Chemical Bonding
Chemical Bonds
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A bond is a force that holds groups of
two or more atoms together and
makes them function as a unit.
Bond energy is the amount of energy
required to break a particular bond.
*determines the strength of a bond
Bonding Review
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What type of elements form ionic bonds?
What type of elements form covalent
bonds?
Where are these types of elements found
on the periodic table in relationship to
each other?
Chemical Bonding: An Overview
• Valence electrons play a fundamental role in
chemical bonding.
• When metals and nonmetals combine, valence
electrons usually are transferred from the metal to the
nonmetal atoms, giving rise to ionic bonds.
• In combinations involving only nonmetals, one or
more pairs of valence electrons are shared between
the bonded atoms, producing covalent bonds.
• In losing, gaining, or sharing electrons to form
chemical bonds, atoms tend to acquire the electron
configurations of noble gases.
Types of Chemical Bonds:
Ionic Bonding
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Ionic bonding occurs when there is an attraction
between oppositely charges ions. Electrons are
transferred to form oppositely charged ions.
Ionic compounds result when an atom that loses
electrons relatively easily (a metal) reacts with an atom
that has a high affinity for electrons (a non-metal).
Example: Sodium chloride, NaCl
Na+ & Cl- ions react to form solid
sodium chloride.
Types of Chemical Bonds:
Covalent Bonding
Covalent Bond: A bond that results from the sharing of electrons between atoms.
Non-polar covalent bonding occurs between identical atoms, like H2, in which the
electrons are shared equally.
Types of Chemical Bonds:
Polar Covalent Bonding
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Polar covalent bonding occurs when atoms are not different enough
to transfer electrons (ionic) but at the same time are not identical
and cannot share electrons equally (covalent); instead there is an
unequal sharing of electrons.
Example: HF, which has an unequal charge distribution or a partial
charge. Electrons are more attracted to, or want to spend more
time with, fluorine than hydrogen. This is what we call bond
polarity.
The partial charge is represented with a delta, δ, and a positive or
negative sign depending on where electrons are positioned.
Bond Polarity & Dipole Moments
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A dipole moment indicates the direction of
charge distribution between two atoms. Can be
represented by a center of positive charge and a
center of negative charge.
The dipole character of a molecule is often
represented by an arrow, which points toward
the negative center and its tail indicated the
positive center.
Any diatomic molecule that has a polar bond has
a dipole moment.
Bond Polarity
Ionic and covalent bonding are the extremes of types of bonding.
Ionic > Polar Covalent > Non-Polar Covalent
NaCl
Cl2
HCl
Electronegativity
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Electronegativity is the relative ability of an atom in a molecule to attract shared electrons
to itself.
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Trend: Electronegativity increases across the period and decreases down the group.
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Fluorine is the most electronegative atom.
Noble Gasses have NO electronegativity!
Electronegativity & Polarity
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The polarity of a bond depends on the difference between the
electronegativity values of the atoms forming the bond.
Difference of Zero
Non-polar Covalent Bond
Intermediate Difference
(greater than 1.0 to 2.0)
Polar Covalent Bond
Large Difference
(greater than 2.0)
Ionic Bond
Lewis Symbols (Electron Dot Symbols)
• In a Lewis symbol, the chemical symbol for the element
represents the nucleus and core electrons of the atom.
• Dots around the symbol represent the valence electrons.
• In writing Lewis symbols, the first four dots are placed singly on
each of the four sides of the chemical symbol.
• Dots are paired as the next four are added.
• Lewis symbols are used primarily for those elements that
acquire noble-gas configurations when they form bonds.
Lewis Structures of Simple Molecules
• A Lewis structure is a combination of Lewis
symbols that represents the formation of
covalent bonds between atoms.
• Lewis structure shows the bonded atoms with
the electron configuration of a noble gas; that is,
the atoms obey the octet rule.
Lewis Structures
• The shared pairs of electrons in a molecule are called bonding pairs.
• In common practice, the bonding pair is represented by a dash (—).
• The other electron pairs, which are not shared, are called nonbonding
pairs, or lone pairs.
Each chlorine atom sees
an octet of electrons.
Multiple Covalent Bonds
• The covalent bond in which one pair of electrons is shared is
called a single bond.
• Multiple bonds can also form:
In a double bond two pairs
of electrons are shared.
In a triple bond three pairs
of electrons are shared.
Note that each atom obeys the octet
rule, even with multiple bonds.
Writing Lewis Structures: A Method
1.
2.
3.
4.
5.
Add together the number of “Happy” electrons.
(H/He = 2e- , B = 6e-, all other atoms = 8e-)
Add up the number of valence electrons.
(Step #1 – Step #2) / 2 = # of Bonds
Put the least electronegative atom in the middle (called the central atom) and
attach all other atoms as terminal atoms.
# bonds = # attachments  all single bonds
# bonds > # attachments  multiple bonds (double, triple)
# bonds < # attachments  EXCEPTION to the octet rule
(attach all attachments with single bonds)
Count electrons to make sure all atoms are “happy”.
(H/He = 2e- , B = 6e-, all other atoms = 8e-)
If short add the correct number of lone pairs to make each atom “Happy”
**If the compound is an exception to the octet rule, make attachments
happy and place extra valence electrons in pairs on the central atom –
may hold more than 8 electrons.
Example
Write a plausible Lewis structure for nitrogen trifloride, NF3.
Example
Write a plausible Lewis structure for carbonate ion, C032-.
Example
Write the Lewis structure for boron tribromide, BBr3.
Exceptions to the Octet Rule
<Expanded Octets>
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1.
2.
3.
4.
Expanded octets (more than 8 electrons) can
only occur in elements that have d-orbitals
(periods 3-7).
Do the math  # bonds < # attachments
Add all attachments -- single bonds and lone
pairs to make happy.
Count number of electrons in structure and
compare with valence electrons (step 2).
If short valence electrons, add them in pairs to
the central atom only.
Example
Write the Lewis structure for iodine pentafluoride, IF5.
Resonance: Delocalized Bonding
• When a molecule or ion can be represented by two or more plausible
Lewis structures that differ only in the distribution of electrons, the
true structure is a composite, or hybrid, of them.
• The different plausible structures are called resonance structures.
• The actual molecule or ion that is a hybrid of the resonance structures
is called a resonance hybrid.
• Electrons that are part of the resonance hybrid are spread out over
several atoms and are referred to as being delocalized.
Three pairs of
electrons are
distributed among
two bonds.
Example
Write three equivalent Lewis structures for the SeO3 molecule
that conform to the octet rule, and describe how the
resonance hybrid is related to the three structures.
Molecular Geometry
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Molecular geometry is
simply the shape of a
A carbon
dioxide
molecule is
linear.
molecule.
Molecular geometry is
described by the geometric
figure formed when the
atomic nuclei are joined by
(imaginary) straight lines.
Molecular geometry is found
using the Lewis structure,
but the Lewis structure itself
does NOT necessarily
represent the molecule’s A water
molecule is
shape.
angular or
bent.
VSEPR
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Valence-Shell Electron-Pair Repulsion (VSEPR) is a simple method
for determining geometry.
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Basis: pairs of valence electrons in bonded atoms repel one another.
These mutual repulsions push electron pairs as far from one another as
possible.
B
B
A
B
B
B
A
B
When the electron
pairs (bonds) are as
far apart as they can
get, what will be the
B-A-B angle?
Two Attachments
Linear
Three Attachments
Trigonal Planar
Four Attachments
Tetrahedral
Five Attachments
Trigonal Bipyramidal
Six Attachments
Octahedral
Example
Use the VSEPR method to predict the shape of the nitrate ion,
NO3-.
Example
Use the VSEPR method to predict the shape of the SF4
Molecular Polarity
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Although most bonds within a molecule are
polar, the arrangement of these bonds may
cause the overall molecule to be non-polar.
IDENTICAL dipoles (same elements b/w bonds)
can cancel at specific angles or combination of
the following.
2 @ 180o
3 @ 120o
4 @ 109.5o
HINT: if the molecule contains less than (<)
bond angles, the molecule is POLAR!!!!
Atomic Orbital Overlap (Hybridization)
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Valence Bond (VB) theory states
that a covalent bond is formed when
atomic orbitals (AOs) overlap.
In the overlap region, electrons with
opposing spins produce a high
electron charge density.
• In general, the more extensive
the overlap between two orbitals,
the stronger is the bond between
Overlap region
two atoms.
between nuclei has
high electron density
Example: Bonding in H2S
The measured bond
angle in H2S is 92°;
good agreement.
The hydrogen atoms’ s orbitals
can overlap with the two halffilled p orbitals on sulfur.
Honors: Test Topics
1. Molecular Geometry:
- shapes
- bond angles
- non(un)bonded (lone) pairs & bonded pairs (attachments)
- polarity
** bent also called angular, seesaw also called sawhorse
2. Electronegativity Trend
- most polar=largest difference
3. Covalent Compounds (2 nonmetals)
- examples
- polar covalent vs. nonpolar covalent
4. Ionic Compounds (metal & nonmetal- opposite charges)
- examples
5. # of valence electrons
6. # of resonance forms
7. Single Bonds (2 e-), Double Bonds (4e-), Triple Bonds (6e-)
# of bonds vs. bond strength
Chemical Bonding
Chemical Bonds


A ______ is a force that holds groups
of two or more atoms together and
makes them function as a unit.
__________ is the amount of energy
required to break a particular bond.
*determines the strength of a bond
Bonding Review



What type of elements form ionic bonds?
What type of elements form covalent
bonds?
Where are these types of elements found
on the periodic table in relationship to
each other?
Chemical Bonding: An Overview
• _____________ electrons play a fundamental role in
chemical bonding.
• When metals and nonmetals combine, valence
electrons usually are __________ from the metal to
the nonmetal atoms, giving rise to ______________.
• In combinations involving only nonmetals, one or
more ____ of valence electrons are ______ between
the bonded atoms, producing _________________.
• In losing, gaining, or sharing electrons to form
chemical bonds, atoms tend to acquire the electron
configurations of noble gases.
Types of Chemical Bonds:
Ionic Bonding



____________ occurs when there is an attraction
between oppositely charges ions. Electrons are
transferred to form oppositely charged ions.
______________ result when an atom that loses
electrons relatively easily (a metal) reacts with an atom
that has a high affinity for electrons (a non-metal).
Example: Sodium chloride, NaCl
Na+ & Cl- ions react to form solid
sodium chloride.
Types of Chemical Bonds:
Covalent Bonding
_____________: A bond that results from the sharing of electrons between atoms.
_______________bonding occurs between identical atoms, like H2, in which the
electrons are shared equally.
Types of Chemical Bonds:
Polar Covalent Bonding



_________________occurs when atoms are not different enough to
transfer electrons (ionic) but at the same time are not identical and
cannot share electrons equally (covalent); instead there is an
___________sharing of electrons.
Example: HF, which has an unequal charge distribution or a ______
_______. Electrons are more attracted to, or want to spend more
time with, fluorine than hydrogen. This is what we call ______
________.
The partial charge is represented with a delta, δ, and a positive or
negative sign depending on where electrons are positioned.
Bond Polarity & Dipole Moments



A ____________indicates the direction of charge
distribution between two atoms. Can be
represented by a center of positive charge and a
center of negative charge.
The dipole character of a molecule is often
represented by an arrow, which points toward
the __________ center and its tail indicated the
positive center.
Any diatomic molecule that has a polar bond has
a dipole moment.
Bond Polarity
Ionic and covalent bonding are the extremes of types of bonding.
Ionic > Polar Covalent > Non-Polar Covalent
NaCl
Cl2
HCl
Electronegativity

_________________ is the relative ability of an atom in a molecule to attract shared
electrons to itself.

Trend: ___________________________________________________________________


___________ is the most electronegative atom.
_______________ have NO electronegativity!
Electronegativity & Polarity

The polarity of a bond depends on the difference between the
electronegativity values of the atoms forming the bond.
Difference of Zero
Non-polar Covalent Bond
Intermediate Difference
(greater than zero to 2.0)
Polar Covalent Bond
Large Difference
(greater than 2.0)
Ionic Bond