Unit 3: Bonding
AP Chemistry
A. Types of Chemical Bonds- A bond will be defined as a
force that holds groups of two or more atoms together.
1.
Why bonds form- chemical bonds form because they
lower the potential energy of the particles in the atom.
a. Repulsion between the negative electron clouds of
each atom. Repulsion between the nuclei.
b. Attraction between the positive nuclei and the
negative electron clouds.
c. As the optimum distance is achieved that balances
these forces, there is a release of energy and a
bond is formed. This is an exothermic process.
Likewise, breaking a chemical bond absorbs
energy and this is considered an endothermic
process.
d. The atoms involved and the energy difference in
their bonding decides which type of bond forms.
2.
Ionic Bonds- this occurs when an atom that loses
electrons relatively easily reacts with an atom that has a
high affinity for electrons. Metals reacting with
nonmetals.
a. A transfer of electrons occurs to form ions.
b. If the atom loses an electron it becomes a cation.
c. If the atom gains the electron it becomes an anion.
d. Unlike charged ions will attract to form the ionic
compound.
3.
Properties of Ionic solids- almost all ionic compounds
are solids at room temperatures and pressures. Think
of: salt, rubies, quartz, etc.
a. The solid has a definite geometric structural unit
called a crystal.
b. Ions in a crystal are arranged to minimize the
repulsion forces.
c. The precise form of the crystal depends on the
kinds of ions in the compound, their sizes, and the
ratio in which they appear. Forms a 3D crystal
lattice.
d. The bonds in an ionic crystal are very strong,
which is why almost all ionic compounds are solid
at room temperature.
e. High temperatures are required to break most ionic
bonds, free the ions from one another, and melt
the crystal to become a liquid.
f. Conducts electricity only when dissolved in water.
g. Hard, brittle, high M.P.
Bond Strengtha. Energy released in the formation of 1 mole of ions is
called lattice energy.
b. Lattice energy depends on:

Charges of ions: greater the charge, lattice
energy increases, more exothermic (more
negative).

Separation distance: larger the ion, smaller the
lattice energy, less exothermic (less negative).
4.

𝐸∝
𝑄1 ∙𝑄2
𝑟
(Coulomb’s Law)
Q = magnitude of charges
r = distance of separation
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5.
Covalent bond- one or more electrons are
simultaneously attracted to two nuclei .
a.
b.
c.
d.
The bonding atoms’ orbitals overlap, which
maximizes attraction between nuclei and bonding
electrons.
The electrons are attracted simultaneously by the
nuclei of each atom.
The shape usually isn’t considered a crystal lattice
as in an ionic bond.
Atoms share 2, 4 or 6 electrons to create bonds.
1. 2 (single), 4 (double), 6 (triple) bond
2. multiple bonds reduce bond distance

bond distance < sum of atomic radii

shorter bond distance = stronger bond
6.
Properties of covalent molecules- can be any state of
matter at room temperature. Think of: candle wax,
water, oxygen gas.
a. Generally lower melting points
b. Don’t conduct electricity
e. Lower melting points and boiling points.
7.
Two types of Covalent bonds:
a. Nonpolar covalent bond- electrons are shared
equally between atoms.

Occurs between atoms of the same element,
diatomic molecules in other words: (H2, O2, N2 …)

Each atom has the same attraction for the shared
pair of electrons.
b.

Polar covalent bond- electrons are not completely
transferred, but unequal sharing of electrons
results.
One atom will have a stronger attraction for the
shared electrons than the other, due to
electronegativity difference.
8.
Metallic bonds- bonding in metals.
a. loosely held valence electrons a.k.a. delocalized
electrons, of metals form a “sea of electrons”
surrounding the metal ions.
b. Properties:

excellent electrical and thermal conductivity

Ductility, malleability

Very strong bonds, high melting point
B. Electronegativity- A property that describes the ability of an
atom in a molecule to attract shared electrons. A “tug-of-war”
for electrons.
1. The difference in electronegativity values between the
atoms involved in a bond determines what type of bond
it is.
a. bond types are on a continuum, nonpolar at one
end, ionic at the other.
b.
0—0.4 = nonpolar
0.4—2.0 = polar covalent
>2.0 = ionic
2. Electronegativity can be estimated from the periodic
table. Increases across a period, and decreases down a
group.
a. Shielding effect explains why EN decreases down a
group. The core electrons shield the nucleus’ ability
to “pull” or attract to the valence electrons.
b. EN increases across a period due to increasing Zeff.
Nucleus increasingly pulls outer energy level in
closer, holds onto electrons tighter.
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3.
Bond Polarity. The greater the  EN of a bond the
greater it’s polarity.
a. It is an imbalance of charge distribution. It is
expressed with a dipole arrow.
b.
c.
Bond strength increases with polarity.
Entire molecules can be described as being polar.
It depends on:
1. the polarity of the individual bonds and
2. the geometry of the molecule.
C. Lewis structures- Since bonding just involves the valence
electrons, the Lewis Structure is a representation of a
molecule that shows how the valence electrons are arranged
among atoms.
1. Add up all the valence electrons from all the atoms
2. Draw a skeleton structure.

H always terminal

more electronegative element terminal

least electronegative element central
3. Use one pair of electrons to form a bond between each
pair of bound atoms. Use a line instead of two dots.
4. Arrange remaining electrons to satisfy the octet rule.
5. Resonance is a condition that occurs when more than
one valid Lewis structure can be drawn.
a. Electrons are actually delocalized over several
atoms or bonds.
b. Delocalized electrons lowers potential energy of
molecule.
6. Formal Charge- fictitious charge to help differentiate
competing lewis structures.
a. FC= valence e–– [nonbonding e- + ½ bonding e-]
b. Preferred strcture:

atoms have formal charges closest to zero

negative formal charge reside on the more
electronegative atom (upper right most on the
periodic table)
7.
Exceptions to the octet rule.

B, Be, can have less than an octet; Be forms two
single bonds and B forms three single bonds.

Third period and beyond. These elements can
have an expanded octet due to the unfilled dorbitals. Any extra electrons are placed around the
central atom.
D. Bond energy- energy required to break 1 mole of the bond
in the gas phase. kJ/mole
1. Greater the bond energy, stronger the bond.
2. Multiple bonds=stronger
3. To calculate the enthalpy (overall energy change) of a
reaction:
∆𝐻𝑟𝑥𝑛 = Σ(𝑏𝑜𝑛𝑑𝑠 𝑏𝑟𝑜𝑘𝑒𝑛) + Σ(𝑏𝑜𝑛𝑑𝑠 𝑓𝑜𝑟𝑚𝑒𝑑)
positive


negative
break bonds, absorb energy, endothermic
form bonds, release energy, exothermic
E. VSEPR theory- allows one to predict the 3-dimensional
shape of molecules from knowledge of their Lewis Dot
structure.
1. The main idea is that the structure around a given atom is
determined by minimizing repulsion between electron pairs.
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2.
This means that the bonding and non-bonding electron pairs
around a given atom are positioned as far apart as possible.
3.
RULES:
a. maximum separation between electron pairs
b. atom positions define molecular geometry
c. lone electron pairs squeeze bond angle (actual angle <
ideal angle)
Electron
Domains
2
Domain
Geometry
–
:
2
0
3
0
2
1
4
0
3
1
2
2
5
0
4
1
180o
109.5o
90o
120o
5
6
Bond
Angle
120o
3
4
Molecular
Geometry
3
2
2
3
6
0
5
1
4
2
90o
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F. Valence Bond Theory (Hybrid Orbitals)- a chemical bond
between two atoms is the result of direct overlap of two
atomic orbitals (one on each atom).
1. Good orbital overlap requires that the atomic orbitals on
each atom (those orbitals overlapping to form the bond)
be oriented directly toward the other atom
2. Orbitals from bonding atoms merge, which allows
single electrons from each atomic orbital to occupy
overlapping area and simultaneously attract both
nuclei.
3. Hybridization- the mixing of two or more atomic orbitals
of similar energies.
a. The number of atomic orbitals mixed to form the
hybrid orbital equals the total number of pairs of
electrons.
b.
c.
Sigma bonds (𝝈)- electrons being shared are
located directly between the two bonded nuclei.
Pi bonds (𝝅)- parallel overlap of p orbitals, the
electrons are spread out across the entire molecule,
they’re delocalized.
G. Nomenclature- rules used to name binary compounds.
1.
2.
3.
Type I- metals and nonmetals. Metal is group 1 or 2.
a. Name metal (cation) then nonmetal (anion), drop
ending of anion and add (ide).
b. To write correct formula be sure oxidation numbers
sum to zero. Use subscripts for additional ions.
Type II- metals and nonmetals, but metal is usually a
transition metal.
a. Name metal and use a roman numeral to indicate
charge of metal, then name anion as before.
b. To write correct formula be sure oxidation numbers
sum to zero.
Type III- nonmetals only.
a. The first element in the formula is named first, and
the full element name is used.
b. The second element is named as though it were an
anion.
c. Prefixes are used to denote the numbers of atoms
present.
d. Prefix mono- is never used for naming the first
element.
e. Prefixes: use the prefixes to determine the
subscripts of the atoms.
mono = 1
di
=2
tri- = 3
tetra = 4
penta = 5
4.
hexa = 6
hepta = 7
octa = 8
nona = 9
deca = 10
Polyatomic ions are named without using any prefix
names.
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