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Chemistry Unit 5
Chemical Bonding
Why Do Atoms Bond?
• To become more stable
• like the noble gases.
• Octet Rule – atoms tend to gain, lose or
share electrons in order to acquire a full set
of valence electrons. (usually 8)
Three Main Types of Bonds
• Ionic Bond – Atoms transfer electrons to fill
their valence shells, oppositely charged
ions are formed, opposites attract.
• Occurs between a metal and a nonmetal
• Covalent Bond – Atoms share electrons to
fill their valence shells.
• Occurs between nonmetals
• Metallic Bonds – Atoms share a “sea of
electrons.”
• Occurs between metal atoms
Properties of Metallic, Molecular
and Ionic Compounds
Metallic
Bonds
Insouluble
Conductor
Mod. High
Melting
Points
Molecular
Ionic
Compounds Compounds
Some sol. in
water/some Soluble in
in nonpolar
water
sol’n
NonConductive
Conductor
in sol’n
Low
Melting
points
Very High
Meting
Points
Ionic Bonding
• Ion – a charged particle
• A neutral atom becomes an ion when it loses or
gains an electron.
• If an atom loses an electron, it becomes a (+) ion
called a cation.
• If an atom gains an electron, it becomes a (-) ion
called an anion.
Ionic Bonding
• Example
Na
To become
more stable,
sodium
must lose
one electron
Cl
To become
more stable,
chlorine
must gain
one electron
Ionic Bonding
• Example
Na
Cl
Sodium loses
Chlorine gains
an electron
an electron and
and becomes
becomes a Cl-1
an Na+1 ion.
ion.
Opposites attract, and an ionic compound is formed…
NaCl
Try Another Example
Al
Aluminum will
become more
stable if it gets
rid of three
electrons.
Br
Bromine will
become more stable
if it receives one
electron.
Are both atoms more stable as a result of this
transefer? No, Al must donate two more… where?
Aluminum & Bromine
Br
Al
Br
Br
Now, each atom has a full valence shell… all are
more stable.
Aluminum and Bromine
Aluminum
donated 3
e-, so it
becomes
Al+3
Br
Al
Br
Br
Each bromine
accepted 1 e-, so they
each become Br-1
The compound that forms is AlBr3
Let’s Wrap it Up
• Ionic bonds are held together by electrostatic
forces.
• The result of an ionic bond is called an ionic
compound.
• Ionic bonds form between a metal and a nonmetal
atom due to large differences in electronegativity.
(1.7 or greater)
• The nonmetal’s EN is so much greater than the
metal’s EN that it removes the metal’s valence
electron. Electrons are transferred.
For Example: Na and Cl
EN of Na = 0.9
EN of Cl = 3.0
Why does Sodium and Oxygen
form an ionic bond?
3.0 EN of O
- 0.9 EN of Na
2.1 Difference in EN
• Difference in electronegativity is 2.1(>1.7)
• An ionic bond will form.
• Chlorine has a greater electronegativity, and
is able to yank electrons away from sodium.
Covalent Bonding
O
O
Each atom of Oxygen needs two more electrons to
become more stable.
They will share two pairs of electrons.
A diatomic molecule of oxygen is formed.
O2
Try another example
Oxygen
needs two
electrons to
become more
stable.
H
O
H
Each atom of hyd
rogen needs one
more electron to
become more
stable.
All atoms become more stable
(have full valence shells). A
molecule of water is made.
H2O
Let’s Wrap it Up… Again!
• Covalent bonds are held together by a mutual
need for the shared electrons (electronegativity)
Their orbits overlap. Each electron is attracted to
the positive charge of the opposite nucleus.
• The result of a covalent bond is called a
molecule.
• Covalent bonds form between two nonmetals due
to a small (or no) difference in EN. (less than
1.7)
• Neither atom’s EN is strong enough to remove
the other atom’s electrons. Electrons are shared.
Polar and Nonpolar Covalent
Bonds
• If one nonmetal has a greater EN than the
other, it can “hog” the shared electrons.
This forms a POLAR covalent bond. (EN
difference greater than 0, but less than 1.7)
• If the nonmetals have the same EN, they
will share equally and form NON POLAR
covalent bond.
(0 diff. in EN)
For Example: N and O
EN of N = 3.0
EN of O = 3.5
Why does Nitrogen and Oxygen
form a Covalent Bond?
3.5 EN of Oxygen
- 3.0 EN of Nitrogen
0.5 = difference in EN
Difference in EN is less than 1.7, therefore a covalent
bond will form.
Difference in EN is greater than 0, therefore the
covalent bond will be polar. (Unequal sharing of e-)
One Final Example
If Chlorine bonds with Chlorine (a diatomic
molecule), the difference in EN would be
“0”, thus a nonpolar covalent bond will
form. (Equal sharing of e-)
Molecular Geometry
• Linear molecules: atoms are connected in a
straight line.
• All molecules with only 2 atoms are linear.
• Many molecules with 3 atoms are also linear.
• Ex. O2, HCl, CO2
O
C
O
Molecular Geometry
• Bent: bonded atoms have a bent shape due
to unshared pairs of electrons.
• Unshared electron pairs exert a greater
repulsion force than the electron pairs in the
bonds.
• Ex. H2O, NH3
H
H
Molecular Geometry
• Tetrahedral: one atom bonded to four other
atoms.
• The angle between any two bonds is 109.5o.
• Ex. CH4 (methane)
H
C H
H
H
Writing Ionic Formulas
Calcium Chloride
• Locate the metal on the periodic table and
write the element symbol with its oxidation
number.
Ca
+2
Writing Ionic Formulas
• Locate the nonmetal on the periodic table
and write the element’s symbol with its
oxidation number.
-1
Cl
+2
Ca
-1
Cl
• Find the common factor between the two oxidation
•
•
•
•
•
numbers.
In this case, 2.
Decide how many of each ion is needed to make
the charge equal to the common factor.
In this case, 1 calcium ion (+2) and 2 chlorine ions
(-1 and –1 = -2). Compounds are neutral.
Use this number of ions as the subscript for the
element, and write the formula.
In this case, Ca Cl2.
Writing Ionic Formulas Part 2
Aluminum Oxide
• Locate the metal on the periodic table and
write the element symbol with its oxidation
number.
Al
+3
Writing Ionic Formulas Part 2
• Locate the nonmetal on the periodic table
and write the element’s symbol with its
oxidation number.
-2
O
+3
Al
-2
O
• Find the common factor between the two oxidation
•
•
•
•
•
numbers.
In this case, 6.
Decide how many of each ion is needed to make
the charge equal to the common factor.
In this case, 2 aluminum ions (+3 and +3 = +6) and
3 oxygen ions (-2 and -2 and -2 = -6). Compounds
are neutral.
Use this number of ions as the subscript for the
element, and write the formula.
In this case, Al2O3.
Try these examples on your own.
•
•
•
•
•
Sodium and Oxygen
Lithium and Sulfur
Aluminum and Chlorine
Potassium and Nitrogen
Magnesium and Fluorine
Naming Ionic Compounds
• Write the name of the metal.
• Write the name of the nonmetal with the ending
changed to –ide.
Example:
Nitrogen = nitride
Oxygen = oxide
Phosphorus = phosphide
Fluorine = fluoride
Sulfur = sulfide
Chlorine = chloride
Iodine = iodide
Bromine = bromide
Naming Ionic Compounds
Al2S3
• Write the name of the metal.
Aluminum
• Write the name of the nonmetal, changing
the ending to –ide.
Sulfide
• Name the compound.
Aluminum Sulfide
Naming Ionic Compounds
BaCl2
• Write the name of the metal.
Barium
• Write the name of the nonmetal, changing
the ending to –ide.
Chloride
• Name the compound.
Barium Chloride
Try these examples on your own.
•
•
•
•
•
BeF
Li20
B2S3
Mg3N2
CaCl2
Transition Metals
Wtg. Formulas / Nmg. Compounds
• Most transition metals can form ions with
more than one charge.
• Examples: Copper atoms can become
Cu +1 and Cu +2 ions
Iron atoms can become
Fe +2 and Fe +3 ions
• Therefore, the oxidation number for the metal
will be given to you as a roman numeral in the
name of the compound.
Writing Formulas w/Transition Metals
Iron (III) Oxide
• Write the symbol for the transition metal.
Ex. Fe
• Take the number in parentheses and write it
as the oxidation number.
Ex. Fe +3
Writing Formulas w/Transition Metals
Iron (III) Oxide
• Write the symbol for the nonmetal.
Ex. O
• Look up its oxidation number on the
periodic table, and add it to the symbol.
Ex.
O -2
Writing Formulas w/Transition Metals
Fe +3
O –2
• Find the common factor between the two
oxidation numbers. In this case = 6
• Decide how many of each ion is needed to make
the charge equal to the common factor. In this
case 2 Fe and 3 O ions.
• Use this number of ions as the subscript for the
element, and write the formula.
Fe2O3
Copper (I) Sulfide
• Write the symbol for the transition metal.
Ex. Cu
• Take the number in parentheses, and write it
as the oxidation number.
Ex. Cu +1
Copper (I) Sulfide
• Write the symbol for the nonmetal.
Ex. S
• Look up its oxidation number on the
periodic table, and add it to the
symbol.
Ex.
S -2
Copper (I) Sulfide
Cu +1
S –2
• Find the common factor between the two
oxidation numbers. In this case = 2
• Decide how many of each ion is needed to make
the charge equal to the common factor. In this
case 2 Cu and 1 S ion.
• Use this number of ions as the subscript for the
element, and write the formula.
Cu2S
Naming Compounds
w/Transition Metals
FeO
• Look up the nonmetal on the periodic table.
Oxygen O-2
• Look up the metal on your ion chart. Find
the possible oxidation numbers.
Fe +2 or Fe +3
Fe
+2
or Fe
+3
-2
O
• Decide which ion will form in the proper ratio
with the known charge on the oxygen ion.
FeO
• Iron bonds in a 1 to 1 ratio with oxygen, therefore,
the iron ion must have a +2 charge. (Fe+2)
• Name the compound, indicating the oxidation
number of the metal in parenthesis.
Iron (II) Oxide
Fe2O3
• Look up the nonmetal on the periodic table.
Find its oxidation number.
Oxygen O-2
• Look up the metal on your ion chart. Find
the possible oxidation numbers.
Fe +2 or Fe +3
Fe2O3
• Decide which ion will form in the proper ratio
with the known charge on the oxygen ion.
Fe +2 or Fe +3
• Iron bonds in a 2 to 3 ratio with oxygen. Three
oxygen atoms will have a charge of -6. Therefore,
two iron ions must equal +6. It must be Fe +3.
• Name the compound, indicating the oxidation
number of the metal in parenthesis.
Iron (III) Oxide
Polyatomic Ions
Writing Formulas / Naming Compounds
• A polyatomic ion is a covalent molecule
•
•
•
•
that has an ionic charge. (As opposed to
being a neutral molecule.)
Poly = many
Atomic = atoms
Ion = charged particle
A charged particle that consists of more
than one atom.
Polyatomic Ions
Examples:
Sulfide 
Nitride 
Phosphide 
Chloride 
Sulfate
Nitrate
Phosphate
Chlorate
=
=
=
=
SO4-2
NO3-1
PO4-3
ClO3-1
• Notice the ending has changed to –ate.
Polyatomic Ions
Examples:
Sulfide 
Nitride 
Phosphide 
Chloride 
Sulfite
Nitrite
Phosphite
Chlorite
=
=
=
=
SO3-2
NO2-1
PO3-3
ClO2-1
• Notice the ending has changed to –ite.
Polyatomic ions
• Not all polyaomic ions end in -ate or -ite.
• Some other examples:
Ammonium
NH4+1
Hydroxide
OH-1
• Some Polyatomic ions contain more than
two elements. Ex. Acetate = C2H3O2-1
Calcium Phosphite
• Write the symbol for the metal. Add the
oxidation number from the periodic table.
Ca+2
• Write the formula for the polyatomic ion
from the ion chart. Add its oxidation
number.
PO3-3
Calcium Phosphite
Ca +2 PO3 -3
• Determine the common factor of the two oxidation
numbers. In this case, 6.
• Decide how many of each ion is needed to equal
the common factor. In this case, 3 calcium ions
and 2 phosphate ions.
• Write these numbers as the subscript for each ion.
Ca3(PO3)2
• Notice that the polyatomic ion must be placed in
parenthesis or, instead of 2 phosphate ions, you
would have 32 Oxygen atoms and 1 Phosphorus
atom.
Ca3PO32
An Additional Example
• Aluminum Hydroxide
Al+3
OH-1
• The least common factor is 3. Therefore, 1
aluminum ion will bond with 3 hydroxide ions to
form a neutral compound.
Al(OH)3
• If you omitted the parenthesis, you would not have
3 hydroxide ions. Instead you would have 3
hydrogen atoms and one oxygen atom. AlOH3
Try these!
• Write formulas for the following compounds.
 Lithium sulfate
 Calcium acetate
 Aluminum nitrite
 Magnesium phosphate
 Sodium carbonate
Answers
•
•
•
•
•
Li2SO4
Ca(C2H3O2)2
Al(NO2)3
Mg3(PO4)2
Na2CO3
Lithium sulfate
Calcium acetate
Aluminum nitrite
Magnesium phosphate
Sodium carbonate
Naming Compounds
w/ Polyatomic Ions
KClO3
• Write the name of the metal.
Potassium
• Write the name of the polyatomic ion from the ion
chart.
Chlorate
• Name the compound.
Potassium Chlorate
Mg3(SO3)2
• Name the metal.
Magnesium
• Name the polyatomic ion from the ion
chart.
Sulfite
• Name the compound.
Magnesium Sulfite
Try these
• Write names for the following compounds.
 CaCO3
 Al2(SO3)3
 Ca(ClO2)2
 K3PO4
 Mg(OH)2
Answers
• CaCO3
• Al2(SO3)3
• Ca(ClO3)2
• K3PO4
• Mg(OH)2
Calcium carbonate
Aluminum sulfite
Calcium chlorate
Potassium phosphate
Magnesium hydroxide
Covalent Molecules
Use prefixes to designate the number of atoms of each element
used in the molecule.
Prefix
mono
di
tri
tetra
penta
Number
1
2
3
4
5
Prefix
Number
hexa
hepta
octa
nona
deca
6
7
8
9
10
Writing Covalent Formulas
Dinitrogen Pentoxide
• Write the symbol of each element.
N
O
• Add the subscript as indicated by the prefixes.
N2O5
Writing Covalent Formulas
Carbon Dioxide
• Write the name of each element.
C
O
• Add the subscripts as indicated by the prefixes.
CO2
Try these examples on your own.
•
•
•
•
•
Sulfur dioxide
Sulfur monoxide
Carbon tetrachloride
Dihydrogen dioxide
Nitrogen triiodide
Naming Covalent Molecules
NH3
• Write the name of the first nonmetal using its
subscript as a prefix.
Nitrogen (No prefix written for the first
element IF it is a one.)
• Write the name of the second nonmetal using its
subscript as a prefix and change the ending to -ide.
Trihydride
• Name of the molecule: Nitrogen trihydride
Naming Covalent Molecules Pt.2
P2O5
• Write the name of the first nonmetal using the
subscript as a prefix.
Diphosphorus
• Write the name of the second nonmetal using the
subscript as a prefix and change the ending to -ide.
Pentoxide
• Name the molecule: Diphosphorus Pentoxide
Try a few examples on your own.
•
•
•
•
•
CO
CO2
SF2
PI3
H2O
Naming Acids
• Acids are water solutions of certain
hydrogen compounds.
• There are two main types of acids:
• Binary Acids – Hydrogen + a nonmetal
• Tertiary Acids – Hydrogen + a polyatomic ion
Binary Acids
• To name the acid.
Hydro _____ ic acid. (The blank is the root
of the nonmetal.)
For example:
HCl = Hydrochloric acid
HI = Hydroiodic acid
HBr = Hydrobromic acid
Binary Acids
• To write the formula:
Hydrofluoric acid
Write the symbol for hydrogen.
Write the symbol for the nonmetal.
Use the ox. numbers to figure out the ratio.
H +1 F-1
HF
Binary Acids
• Hydrosulfuric acid
Write the symbol for hydrogen.
Write the symbol for the nonmetal.
Use the ox. Numbers to figure out the ratio.
H+1 S-2
H2S
Tertiary Acids
Hydrogen plus a polyatomic ion.
If the anion ends in –ate, _____ic acid.
If the anion ends in –ite, _____ ous acid.
(I ate it and it was icky.)
(Rite ous!)
Tertiary Acids
• To name an -ate acid:
H3PO4 = Phosphoric acid (phosphate ion)
HClO3 = Chloric acid (chlorate ion)
HNO3 = Nitric acid (nitrate ion)
• To name an –ite acid:
H2SO3 = Sulfurous acid (sulfite ion)
HClO2 = Chlorous acid (chlorite ion)
HNO2 = Nitrous acid (nitrite ion)
Tertiary Acids
• To write formulas for ___ ic acids:
Carbonic acid: Write symbol for hydrogen.
Write polyatomic –ate ion.
H+1 CO3-2 (carbonate ion)
H2CO3
• To write formulas for ___ous acids:
Clorous acid: Write the symbol for hydrogen
Write the polyatomic –ite ion.
H+1 ClO2-1 (chlorite ion)
HClO2
Properties of Water
• The unique properties of water are due to the
strong intermolecular HYDROGEN BONDS
that are formed between the polar water
molecules. (Opposite poles attract. The
positive hydrogen end of one molecule of
water attracts to the negative oxygen end of
another molecule of water.)
Hydrogen Bonding
• This hydrogen bonding is more
extensive in ice than it is in liquid
water. For this reason, ice is less
dense than water.
• Hydrogen bonding in water gives it a
high surface tension.
Properties of Water
• Water is called the universal
solvent due to its polarity. (Pulls
ionic compounds apart)
• Water has a very high specific heat.
(It takes a lot of energy to
overcome the forces between the
molecules and make them move
faster and heat up) This moderates
climates near large bodies of water.
Properties of Water
• Water molecules are adhesive, they
stick to other things. (Forms a
meniscus, leaves containers wet)
• Water molecules are also cohesive,
they stick to each other. (Capillary
action – draws water up through
plants, trees, etc.)