Chapter 6: Chemical Bonding and
Molecular Structure
Week 2 – Bonding, Lewis structures and
shapes of molecules
6.1 Fundamentals of Bonding
Fundamentals of bonding
There are three types of interactions within a molecule:
• Electrons and nuclei attract one another
• Electrons repel each other
• Nuclei repel each other
The hydrogen molecule H2
Fundamentals of bonding
• These three interactions are balanced to give the molecule its
greatest stability
• This balance is achieved when the electron density is situated
between the nuclei of bonded atoms
• This shared electron density is called a covalent bond
• Attractive energy between nuclei and electrons overcomes
repulsions from nuclei-nuclei and electron-electron interactions
• Covalent bond = “chemical bond in which two atoms share one
or more pairs of electrons
Fundamentals of bonding
Bond length and bond energy
• Bond length is the distance at which the molecule has the
maximum
• energetic advantage over the separated atoms
• Bond energy is the energy required to break the bond (kJ
mol-1), it is
• always positive
• Each different chemical bond has a characteristic bond
length and energy
Bond length
• >300 pm apart = no interaction
• 74 pm = maximum stability
• Closer than that has too much repulsion!
• Bond length = separation distance at which the
molecule has the maximum energetic advantage over
the separated atoms
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Bond energy
• Bond energy = the amount of energy required to break
the bond
• Measured in kJ mol-1
• Multiply the energy of one bond (7.22 × 10-19 J) by
Avogadro’s constant (6.022 × 1023 mol-1)
• Bond energy of H2 = 435 kJ mol-1
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Covalent
Bonds
• Covalent bond—A bond formed by sharing electrons
between atoms
• Molecule—A group of atoms held together by
covalent bonds
Covalent Bonds
• Main group elements undergo reactions that leave
them with eight valence electrons (or two for
hydrogen), so that they have a noble gas electron
configuration.
• Nonmetals can achieve an electron octet by sharing
an appropriate number of electrons in covalent
bonds.
Covalent Bonds
Covalent bonding in hydrogen (H2):
• Spherical 1s orbitals overlap to give an egg-shaped region.
• Two electrons between the nuclei, providing 1s2
configuration of helium.
• H-H, H:H and H2 all represent a hydrogen molecule.
Multiple Covalent Bonds
• Single bond—A covalent bond formed by sharing one
electron pair.
• Represented by a single line:
H-H
• Double bond—A covalent bond formed by sharing two
electron pairs.
• Represented by a double line:
O=O
• Triple bond—A covalent bond formed by sharing three
electron pairs.
• Represented by a triple line:
N≡N
POLARITY (Unequal electron sharing)
• H2 and F2 share electrons equally
• What about HF?
• Effective nuclear charge of F greater than that
of H (F has greater pull for electrons)
• Known as a polar covalent bond
• “a bond that possesses an asymmetric
distribution of electrons”
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Partial charges (δ+ and δ-)
• Remember how electrons have a negative
charge?
• Fluorine has a partial negative charge (δ-)
• Hydrogen has a partial positive charge (δ+)
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Partial charges (δ+ and δ-)
• We signify that this molecule has a variation in
electron density by the arrow you see under
the diagram
• Arrow points from δ+ to δ• We make the δ+ end into a plus sign!
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Fundamentals of bonding
Unequal electron sharing
Electronegativity gives a numerical value of
how strongly an atom attracts the electrons
in a chemical bond
Trend in the periodic table
Electronegativity
• Atoms of each element have a characteristic
ability to attract electrons
• Electronegativity = “A measure of the ability of
an atom in a molecule to attract the shared
electrons in a chemical bond”
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Electronegativity trends
• Metals = low electronegativities
• Nonmetals = high electronegativities
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Electronegativity = χ (chi)
• Polar bonds = bonds between atoms of different
electronegativities
• Δχ= χ (F) –χ(H) = 4.0 - 2.1 = 1.9
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Polar Covalent Bonds and Electronegativity
• Electronegativity differences of less than 0.5 result in
nonpolar covalent bonds.
• Differences up to 1.9 indicate increasingly polar
covalent bonds.
• Differences of 2 or more indicate substantially ionic
bonds.
• There is no dividing line between covalent and ionic
bonds; most bonds fall between these categories.
Polar Molecules
• Molecular polarity is due to individual bond polarities
and lone-pair contributions. Electrons are displaced
toward the more electronegative atom.
Polar Molecules
• Molecular polarity depends on the shape of the
molecule.
• Symmetrical molecules can have polar bonds and be
non-polar overall.
Polar Molecules
• Polarity has a dramatic effect on the physical properties of
molecules particularly on melting points, boiling points, and
solubilities.
6.5 Properties of Covalent Bonds
Properties of covalent bonds
• Dipole moment:
• Most chemical bonds are polar (one end slightly positive, the other slightly
negative).
• Bond polarities can lead to molecules with dipole moment.
• Dipole moment depend on bond polarities (Δχ) and on molecular shape.
Properties of covalent bonds
Properties of covalent bonds
• Bond length:
• Bond length of a covalent bond is the nuclear separation distance at which
the molecule is most stable.
• At this distance, attractive interactions are maximised relative to repulsive
interactions.
• Bond lengths vary between 70 and 250 pm (1pm = 10-12m).
Properties of covalent bonds
• Bond energy:
•
•
•
•
It is the amount of energy that must be supplied to break a chemical bond.
Bond energies increase as more electrons are shared between the atoms.
Bond energies increase as the Δχ between bonded atoms increases.
Bond energies decrease as bonds become longer.
6.2 Ionic Bonding
Ionic bonding
• Bonds formed between atoms of very different
electronegativities are predominantly ionic in
character
• Cations (+ve) from Groups 1 & 2
• Anions (-ve) from Groups 16 & 17
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Ionic bonding
• Compounds formed between elements of very different
electronegativities are ionic
• Most ionic compounds are solids with high melting points
• They are held together by the attractive forces between oppositely
charged ions
Ionic compounds don’t share electrons in a
bond (NOT SIGMA BONDS)
• Form a lattice
Lattice energies
• Amount of energy required (in kJ mol-1) to
break the lattice apart
• Depend on the charge and size of the ions
• Increase cation size, decrease lattice energy
• Increase anion size, decrease lattice energy
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6.3 Lewis Structures
4.6 Molecular Formulas and Lewis Structures
• Lewis structure—A molecular representation that
shows both the connections among atoms and the
locations of lone-pair valence electrons
• Lone pair—A pair of electrons that is not used for
bonding
4.6 Molecular Formulas and Lewis Structures
Drawing Lewis Structures
• Lewis Structures for Molecules Containing C, N, O, X
(Halogen), and H
• C forms four covalent bonds and often bonds to other
carbon atoms.
• N forms three covalent bonds and has one lone pair of
electrons.
• O forms two covalent bonds and has two lone pairs of
electrons.
• Halogens (X) form one covalent bond and have three lone
pairs of electrons.
• H forms one covalent bond.
Lewis Structures
• Show how compounds are bonded
• Differentiate between bonding and nonbonding
electrons
• The first step in developing a bonding
description of a molecule
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Lewis structures
The conventions
• An atom is represented by its elemental symbol
• Only the valence electrons appear
• A line represents one pair of electrons that is shared
between two atoms (double bond : 2 lines, triple
bond : 3 lines…)
• Dots represent the nonbonding electrons on that
atom (nonbonding pairs are called lone pairs)
Lewis structures
The conventions
Lewis structures (5 Step Procedure)
• Step 1: Count the valence electrons
• Step 2: Assemble the bonding framework using
single bonds
• Step 3: Place three nonbonding electron pairs
on each outer atom except H
• Step 4: Assign the remaining valence electrons
to inner atoms
• Step 5: Minimise formal charges on all atoms
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EXAMPLE = SO2
• Step 1: Count the valence electrons. If the
species is an ion, add or subtract one electron
for each negative or positive charge
respectively
• S=6
• O=6
• SO2= [6+(2×6)] = 18 valence electrons
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Electronegativity trends
• Metals = low electronegativities
• Nonmetals = high electronegativities
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EXAMPLE = SO2
• Step 2: Assemble the bonding framework using
single bonds
• Outer atoms are usually the more electronegative
• Usually one inner atom attached to two or more
other atoms
• Is sulfur or oxygen more electronegative?
4 electrons used
O
S
O
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EXAMPLE = SO2
• Step 3: Place three nonbonding pairs of
electrons on each outer atom except H
• Each outer atom (except H) is associated with
eight electrons (four pairs of electrons)
• Octet rule
• Nonbonding pairs are called lone pairs
16 electrons used
O
S
O
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EXAMPLE = SO2
• Step 4: Assign the remaining valence electrons
to inner atoms
• We’ve ‘used’ 16 electrons, and Step 1 told us
there is 18 electrons in this molecule
• 2 electrons left
18 electrons used
O
S
O
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EXAMPLE = SO2
• Step 5: Minimise formal charges on all atoms
• Making sure our structure makes sense!
Formal charge  (valence electrons of free atom) - (electrons assigned in Lewis structure)
• Lone pair electrons ‘belong’ to the atom
• Electrons in bonds are shared between the atoms
O
S
O
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EXAMPLE = SO2
O
S
O
Formal charge  (valence electrons of free atom) - (electrons assigned in Lewis structure)
Valence electrons of atom
S
O
6
Electrons assigned in Lewis structure 4
Formal charge
(6-4) = +2
Valence electrons of atom
6
Electrons assigned in Lewis structure 7
Formal charge
(6-7) = -1
• We can minimise these formal charges to give
a better structure!
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EXAMPLE SO2
• We can minimise these formal charges!
• Convert one lone pair from each O into a bond
O
S
O
Valence electrons of atom
S
O
S
O
O
6
Electrons assigned in Lewis structure 6
Formal charge
(6-6) = 0
Valence electrons of atom
6
Electrons assigned in Lewis structure 6
Formal charge
(6-6) = 0
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More about Formal Charges
• Formal charges can’t always equal zero
• Make sure negative formal charges are on the
more electronegative atoms
• Formal charges are NOT THE SAME as partial
charges (δ+ or δ-)
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Resonance Structures
• Sometimes there’s more than one Lewis
structure possible
• Resonance structures
• Eg. NO3• 24 valence electrons
• Three N-O bonds use 6 electrons
• Three lone pairs on each O uses another 18
• Formal charges: O = -1, N = +2
• We can minimise these formal charges
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Resonance Structures
• Move one lone pair from O to form a bond
• We can do that three times!
• But which one is right?
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Resonance Structures
• NONE- no single Resonance structure is an accurate
representation of NO3• Experiments show each N-O bond is of equal length
(but double bonds are SHORTER than single bonds!)
• Double-headed arrow emphasises that a complete
depiction requires ALL the resonance structures
• Electrons don’t ‘flip back and forth’
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Lewis structures
Resonance structures
• Composites of equivalent Lewis structures
• Resonance structures differ only in the
position of the electrons, not atoms
CO and NO: Pollutants or Miracle Molecules?
• CO is a poison and, NO, a pollutant in the
environment.
• In 1992, it was also discovered that these
molecules are key chemical messengers in
the body.
• CO and NO are highly soluble and can
diffuse from one cell to another, where
they stimulate production of guanylyl
cyclase.
• Guanylyl cyclase controls the production of
cyclic GMP, which regulates many cellular
functions.
CO and NO: Pollutants or Miracle Molecules?
• CO is associated with long-term
memory. When CO production is
blocked, long-term memories are no
longer stored, and memories that
previously existed are erased. When CO
production is stimulated, memories are
again laid down.
• NO fights infections and tumors,
transmits messages between nerve
cells and is associated with learning and
memory, sleeping, and depression. It is
also a vasodilator, a substance that
allows blood vessels to relax and dilate.
6.4 Valence-shell-electron-pair-repulsion
theory
Valence Shell Electron Pair Repulsion (VSEPR) theory
• Gives the 3D structure
• VSEPR determines the shape of a molecule based
on the repulsion between pairs of electrons
• Electrons around an inner atom within a molecule
will be situated as far apart as possible in the
preferred 3D structure
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VSEPR Procedure
• 1. Draw the Lewis structure of the molecule
• 2. Count the number of sets* of bonding pairs and
lone pairs of electrons around the inner atom
Number of sets of electron pairs
Geometry of sets of electron pairs
2
Linear
3
Trigonal planar
4
Tetrahedral
5
Trigonal bipyramidal
6
Octahedral
* VSEPR doesn’t distinguish between single, double
and triple bonds- each is treated as one set
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VSEPR Procedure
• 3. Modify the geometry to account for the
magnitudes of repulsions between sets of
electron pairs, depending on if they are
bonding (BP) or lone pairs (LP)
• Repulsions in the order
LP-LP > BP-LP > BP-BP
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Two sets: Linear geometry
• Beryllium hydride (BeH2)
H Be H
• Carbon dioxide (CO2)
O
C
O
• Both Linear, 180° between each atom
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Three sets: Trigonal planar geometry
• Boron trifluoride (BF3)
F
B
120° between each bond
F
F
• Nitrite ion (NO2-)
N
N
O
O
O
O
Due to lone pair has ‘bent’ molecular shape
115° between each bond
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Four sets: Tetrahedral geometry
• Methane (CH4) has tetrahedral molecular shape
H
H
C
H
109.5° between each bond
H
• Ammonia (NH3) has trigonal pyramidal shape
H
N
H
107° between each bond
H
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Four sets: Tetrahedral geometry
• Water (H2O) has bent molecular shape
O
H
H
104.5° between each bond
• Tetrahedral geometry has three shapes
• Dependant on number of lone pairs
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Five sets: Trigonal bipyramidal
• Phosphorous pentachloride (PCl5) has trigonal bipyramidal
geometry
Cl
Cl P Cl
Cl Cl
• Sulfur tetrafluoride (SF4) has two possibilities- trigonal
pyramid or seesaw molecular shape- but seesaw is more
stable
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Five sets: Trigonal bipyramidal
• Chlorine trifluoride (CF3) has T-shaped molecular
shape
F
Cl
F
F
• Triiodide ion (I3-) has linear shape
I I I
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Five sets: Trigonal bipyramidal
• Trigonal bipyramidal geometry has four shapes
Depend on number of lone pairs
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Six sets: Octahedral geometry
• Sulfur hexafluoride (SF6) has octahedral shape
• Replacement of any atom is equivalent
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Six sets: Octahedral geometry
• Chlorine pentafluoride (ClF5) has a square
pyramidal shape
F
F
F
F Cl
F
• Xenon tetrafluoride (XeF4) has square planar
molecular shape
F
F
F Xe
F
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Six sets: Octahedral geometry
• Three possible molecular shapes for octahedral
geometry
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VSEPR Conclusion
• Draw the Lewis structure
• Count the number of sets of electron pairs,
NOT the number of atoms!
• Eg. SF4 is NOT tetrahedral like CH4
it’s based on trigonal bipyramidal geometry!
• Base the geometry on the repulsions
Repulsions are in the order
LP-LP > BP-LP > BP-BP
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