additional
electrons it can attract.
Examples
H2
H H
H2O
O H
H
More Examples
CH4
H
H C H
H
NH3
H N H
H
Multiple Bonds
• Sometimes two atoms share more than just
one pair of electrons which results in a double
bond or triple bond.
O2
O O
CO2
O C O
Writing Formulas for Covalent Compounds
(2 non-metals bonded together)
1. Write the symbol for the first element listed
• If there is a prefix on the first element, write the
prefix value as a subscript attached to the first
element’s symbol
2. Write the symbol for the second element listed
• If the prefix on the second element is “di-” or
greater, write the prefix value as a subscript attached
to the second element’s symbol
Ex: Diphosphorous monoxide 
Writing Formulas for Covalent Compounds
(2 non-metals bonded together)
Examples:
Diphosphorous hexoxide
Carbon tetrachloride
Nitrogen trihydride
Naming Covalent Compounds
(2 non-metals bonded together)
Given the formula:
1. Write the names of the 2 elements
2. Attach a Greek prefix (see chart) to the beginning of
the first element only if it has a subscript greater
than or equal to 2
3. Attach a Greek prefix to the second element no
matter what!
4. Change the ending of the second element’s name to
an –ide ending
Ex: Se2O3 
Naming Covalent Compounds
(2 non-metals bonded together)
Examples:
P2O6
CCl3
NO
In review…
Chemical Bonds
– Force that holds atoms together
– It’s all about the electrons (e-)
Types of Chemical Bonds
– Ionic
• Bond between metal and nonmetal due to “electrostatic
interactions”
• Electrons are transferred from metal to nonmetal
In review…
Covalent Bonds
– Bonds in which e- are shared
– Most common type
Metallic Bonds
– Atoms are bonded
to one another
(not to other
elements)
– Positive ions in a “sea” of negative charge (e-)
– Example: Cu, Ag