Unit III: Bonding
Textbook Chapters 5,6 & 11
Bonding
 materials held together by the simultaneous
attraction of electrons to two nuclei
What electrons?
How attracted?
2 ways!
Chemical Energy
-a form of potential energy (PE)

PE- energy stored in
molecules
 Substances possess
energy (chemical) due
to their composition
and structure
Energy Changes in Bonding
When 2 atoms are held together by a
chemical bond, generally at a lower
energy condition than when separated.

When a chemical
bond is formed,
energy is released.
 When a chemical
bond is broken,
energy is absorbed.
Bonding and Stability

Generally, systems at
low energy levels are
more stable than at
high energy levels.
 So bonding will more
often occur if change
leads to lower energy
condition (MORE
STABLE).
Bond is formed
-more energy released means
more stable, stronger bond
-less energy given off means
less stable weaker bond

Is Oxygen complete?
 Forms bond, releases energy
exothermic reaction
 Later, breaks bond and
decomposes endothermic rxn
(reverse of exothermic)
4Al(s) + 3O2(g)  2Al2O3(s)
-3351kJ
*from Ref Table I
2Al2O3(s)  4Al(s) + 3O2(g)
+3351kJ
Exact same amount of energy required for reaction
 either released or absorbed (3351 kiloJoules)
Electronegativity

Measure of the ability
of an atom to attract
the electrons that form
a bond between it and
another atom.
 Highest
electronegativity is 4.0
 Reference Table S
Can use electronegativities to predict type of bonds formed.
Bonds Between Atoms

Electrons (valence)
involved in bond
formation can be
transferred, shared
equally or shared
unequally between 2
atoms.
Why Bond?

Become more like
“inert” or noble gases
by completing the
valence shell
 Most stable
Ionic Bonds
aka Electrovalent/Electrostatic

Transfer of one or more electrons from metals to
nonmetals
 form ions
 ions attract (+ and -) {electrostatic force}
 ionic bond formed
Ionic Bonds continued

Form between
elements with
electronegativity
difference  1.7 with
few exceptions.
 *Remember metals
with nonmetals.
Ionic bonds may form between monatomic or
polyatomic ions.
 Polyatomic ions
Monatomic ions
 one atom with charge  compound of 2 or
more covalently
 Na+, F-, Al+3
bonded atoms with a
charge.
 OH-, NH4+, S2O3-2
 Reference Table E.

AgNO3(aq) + NaCl(aq) --> AgCl(aq) + NaNO3(aq)
What is polyatomic ion in above reaction?
What happens to compounds above in water?
Lattice (Binding) Energy
of an Ionic Solid

Measure of the energy
required to completely
separate a mole of a
solid ionic compound
into its separate ions.
 The higher the lattice
energy, the stronger
the ionic bond.
Taken from mikeblaber.org on 7/27/11.
Ionic Solids

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Structural unit: made up of ions.
Ionic bonding
High melting points-->strong forces
Geometric structure, ions held in crystal lattice by
electrostatic attraction
Do not conduct electricity.
When melted or dissolved in water, crystal lattice
is destroyed and ions move freely allowing for
electrical conductivity.
Brittle
Ex: NaCl, KClO3,MgO, KBr, Li2SO4
Covalent Bonds

Simultaneous
attraction of 2 nuclei
for the same electrons
resulting in the sharing
of those electrons.
 Difference in
electronegativities is
less than 1.7(some
exceptions).
Bonding Continuum
Building Molecules with Lewis Dot Structures
Nonpolar Covalent Bonds

Electrons shared
equally between atoms
of the same element.
 Ex: Diatomics
“identical twins”
 H2,N2,O2,F2,Cl2,Br2,I2
 Difference in
electronegativities is
ZERO.
Diatomics can have single,
double or triple covalent bonds.
Structures are symmetrical/nonpolar.
Building Molecules with Lewis Dot Structures
Sigma (s) and Pi (p) Bonds
•Single bonds are sigma
bonds, electron density
is concentrated along the
line that represents the
bond joining the two
atoms.(overlapping s
orbitals)
•Double bonds contain
one sigma and one pi
bond. A pi bond occurs
when the electron
density is concentrated
above and below the line
that represents the bond
joining the two atoms.
(overlapping p orbitals)
•Triple bonds are one
sigma and two pi bonds.
Polar Covalent Bonds

Electrons are shared unequally between atoms of
different elements.
Polar Covalent bonds can help
create both polar(asymmetrical)
and non-polar(symmetrical)
structures.
Remember that the words polar
and nonpolar can be used to
describe both bonds and overall
symmetry of molecules!
Polar Molecules due to either bonding or symmetry
Types of Molecular Shape that
influence Overall Symmetry
Molecule
Bonding
Diagram
Shape
3D
image
Water
Carbon
dioxide
polar
polar
Symmetry of Molecule
bent
asymmetrical
(polar)
linear
symmetrical
(non-polar)
3D
image
Methane
polar
tetrahedral
symmetrical
(non-polar)
Resonance Structures

A hybrid of the possible drawings because
no one Lewis structure can represent the
situation.
Taken from sv.wikipedia.org on 7/27/11.
Taken from en.wikipedia.org on 7/27/11.
The VSPER Model
-molecule will assume the shape that most minimizes electron pair repulsions
Total number of single bonds, double
bonds, and lone pairs on the central
atom
Structural pair geometry
2
3
Linear
planar
Trigonal Planar
4
Tetrahedral
5
Trigonal Bipyramidal
6
Octahedral
VSPER is Valence Shell Electron Pair Repulsion
.
Shape
Coordinate Covalent Bonds

When 2 shared electrons forming covalent bond
are both donated by one of the atoms
 “Free loader”
 Once formed, same as ordinary covalent bond
 Often involved with forming polyatomic ions
Another Coordinate Covalent Bonding Example
Bonding by Jarod Gagnon
Electronegativity
It is the ability
of an atom to take
some electrons to make
it’s outer shell stable
But how is it able?
What is electronegatvity for?
Sometimes the difference is 1.7 or more
That means one atom is so much stronger
it takes the other’s electrons who has them no longer.
If the number is lower than that
the electrons just share and sit where they sat
Then a covalent bond has begun,
but that isn’t the end of the fun.
Sometimes a freeloader comes to the table
and shares some electrons so it can be stable.
The previous compound had no electrons to lack,
which is good for the freeloader gives nothing back.
Covalent Bonds form molecules.

Molecule-discrete particle formed by covalently
bonded atoms where atoms share electrons so that
final electron configuration of each atom is similar to
an inert gas
 Examples: O2, HCl, H2O, CH4, NH3, C6H12O6, CO2
Molecules form molecular substances.

Molecular substances may be gases, liquids or
solids depending on attraction that exists between
the molecules.
Just as we did for Ionic Bonding, let’s compare properties &
characteristics of Molecular substances as solids!
Covalent or Molecular Solids
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Predicting Bonds and Structure?
Building Molecules with Lewis Dot Structures
Structural unit: made
of molecules (covalent
bonds)
Low melting points
Relatively weak forces
Soft
Poor heat conductors
Do not conduct
electricity (good
electrical insulators)
Ex: I2, H2O, CO2,
cellulose C5H10O5
Network Bonding

Certain solids consist of covalently bonded atoms
linked in a network that extends throughout
sample with an absence of discrete particles
 “One Big Giant” Molecule
Just as we did for Ionic & Covalent bonding, let’s
compare properties & characteristics of Network solids!
Network Solids
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Structural unit: made
up of atoms, “One Giant
Molecule”
Very high melting pts.
Very strong covalent
bonds (Network bonds)
Do not conduct in any
phase
Very hard
Ex: Diamond (C),
Quartz/Sand (SiO2),
Silicon Carbide (SiC)
Silicon Carbide grinding wheels
Metallic Bonding

Occurs between atoms
that have a small
number of valence
electrons (metals)
leaving them with
many vacant valence
orbitals and low
ionization energies
 “Electron Sea Model”
Electron Sea Model

Held together by positive kernel and negative
valence electrons. Do you remember kernel?
 Valence e- free to move from atom to atom. How
can this explain a metal’s conductivity, ductility
and malleability?
Just as we did for Ionic, Covalent & Network bonding, let’s
compare properties & characteristics of Metallic solids!
Metallic Solids
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Structural units: made of
positive kernels and
valence electrons.
Intermediate melting pts.
Relatively intermediate
forces (metallic bonds)
Conduct electricity and
heat in all phases.
Malleability, Ductility
and Luster.
Ex: Cu, Na, Fe, K, Au
Can you complete a summary chart that compares
characteristics of all four types of solids?
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4 Types of Solids: Ionic,
Molecular, Network &
Metallic solids
At least 6
properties/characteristics:
type of bonding
structural unit
melting point
conductivity
examples
other
So far we have talked about bonds between
atoms inside molecules or particles!

Intramolecular Forces:
found within molecule
or particle called
chemical bonds.

Think intramurals or intravenous!

6 examples:
Ionic
Polar Covalent
Non-polar covalent
Coordinate covalent
Network
Metallic
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Molecular Attractions aka
Intermolecular Forces
forces beween molecules
Think international or interscholastic

Bonds (Intramolecular forces) build particles.
 Now how do these particles come together to build
something we can actually see?
Dipoles

Dipoles are polar
molecules “2 poles.”
 Have asymmetrical
distribution of
electrical charge
within molecule.
 Dipole-dipole
attraction is force of
attraction between
polar molecules.
Hydrogen Bonding

Special case of dipoledipole that occurs
when hydrogen is
bonded to a small,
highly electronegative
atom (N, O, F)
 Slightly stronger than
other dipole-dipole
attraction.
Hydrogen bonding is FON!!! Or
I’ve had NOF of Hydrogen bonding!!!
Why doesn’t H2O fit with pattern of others?

All are polar, same family, all have same
bonding.
 Smaller mass, should be lower boiling pt.?
 H2O has hydrogen bonding.
London Dispersion Forces
aka Van Der Waals

Weak attractive forces exist
between non-polar molecules
(no dipoles, no H bonding)
 Caused by momentary dipole,
a chance distribution of
electrons.
 Make it possible for small,
non-polar molecules to exist as
solids or liquids under low
temp and high pressure ex.
H2, He, O2, N2
London Dispersion Forces act
over short distances.

Can these forces be
increased?
 London Dispersion
forces increase with
increase in size of
molecules and # of
electrons or a decrease
in distances between
molecules.
Pentane
Changing London Dispersion
Forces influences phase of
matter. Look at halogens for
example!
Molecule-Ion Attraction

Ionic compounds
(salts) are generally
soluble in polar
solvents.
 Remember (aq)!
 Why soluble?
Molecule-Ion Attraction Explained

Polar solvents are asymmetrical (+ & - ends).
 + & - end of liquid are attracted to + & - ends of
ionic salt. This pulls apart ions, breaking crystal
lattice structure (salt is dissolved and broken!)
 If water is solvent, creates hydrated ions: water
molecule surrounded by ions.
 Now that ions are broken free of lattice, what can
the solution now do?
Chemical Symbols

How much lithium is
represented?
 One Li atom or One
mole of Li atoms
(6.02 X 1023)
Chemical Formulas

Both a qualitative and
a quantitative
expression of the
composition of an
element or compound.
 How much sulfuric
acid is represented?
One H2SO4 molecule containing 2 H atoms, 1 S atom & 4 O atoms
OR One mole of H2SO4 molecules containing 2 moles of H atoms,
1 mole of S atoms and 4 moles of O atoms.
2 Types of Chemical Formulas

Molecular Formula:
indicates total number
of atoms of each
element needed to
form a molecule.
 Empirical Formula:
simplest ratio in which
atoms combine to
form a compound
 Empirical formulas do
not always exist in
nature.
For the next section in this unit,
you will need to do the following:

Write chemical
formulas.
 Name chemical
compounds.
 Balance chemical
equations.
Writing Formulas for Ionic Compounds

If you use the criss-cross method, remember to
simplify to lowest values. (Use parentheses if
multiple polyatomic ions are needed).
Write the symbol for the metal ion.
Write the symbol for the nonmetal or
polyatomic ion.
3. Check the oxidation numbers of each ion. If they
add up to zero, this is the formula.
1.
2.
 A Roman numeral after the name of the metal ion
denotes its oxidation number.
Writing Formulas for Ionic Compounds continued
4.
Use the proper subscripts after the symbol
for each ion so that when multiplied times
the oxidation #, the total algebraic sum is
zero. (Use parentheses if multiple polyatomic ions are
needed).

Examples: calcium carbonate, ammonium
sulfite, Nickel (III) sulfide, Copper (II)
chloride
Naming Ionic Compounds Text p.176
1.
2.
3.
4.
Write the name of the cation (metal or
ammonium ion).
Write the name of the anion (nonmetal or
polyatomic ion). Nonmetals ending to ide.
If the metal can have more than one charge
(oxidation number) place a Roman Numeral
after its name to denote the charge.
Examples: FeBr3, K2Cr2O7, Mn(C2H3O2)3
Naming Covalent Compounds Text pp.206-207

For covalent compounds composed of two
elements, name 1st element and then 2nd element’s
name ending is changed to –ide.
 Use prefixes to indicate number of atoms of each
element.
 Examples: CO, CO2, CCl4, P2S5,
1 = mono-
2 = di-
3 = tri-
4 = tetra-
5 = penta-
6 = hexa-
7 = hepta-
8 = octa-
9 = nona-
10 = deca-
Balancing Chemical Equations
Text pp.267-274

Equations must be balanced to support the
Law of Conservation of Mass (matter or mass
can not be created nor destroyed, only rearranged).

To balance, you need to make # of atoms of
each element the same on both reactants’
and products’ side.
Balancing Chemical Equations Continued
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Can’t change the formula, only the coefficients.
Coefficient- a small whole number that appears as
a factor in front of a formula in a chemical rxn.
Hint: When balancing, start with uncommon
elements first. If polyatomic ions appear on both
reactant and product side, balance as a group.
Often it is helpful to save H and O until the end.
When finished, your answer should be simplified
into smallest whole #’s.
Let’s try some examples.
Balancing Equations Examples
3.
Fe3O4 + H2  Fe + H2O
Al + Pb(NO3)2  Al(NO3)3 + Pb
Fe2O3 + CO  Fe + CO2
4.
Ca(OH)2 + (NH4)2SO4  CaSO4 + NH3 + H2O
5.
H3PO4 + CaCl2  Ca3(PO4)2 + HCl
1.
2.