Unit 3: Chemical Formulas
and
Bonding
Electron Dot diagrams
• Electron dot diagrams show the valence
electrons around an atom. In most
molecules and compounds a complete octet
is achieved for each atom:
Al
N
• Most monatomic ions have an electron
configuration of noble gases:
7 valence e-s
F
+ e- 
F
8 valence e-s
Drawing Lewis Dot Structures
• To visualize valence e-, we will use Lewis Dot
Structures.
• Step 1: The element symbol represents the
nucleus and all e- except valence.
• Step 2: From the periodic table, determine
the number of valence e-.
• Step 3: Each “side” of symbol represents an
orbital. Draw two dots on one side, then one
for each of the remaining three sides.
Additional electrons should then be paired.
Lewis Dot Structures
• Ex: carbon
step 1: C
step 2: 4 valence e-s
step 3:
C
Lewis Dot Structures
• Ex: bromine
step 1: Br
step 2: 7 valence e-s
step 3:
Br
Chemical Bonding
What holds things together?
Let’s examine the melting point of compounds
across two periods. What is the trend?
Conductivity - high
Conductivity - low
Nonconductive
Chlorides of Period 2
compound
LiCl
BeCl2
BCl3
CCl4 NCl3 OCl2
melting point
610
415
-107
-23
-40
Cl2
-121
-102
NaCl MgCl2 AlCl3 SiCl4 PCl3
SCl6
Cl2
801
-51
-102
Chlorides of Period 3
compound
melting point
714
high
193
-69
-112
low
Bonding
How can we explain the melting point behavior
across a period?
Bonding between atoms changes across a period…
• Bonding involves the valence electrons or
outermost shell (or highest shell) electrons
• Atoms form bonds to become more stable –
electrons are gained, lost or shared to achieve
stability.
• The properties of a compound are different from
the properties of the atoms that make up the
compound. Ex: NaCl
Types of Bonds
1. Ionic bond
Transfer of e- from a metal to a
nonmetal and the resulting electrostatic
force that holds them together forms
an ionic compound.
EX: Na+ + Cl-  NaCl
(neutral)
Ionic Bonding
Ionic bonds involve the formation of positive
and negative ions that then attract each
other.
Metals form positive ions by losing electrons
Nonmetals form negative ions by gaining
electrons
Next Slide
Ionic Bonding Example 1
Sodium has 1 valence electron which it needs to
lose.
Na
Cl
Chlorine has 7 valence electrons and
needs to pick up 1 electron.
Next Slide
Ionic Bonding Example 1
The sodium loses its electron to the chlorine.
+1
-1
Na
Cl
This makes the sodium +1 and the chlorine -1
They attract each other forming the compound
NaCl
Next Slide
Ionic Bonding Example 2
Magnesium has 2 valence electrons which it needs
to lose.
O
Mg
Oxygen has 6 valence electrons, It
needs to pick up 2 electrons.
Next Slide
Ionic Bonding Example 2
Magnesium loses both of its outer electrons to the
oxygen.
Mg
Next Slide
O
Ionic Bonding Example 2
This gives the magnesium a +2 charge
and the oxygen a -2 charge
-2
+2
Mg
O
They join together to form the compound
MgO.
Next Slide
Exchange of Electrons
Ionic Bonding
When atoms bond, the properties of the new
compound are DIFFERENT from the properties of
the elements that made them up.
Ionic compounds have several characteristics in
common due to the presence of the ionic bond.
These characteristics include:
Crystalline structure (the formula gives the ratio between the
ions making up the substance)
High melting points, making them solids at room temperature
Usually water soluble (can dissolve in water)
Electrolytes when in solution (conduct electricity)
Ionic Bonding
Sodium chloride (NaCl) is held together by an
ionic bond.
The properties of sodium chloride are:
Sodium chloride forms a cube
shaped crystalline solid.
Melting point = 801˚C
Boiling point = 1413˚C
Highly soluble in water
Strong electrolyte
Cl-
Na+
Na +
ClNa +
Cl-
Na +
Cl-
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Types of Bonds
2. Covalent bond
Formed from the sharing of e- pairs
between two or more nonmetals
resulting in a molecule.
EX: H2 + O  H2O
Covalent Bonding
Definition - bond formed due to the sharing
of electrons between nonmetals.
The high attraction for electrons of
nonmetals results in the nonmetals
attempting to remove electrons from each
other.
Since neither nonmetal is able to give up
electrons they are forced to share the
electrons.
Covalent Bonding Example 1
Bromine and Fluorine both have 7 valence
electrons and very high attraction for
electrons.
Br
F
Covalent Bonding Example 1
Since neither fluorine or bromine are able to
lose electrons they get drawn together until
their outer orbits overlap and one electron
from each atom goes back and forth between
the two atoms.
Br
F
This creates the compound BrF.
Covalent Bonding Example 2
Hydrogen and Oxygen can both pick up electrons.
(If hydrogen loses its only electron it will end up
as a nucleus with no electrons around it.) They
will share electrons to form covalent bonds.
This results in the formula H2O
H
O
H
Covalent Bonding Example 3
Seven of the elements have such high attraction
for electrons that they will never exist as
individual, unattached atoms. Anytime these
elements are present in pure form they will bond
to other atoms of the same element.
For example a fluorine atom will readily bond to a
second fluorine atom.
Resulting in F2.
F
F
Covalent Bonding
These elements are called diatomic
elements, and the molecules they form are
called diatomic molecules. The definition
of a diatomic molecule is:
A molecule made up of 2 atoms of the same
element.
Covalent Bonding
The seven diatomic elements are:
Hydrogen, H2
Nitrogen, N2
Oxygen, O2
Fluorine, F2
Chlorine, Cl2
Bromine, Br2
Iodine, I2
Covalent Bonding
Covalent compounds are made up of small
units called molecules. The formula for a
covalent compound tells the actual number
of atoms, of each element, found in each
molecule.
For example: The formula for water is H2O.
This formula indicates that each water
molecule is made up of two Hydrogen atoms
and one Oxygen atom.
What is a “polar” covalent bond?
• Covalent bonds involve sharing
electrons.
• The electrons may be shard equally
(nonpolar covalent) or unequally (polar
covalent).
• Example: H2 shares electrons equally,
but HCl does not. Therefore, H2
contains a nonpolar covalent bond and
HCl contains a polar covalent bond.
Polarity of water
• Take for example H2O. When we draw
the structure it looks like:
O
H
H
• The oxygen atom pulls electrons away
from the hydrogen atoms. This unequal
sharing results in polar bonds, which have
a “more negative” end and a “more
positive” end.
Polar Water Molecule
more negative
EX:
O
more positive
H
H
H2O is a polar molecule.
Bond polarity affects the properties of a
material such as melting and boiling
points, crystal structure and acidity.
Covalent Bonding
The characteristics shared by covalent
compounds are:
Molecular structure –individual units
Low melting and boiling points, most covalent
compounds are gases or liquids at room temperature
(the larger the molecule the higher its melting and
boiling point)
Soluble in covalent solvents such as alcohol or
benzene.
Nonelectrolytes
gas atoms and molecules
Covalent Bonding
Paradichlorobenzene (moth balls)
(C6H4Cl2) is a covalent compound.
Its properties are:
Molecular solid
Cl
Melting point = 53.1˚C
H
Boiling point = 174.55˚C
Soluble in alcohol, ether, acetone
and benzene
Nonelectrolyte
H
C
H
C
C
C
C
C
Cl
H
Comparison of Bonding Types
ionic
covalent
ions
molten salts
conductive
molecules
nonconductive
Both
transfer of determined sharing of
by valence electrons
electrons
electrons
high mp
low mp
not usually
water
water soluble
soluble
The properties of a material depend on the structure
-different bond types result in different properties.
Polyatomic Ions/Radicals
Some groups of atoms are covalently bonded
together so strongly that the stay together
during chemical reactions and act as a single
unit.
These groups of atoms become charged, with
the charge being spread out through out the
group.
These groups are called polyatomic ions or
radicals.
Polyatomic Ions/Radicals
Definition - groups of atoms bonded together
that act as a charged unit.
Examples:
Ammonium - NH4+1
Sulfate - SO4-2
Acetate - C2H3O2-1
Phosphate - PO4-3
What kind of bond do you think hold the
atoms in a polyatomic ion together?
What kind would hold two polyatomic ions
together?
Type of bond? – Ionic, Polar Covalent, or
Nonpolar Covalent?
TiO2
CH4
NaI
CS2
O2
KCl
CsF
HBr
AlCl3
Types of Bonds
3. Metallic bond
Metals bonding with other metals do not
gain or lose e- or share e- unequally.
These bonds are created from the
delocalized e- that hold metallic atoms
together.
Chemical Formulas
• A chemical formula is a combination of
symbols that represents the
composition of a compound.
• Chemical symbols are used to indicate
types of elements present.
• Subscripts are used to indicate the
number of atoms for each element
present.
What are the “parts” of a
formula?
chemical symbols
C8H18
number of atoms of each element
• 8 atoms of carbon
• 18 atoms of hydrogen
Charges of Monatomic Ions
• Because atoms want to reach an
octet of valence electrons, the
oxidation numbers, (positive or
negative charges) can be predicted
for single atoms (monatomic).
• Metals tend to have positive
oxidation numbers. (lose e-)
• Nonmetals tend to have negative
oxidations numbers. (gain e-)
0
1+
2+
3+
varies +1 to +7
varies
3- 2- 1-
Oxidation Numbers (charges)
of Polyatomic Ions
• Polyatomic ions are ions that are made
up of two or more atoms.
• Refer to your table of Polyatomic ions.
• Polyatomic ions generally have the
following endings: “ate” or “ite”
Ex:
NO2- nitrite
PO43- phosphate
SO42- sulfate
Polyatomic Ions
Oxidation Numbers
• The sum of the oxidation numbers in
a compound must equal zero.
Ex:
CaCl2 = Ca+2 + Cl- + Cl2 positive charges –
2 negative charges = 0
• The charge on a monatomic ion is its
oxidation number.
Ex: Ba+2 has an oxidation of +2
Cl- has an oxidation of -1
What happens when the
predicted charge can vary??
The oxidation number of a transition
element is shown using Roman numerals
to indicate the charge.
The Roman numeral indicating oxidation.
Ex: iron (II) is Fe+2
iron (III) is Fe+3
Why is aluminum oxide Al2O3?
Writing Ionic Formulas
Ionic compounds are composed of
metals and nonmetals.
• Ionic compounds are made from the
gaining or losing of electrons and
the resulting electrostatic force
that holds the ions together.
• The sum of the oxidation numbers
in a compound must equal zero.
Writing Ionic Formulas
• When writing formulas, the cation
(metal ion) is always written before
the anion (nonmetal ion).
• When using polyatomic ions, refer to
charge given on your table.
NOTE: There is only one polyatomic
cation (NH4+).
The rest are all are polyatomic
anions.
Criss Cross Method of Writing
Formulas
Notice a trend between the
oxidation/charges of ions and the
subscripts of elements.
Ex:
Mg+2 and Cl- gives
MgCl2
Ex:
Al+3 and SO4-2 gives
Al2(SO4)3
Criss Cross Method of Writing
Formulas
• This method crosses charges and
subscripts to form neutral
compounds.
Al +3 and O-2
Al
and O
Al2O3
(neutral)
Criss-Cross Method of
Writing Formulas
Ex:
Lead (II) phosphate
Pb+2 and PO4-3
Pb
and PO4
Pb3(PO4)2
Nomenclature
• Nomenclature is defined as a naming
system.
• Chemistry uses nomenclature to
standardize names of chemicals.
• Let’s take a look.
Naming Binary Ionic
Compounds
Rules for naming binary (composed of
two) ionic compounds:
1. Name of cation is given first. The
name of the cation is the same as
the element.
2. Name of anion is given last. The
name of the anion is the same as
the element, but with an “ide”
suffix.
Naming Ionic Compounds
Ex:
Al2O3
Aluminum and Oxygen
Aluminum oxide
cation
anion
Naming Ionic Compounds
Ex:
Ni2O3
Note that the cation has MORE THAN
ONE possible oxidation state, so
Roman numerals are needed to
identify the ion.
Nickel and Oxygen
+3
-2
Nickel (III) oxide
Naming Ionic Compounds
EX: AgCl is __________________
Silver chloride
Na2O is __________________
Sodium oxide
CaBr2 is ___________________
Calcium bromide
PbO2 is ___________________
Lead (IV) oxide
Naming Ionic Compounds with
a polyatomic ion
Rules for naming compounds that
contain a polyatomic ion.
1. Cation rule from binary applies.
2. Anion takes the name of the
polyatomic ion as found on
the table.
Ex:
Al2(SO4)3
aluminum sulfate
Ex:
Mg(OH)2
magnesium hydroxide
Naming Ionic Compounds
Ex: Li2CO3 is ____________________
Lithium carbonate
Ba(OH)2 is ___________________
Barium hydroxide
Zn(NO3)2 is __________________
Zinc nitrate
KClO3 is ___________________
Potassium chlorate
Naming continued…
• Name the following compound:
• Ba(Na)2
Banana
Naming Binary Molecular
Compounds
• Unlike ionic compounds, molecular
compounds are composed of
individual covalently bonded units, or
molecules.
• Covalent compounds are formed
between nonmetals.
• Prefixes are used to indicate number
of each type of element in the
compound.
• Write the prefixes as indicated on
the next slide….
Molecular Prefixes
1
mono-
6
hexa-
2
di-
7
hepta-
3
tri-
8
octa-
4
tetra-
9
nona-
5
penta-
10
deca-
(decade)
Naming Binary Molecular Compounds
Follow these rules:
1. The element written first is given a
prefix if it contributes more than one
atom to the molecule.
2. The second element is named by
combining (a) a prefix indicating the
number of atoms contributed by the
element, (b) the root of the name of the
second element, and (c) the ending -ide.
3. The o or a at the end of a prefix is
usually dropped when the word following
the prefix begins with another vowel.
Ex: monoxide or pentoxide
Naming Binary Molecular
Compounds
Ex:
P4O10
1. P has more than one atom in this
molecule.
Tetraphosphorus
2. O is named by combining prefix,
root name, and -ide ending.
decoxide (a is dropped from prefix)
*combine to form: Tetraphosphorus
decoxide
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Naming Binary Molecular
Compounds
Ex: SO3 is ____________________
sulfur trioxide
PBr5 is ____________________
phosphorus pentabromide
ICl3 is _____________________
iodine trichloride
H2O is _____________________
dihydrogen monoxide
Sb2O3 is _____________________
diantimony trioxide
(*metalloid)
• THE END
Strange Names for Molecules
BUCKMINSTER
FULLERENE
MORONIC ACID
Paper Chromatography.mov
Chromatography Lab
Paper Chromatography Paper chromatography is a
method chemists use to separate compounds from one
another, but not change them. In this lab we will explore
how this separation is made using different dye
compounds. Molecules with similar polarities or molecular
structures are attracted to each other. Water molecules
have a polar structure.
Because of this structure the oxygen end of the molecule
has a small negative electrical charge and the hydrogen
end has a small positive charge. Liquid water is held
together by the attraction between the charges on
different molecules.
• A more complex, yet still similar
molecule is cellulose, a molecule which is
the basic component of paper. It is a
very long molecule (a polymer) in which
thousands of rings of six atoms each
are linked together like beads. A portion
of a cellulose molecule is shown below.
• Paper Chromatography Paper chromatography
is a method chemists use to separate
compounds from one another, but not change
them. The polar regions of these molecules
are attracted to polar regions of the cellulose
chains (which help to hold the fibers together
in paper). Not surprisingly, water molecules,
being polar, are also attracted to these
regions and when paper is wet it loses
strength because the water molecules get
between the cellulose chains and weaken the
attraction between them.
• When water molecules move up paper that is
dipped in water, the molecules which might be
dissolved in the water will also be carried
along up the paper. This is applied to the
separation of dyes in a technique known as
paper chromatography.
•A spot of dye is placed on the paper above the level of
the water. As the water moves up, the dye molecules will
move with it if they are more strongly attracted to the
water molecules than to the paper molecules. If the dye
molecules are more strongly attracted to the paper than
to the water, they will move more slowly than the water
or even not at all.
•What if the dye is a mixture? If two or more dyes have
been mixed, then they may move at different rates as
the water moves up the paper. If this happens, they will
separate and we can identify them . This is depicted in
the sketches below.
Yellow #5 Structure
• C16 H9 N4 Na3 O9 S2
• Skittles ingredients: sugar, corn syrup,
hydrogenated palm kernel oil, apple juice from
concentrate, citric acid, dextrin, natural and
artificial flavors, gelatin, food starch,
coloring (includes Yellow 6 lake, Red 40 lake,
Yellow 5 lake, Blue 2 lake, Blue 1 lake, Yellow
5, Red 40, Yellow 6, Blue 1), ascorbic acid
(Vitamin C).
• M&M ingredients: Milk Chocolate (Sugar,
Chocolate, Cocoa Butter, Skim Milk, Milkfat,
Lactose, Soy Lecithin, Salt, Artificial
Flavors), Sugar, Cornstarch, Corn Syrup,
Dextrin, Coloring (Includes Blue 1 Lake, Red
40 Lake, Yellow 6, Yellow 5, Red 40, Blue 1,
Blue 2 Lake, Yellow 6 Lake, Yellow 5 Lake, Blue
2), Gum Acacia.
* The “lake” part means that the dye is attached as a coating