Chapter 12: Chemical Bonding
• General info on Chemical Bonds:
1. Chemical Bond: the force that holds 2 or more
atoms together, and makes them function as a unit.
a. Involve the valence electrons of elements.
b. These valence electrons can be represented by
using an electron dot structure (diagram).
c. Electron dot diagrams are also useful in
illustrating the formation of a chemical bond.
2. The strength of a bond can be determined by
measuring the energy required to break that bond
(called “bond energy”)
Electron Dot Diagrams
• Because valence electrons are involved with
chemical bonds, chemists use a short-hand method
to represent atoms, called “electron dot diagrams
(structures)”.
• Electron dot diagram: has the following:
1. The elements symbol
a. This represents the atoms nucleus and its
inner-core electrons.
2. Dots surrounding the symbol, representing the
atoms valence electrons.
a. There will be 1-8 dots, depending on the element
Electron Dot Diagrams
• Method of creating electron dot diagram:
1. the “dots” representing the atoms valence
electrons are placed one at a time on all 4 sides
of the elements symbol, and only paired after all
4 sides have 1 dot.
2. It doesn’t matter where you place the initial
dot.
Electron Dot Diagrams
• Valence electrons: For the representative
elements, the group number is also the # of
valence electrons.
1. Group 1A has 1 valence electron
2. Group 2A has 2 valence electrons
3. Group 3A has 3 valence electrons
4. Group 4A has 4 valence electrons…….
5. Group 8A has 8 valence electrons
• These valence electrons are represented as the
“dots” drawn around the elements symbol.
Electron Dot Diagrams
• The electron dot diagrams for noble gases will
show a full octet (except for helium)
Electron Dot Diagrams
• Examples:
Types of Chemical Bonds
• Two types of chemical bonds exist:
1. ionic bonds, forming ionic compounds.
2. Covalent bonds, forming covalent compounds.
• Ionic Bonds: formed when a metal reacts with a
non-metal.
1. The metal atom(s) easily give-up one or more of
their valence electrons to expose a full octet of inner
s & p orbitals.
2. The non-metal, having an affinity (desire) for
electrons, will take on the extra electron(s), which
will give it a full octet in its outer s & p orbitals.
Types of Chemical Bonds
• Electron Affinity: a measure of the attraction of an
atom for electrons
1. noble gases have low electron affinity ( already
have a stable octet)
• Reactivity of an element is directly related to its
valence electrons.
1. elements will “react” in order to acquire the
stable electron structure of a noble gas.
2. Forming a positive ion: atoms will lose 1+ of their
valence electrons in order to obtain a noble gas
configuration.
3. Forming a negative ion: atoms will gain 1+
electrons to obtain a noble gas configuration.
Covalent Bonds
• There are compounds that form where both
elements making up the compound need to gain
electrons in order to have a stable octet.
1. They will “share” electrons to become stable.
2. Bonding follows the “Octet rule”.
Octet rule: Elements will form bonds with other
elements in order to obtain a stable octet of
valence electrons.
3. this is called a covalent bond.
Covalent Bonds
• Covalent Bond: A chemical bond that results from
the sharing of valence electrons
1. Shared electrons are considered to be a part of
the complete outer energy level of both atoms.
2. Generally:
a. Mostly occurs in bonds between non-metals.
b. Occurs with elements close to each other on the
periodic table.
3. Molecule: formed when 2 or more atoms bond
covalently.
4. most common materials we encounter have
covalent bonds (clothing, food)
Covalent Bonds
• Formation of Covalent bonds:
1. Diatomic molecules: all of them share electrons
(H2, N2, O2, F2, Cl2, Br2, & I2)
2. Example: Fluorine (group 7A; 7 valence electrons)
a. As 2 fluorine atoms approach each other, the
protons of one will be attracted to the electrons of the
other.
b. They will get close enough to achieve maximum
attraction (too close, and they will repel each other)
c. At this point, they are sharing an electron, so both
fluorine atoms have a stable octet.
d. Each fluorine atom will now have 1 “bonding pair”
of electrons, and 3 “lone” (unshared) pairs of electrons.
Covalent Bonds
• Types of Covalent Bonds:
1. Single covalent Bonds: A covalent bond in
which only a single pair of electrons is shared.
a. The shared electron pair = “Bonding pair”
b. Bonding pairs are represented in electron dot
diagrams called Lewis structures.
Lewis structure: model that uses electron-dot
structures to show how electrons are arranged
in molecules.
c. Bonding pairs are represented as a pair of
dots, or a line drawn between atoms.
Covalent Bonds
• Examples of single covalent bonds:
1. Hydrogen molecule (H2) = H-H or H:H
2. All halogens (group 7A) have 7 valence electrons,
need only 1 more to form an octet, and therefore
will share one pair of electrons, forming a single
covalent bond.
a. Can form a single covalent bond between
themselves (F2, Cl2, I2), and with other elements.
3. Group 6A has 6 valence electrons, and need 2
more to form an octet. (will share 2 pairs to do it)
a. Will form 2 separate single covalent bonds.
b. Ex. H2O (on next slide)
Single covalent Bonds
• Example H2O Lewis Structures:
• A set of dots can be drawn to represent the
bonding pair of electrons, or a single line.
Single Covalent Bonds
• Group 5A elements have 5 valence electrons, and
need 3 more to form an octet.
1. they will “share” 3 electrons with other
atoms, forming 3 single covalent bonds.
2. Example: ammonia NH3
3. Example:
nitrogen triflouride NF3
Single Covalent Bonds
• Group 4A elements have 4 valence electrons, and
need 4 more to form an octet.
1. they will “share” 4 electrons with other atoms,
forming 4 single covalent bonds.
2. Example: methane CH4
• Single covalent bonds are also called “sigma bonds”
1. symbolized by greek letter sigma σ
2. the shared electrons will be concentrated into a
“bonding orbital” between the 2 atoms
Bonding orbital: localized region where bonding
electrons are most likely found.
Multiple covalent bonds
• Multiple covalent bonds: a bond in which atoms
share more than one pair of electrons.
1. Elements that commonly form multiple
covalent bonds are carbon, nitrogen, oxygen &
sulfur.
2. Double covalent bond: one in which 2 pairs of
electrons are shared between 2 atoms.
a. Two lines are drawn between 2 atoms to
indicate the presence of a double covalent bond.
b. Ex: oxygen molecule O2
Multiple covalent bonds
• Triple covalent bond: one in which 3 pair of
electrons are being shared between 2 atoms.
1. ex: nitrogen molecule N2
• Bond dissociation energy:
1. energy is released when a
bond forms.
2. energy must be added to break a bond.
Bond Dissociation energy: the amount of energy
required to break a specific covalent bond.
3. triple bonds require more than double, which
requires more than single covalent bonds.
Molecular structure
• Molecular structure: the use of models to predict
the arrangement of atoms in a molecule.
1. Structural formula: Uses elements symbols
and bonds (using dots or lines) to show relative
positions of atoms.
2. The structural formula can be predicted for
many molecules by drawing the Lewis Structure
for the molecule.
3. There are steps that should be followed to
draw a Lewis Structure for a covalent compound
Steps to determine Lewis Structures
• Step 1: predict the location of certain atoms.
1. Hydrogen is always a terminal (end) atom.
2. The atom with the least attraction for shared
electrons in the molecule will be the central
atom.
a. The element is usually the one closest to the
left side of the periodic table (if a choice is there).
b. Central atom: located in the center of the
molecule, with all other atoms being terminal
atoms.
Steps to determine Lewis Structures
• Step 2: Find the total number of electrons
available for bonding (bonding pairs).
1. Count the total number of valence electrons of
all the atoms in the molecule.
2. Divide this number by 2 to determine the
number of bonding pairs.
• Step 3: apply the bonding pairs of electrons in a
specific way the molecule.
1. Place one bonding pair (representing a single
bond) between the central atom and each of the
terminal atom.
Steps to determine Lewis Structures
• Step 3 (cont.):
2. Subtract the number of pairs used from the
original total number of bonding pairs.
a. The remaining electron pairs will be applied as
lone pairs of electrons, or used in double or triple
bonds.
• Step 4: Place lone pairs around each terminal atom
that is bonded to the central atom
1. This is done to satisfy the octet rule for each
terminal atom.
2. The exception to this is hydrogen as a terminal
atom.
Steps to determine Lewis Structures
• Step 5: Assign any remaining bonding pairs of
electrons to the central atom
• Step 6: Evaluate if the central atom has an octet
around it (4 electron pairs).
1. If central atom has an octet, and all bonding pairs
have been used, you are done!
2. If central atom doesn’t have an octet, then you
must convert one or more of the lone pairs of
electrons on the terminal atom(s) to a double or
triple bond between that terminal atom and the
central atom.
a. Often, carbon, nitrogen, oxygen or sulfur are the
central atoms (for double & triple bonds)
Steps to determine Lewis Structures
• Additional step: if Lewis structure being drawn is for
a polyatomic ion.
1. The atoms within the polyatomic ion are
covalently bonded to each other.
2. When determining the number of bonding pairs
of electrons that exist for the polyatomic ion, you
must take into account the charge of the ion.
3. Find the total number of valence electrons of all
involved atoms, then add/subtract electrons to
match the charge of the polyatomic ion.
a. Apply the electrons pairs as previously described.
Example of a Lewis Structure for a
Polyatomic Ion
• Chlorite ion: ClO2-
Lewis structures with 2 central atoms
• Some molecules have 2 central atoms that are
bonded to each other, and have terminal atoms on
each of them.
• Ex: C2H4 (ethene)
1. follow same steps to form the Lewis structure
• Exceptions to octet rule:
1. Incomplete octets
2. Expanded octets
a. always apply extra pairs to central atom
Exceptions to Octet Rule
• Some molecules do not obey the octet rule.
• 3 Reasons:
1. There are a small group of molecules that have
an odd number of valence electrons, and can’t
form an octet around each atom:
a. Ex: NO2N = 5 valence electrons
O = 6 valence electrons
*5+6+6 = 17 valence electrons, therefore you
can’t make exact pairs (one left over)
b. CLO2- and NO are additional examples
Exceptions to Octet Rule
• Reason # 2: certain compounds will form with
fewer than 8 electrons present around its central
atom (rare)
1. Ex: BH3
• Reason # 3: certain compounds have central
atoms that have more than 8 valence electrons,
called an “expanded octet”.
1. Ex: PCl5
• Any lone pairs not used should be applied to
central atom.
Resonance Structures
• When there are more than one correct Lewis
structures that can be drawn for a molecule, that
molecule is said to “show resonance”.
Resonance: a condition that occurs when more
than one valid Lewis structure can be written for
a molecule or ion.
1. Often seen when a molecule has a double or
triple bond.
2. The two or more correct Lewis structures
representing a single molecule/ion are called
“resonance structures”.
Resonance Structures
• When drawing resonance structures, the position
of the atoms will remain the same.
1. the difference between the resonance
structures is the position of the lone and bonding
pairs of electrons
2. Ex: NO33. Ex: CO2
Molecular Shape
• General info:
1. The shape of a molecule determines its physical &
chemical properties.
2. The shape of a molecule is determined by the
overlap of orbitals that share electrons.
3. Many chemical reactions depend on the ability of
2 compounds to contact each other.
a. The shape of molecules dictates how close they
will get to one another
4. A Lewis structure must be drawn first in order to
determine molecular shape.
5. The model used to determine molecular shape is
Valence Shell Electron Pair Repulsion Model (VSEPR)
VSEPR Model
• Model is based on an electron arrangement that
minimizes the repulsion of shared/unshared pairs of
electrons around the central atom.
1. This repulsion results in atoms being oriented at
fixed angles to each other.
2. The angle formed by any 2 terminal atoms and
the central atom = Bond Angle
• Shared pairs of electrons occupy a bonding orbital
which is smaller
• Lone pairs of electrons are not shared by 2 different
nuclei, and therefore occupy a larger orbital
1. They will push the shared bonding orbitals closer
together (and give shape to the molecule)
VSEPR Model
• Molecules will have different geometric shapes,
depending on:
1. The number of bonding pairs of electrons
between the central atom & terminal atoms.
2. The presence or absence of lone pairs of
electrons on the central atom.
• Molecular Shapes:
1. Linear
2. Trigonal Planar
3. Tetrahedral
4. Trigonal Pyramidal
5. Bent
Molecular Shape
• Molecules containing NO lone pairs of electrons on
the central atom:
1. Ex: Beryllium Chloride
2. Bonding electrons will seek
maximum separation.
3. Molecular shape will be Linear
4. Bond Angle: 180°
5. Has 2 pairs of electrons shared off the central
atom.
6. Doesn’t exhibit polarity (non-polar molecule)
Molecular Shape
• Molecules containing NO lone pairs of electrons on
the central atom:
1. Ex: Carbon Dioxide
2. Bonding electrons will seek
maximum separation.
3. Molecular shape will be Linear
4. Bond Angle: 180°
5. Has 4 pairs of electrons shared off the central
atom
6. Doesn’t exhibit polarity (non-polar molecule)
Molecular Shape
• Molecules containing NO lone pairs of electrons on
the central atom:
1. Ex: Aluminum Chloride
2. Bonding electrons will seek
maximum separation.
3. Molecular shape will be Trigonal Planar
4. Bond Angle: 120°
5. Has 3 pairs of electrons shared off the central
atom
6. Doesn’t exhibit polarity (non-polar molecule)
Molecular Shape
• Molecules containing NO lone pairs of electrons on
the central atom:
1. Ex: methane
2. Bonding electrons will seek
maximum separation.
3. Molecular shape will be Tetrahedral
4. Bond Angle: 109.5°
5. Has 4 pairs of electrons shared off the central
atom.
6. Doesn’t exhibit polarity (non-polar molecule)
Molecular Shape
• Molecules containing ONE lone pair of electrons
on the central atom:
1. Ex: PH3
2. Lone pair of electrons takes up greater space,
forcing the 3 single covalent bonds away.
3. Molecular shape will be Trigonal pyramidal
4. Bond Angle: 107.3°
5. Has 3 pairs of electrons shared off the central
atom, and 1 lone pair.
6. Exhibits polarity (polar molecule)
Molecular Shape
• Molecules containing TWO lone pairs of electrons
on the central atom:
1. Ex: Water
2. Two lone pairs take up greater space, forcing
the 2 single covalent bonds away.
3. Molecular shape will be Bent
4. Bond Angle: 104.5°
5. Has 2 pairs of electrons shared off the central
atom, and 2 lone pairs.
a. Lone pairs will both be on one side.
6. Exhibits polarity (polar molecule)
Electronegativity Difference & Bond
Character
• The type of chemical bond that will form
between 2 elements can be predicted by
determining the electronegativity difference of
the 2 atoms that are bonded.
• When covalent bonds form between atoms of
different elements, there will be an unequal
sharing of their electron pairs, due to differences
in electronegativity values.
1. This unequal sharing of the bonding pairs of
electrons results in a Polar Covalent Bond
Electronegativity Difference & Bond
Character
• When covalent bonds form between atoms of
different elements, large differences in
electronegativity values between the atoms will result
in an ionic bond.
1. Bonding is therefore often not clearly ionic or
covalent.
2. As the difference in electronegativity increases, the
bond will become more ionic.
3. With atoms of the same type (diatomic), the
electronegativity difference is zero, resulting in a Pure
Covalent Bond.
4. Smaller electronegativity differences will result in a
Polar Covalent Bond.
Electronegativity Difference & Bond
Character
Electronegativity Difference & Bond
Character
• Electronegativity difference of zero = pure covalent
• Electronegativity difference of 1.70 = 50% covalent
50% ionic
• Electronegativity difference above 1.70 = more ionic
Polar Covalent Bonds
• Between pure covalent bonds and ionic bonds, there
is a range of bonds that are polar covalent.
1. One of the atoms in the bond are more
electronegative than the other, causing an un-equal
sharing of the bonding electron pair (like un-equal
sides in a tug-of-war contest).
2. These shared electrons will be pulled closer to
the more electronegative atom, causing a “polarity”
to the molecule (one side of molecule will be slightly
positive, the other side slightly negative.
3. The delta symbol “δ” is used with a +/- sign
drawn next to the molecule to show which side is
slightly positive, and which is slightly negative.
Polar Covalent Bonds
• The result of a polar covalent bond: a Dipole (twopole) molecule.
1. one side will be slightly negative δ-, and the
other side slightly positive δ+.
2. Example: H2O
3. polar molecules will be
attracted by an electrical field.
4. Some molecules exhibit polarity between their
central & terminal atoms, but their overall polarity
is balanced, so that the molecule is considered to
be non-polar.
Polar vs Non-polar molecules
• Polar molecules:
1. Attracted by an electrical field
2. Have lone pairs of electrons, which push away
bonding pairs.
3. Have slightly positive and slightly negative sides.
4. Are asymmetrical in shape.
• Non-polar molecules:
1. Are not attracted by an electrical field
2. Do not have any lone pairs of electrons.
3. Are symmetrical in shape.