MED109 Chemistry
Lecture 02a – Chemical Bonding
Introduction & Covalent Bonds
Dr Olivier Missa
Staff Room, Dept. of Medicine
OMissa@dwu.ac.pg
phone ext: 847
1
Lesson 2a – Learning Objectives
• Draw the Lewis dot symbol for an element
• Explain the formation of atomic bonds
through the principle of the octet rule
• Describe how different types of covalent
bonds are formed
• Illustrate the formation of molecular orbitals
• Draw the Lewis dot structures of a molecule
2
Chemical Bonding
The tendency of atoms to react with other
atoms leads to the formation of chemical
bonds
A chemical bond between atoms involves
their valence electrons & can be explained
through different interactions between the
atoms:
Ø covalent
bond (sharing electrons)
Ø ionic bond (transfer of electrons):
electrostatic
3
Lewis dot symbols
The only electrons involved in bonding
are the valence electrons & can be
represented by the Lewis dot symbols
These are represented by the symbol of an
element & one dot for each valence
electron
4
Lewis dot symbols
•
The electron configuration shows all
electrons; Lewis dot symbols only
show valence electrons:
example:
N
1s2 2s22p3
core
electrons
5 valence
electrons
••
•N•
•
5
Lewis dot symbols
Problem: Which of the following is an incorrect
Lewis dot symbol?
•
Ca
•
•
•C•
•
2A
••
• O•
••
••
1A
••
Ne
••
••
•
H
••
• O•
•
8A
3A 4A 5A 6A 7A
6
Lewis dot structure
for Oxygen
https://www.youtube.com/watch?v=o4J9mLbvWJA
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Chemical Bond
A chemical bond is formed because the bound
atoms reach a more stable state (lower energy
compared to unbound atoms)
Noble gases (He, Ne, Ar, …) have a very low
tendency to form chemical bonds (maximum
number of electrons in their outermost shell =
filled valence shell)
outermost
electronic
configuration
He: 1s2
Ne: 2s22p6
Ar: 3s23p6
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Octet rule
Noble gases have 8 electrons (only 2 for He) in the
outermost shell; this configuration gives them great
stability (no reactivity, no chemical bond)
When forming chemical bonds, elements try to reach
(mimick) the configuration of noble gases
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Octet rule
Octet rule: when forming a chemical
bond, an element loses, gains, or
shares its electron(s), in order to have
8 electrons in its outermost shell (or
only 2 electrons for elements near He)
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Octet rule
Examples
H (1s1): shares 1 e- (He)
Na (3s1): loses 1 e- (Ne)
Ca (4s2): loses 2 e- (Ar)
N (2s22p3): shares 3 e- (Ne)
S (3s23p4): gains 2 e- (Ar)
C (2s22p2): shares 4 e- (Ne)
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Covalent bond
• Occurs between two nonmetal atoms sharing their
electrons
• Covalent compounds are also known as molecules
• In the valence bonding theory the driving force for covalent
bond formation derives from the octet rule (both atoms
reaching the stable configuration of noble
gases)
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Covalent bond
• If the bound atoms are identical
the same electronegativity *) the
called non-polar #
(or have
covalent bond is
• In case of different atoms (having
different electronegativity) the covalent bond
is called polar #
Ø * Lecture 2b
Ø # Lecture 2c
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How is a covalent
bond formed ?
Example: binding between two H atoms
(1s1: each have one unpaired electron)
Ø at a proper distance between the atoms
(inter-nuclear bonding distance), each
bonding electron is attracted by the
nucleus of the other atom
➡
H• ➡
•H
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How is a covalent
bond formed ?
Ø When the two H atoms get close to each
other, their atomic orbitals with unpaired
electrons interact, leading to the
formation of a covalent bond
H• •H
H➖H
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Covalent bond
strength & length
Attraction of bonding electrons by both nuclei
determines the bond strength
Strong repulsive forces keep the positive charge
of nuclei apart after they have reached a distance
corresponding to the minimum energy
This distance between atoms at
the point of minimum energy
corresponds to the bond length
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Single, double
or triple bonds
A covalent bond between two atoms may be:
Ø single (1 shared pair of electrons)
Ø double (2 shared pairs of electrons)
Ø triple (3 shared pairs of electrons)
In all cases the two elements involved (single,
double or triple bond) reach a stable
configuration, according to the octet rule
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Single, double
or triple bonds
Instead of representing a covalent bond as two dots, it
is represented by a line connecting the two atoms
The single line corresponds to a single bond with a
shared pair of electrons
ex: H➖H
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Single, double
or triple bonds
Examples:
H• + •H → H• •H
H➖H (single bond)
H• + •F → H• •F
H➖F (single bond)
O: + :O → O: :O
O=O (double bond)
:C: + 2 :O → O::C::O
N⋮ + ⋮N → N⋮ ⋮N
O=C=O (2 double bonds)
N≡N (triple bond)
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Double bond formation using
the Lewis dot structure?
Formation of CO2:
Ø carbon:
4 valence electrons
Ø oxygen:
6 valence electrons
•
••
• O• • C •
••
•
•
• C•
•
••
• O•
••
••
• O•
••
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Double bond formation using
the Lewis dot structure?
1. Formation of a single bond:
only 7 electrons
(one shared with C)
•• • ••
• O • • C• • O •
•• • ••
only 7 electrons
(one shared with C)
only 6 electrons
(one shared with one O,
one shared with the other O)
Ø octet rule not respected
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Double bond formation using
the Lewis dot structure?
2. Create new bonds: double bonds
C
••
O•
•
•• ••
•• ••
•
•O
••
8 electrons
around each atom
Lewis dot
structure for CO2
••
•
•
• •O = C = O •
•
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Lewis dot structures
for molecules
In the Lewis structure of a molecule, all valence electrons
are shown, including the ones involved in the bond (2) & the
others also referred as lone pairs (non-bonding pairs)
Cl2
••
••
Cl➖Cl
••
••
example 1:
••
•
••
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Lewis dot structures
for molecules
example 3:
N2
O=O
••
••
N≡N
••
O2
••
example 2:
•• ••
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How to draw Lewis structures
https://www.youtube.com/watch?v=1ZlnzyHahvo
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Formation of a
Molecular Orbital
When two atoms get close, a new space
accommodating one electron from each atom
is formed through a combination of their
atomic orbitals
A new orbital named molecular orbital is
formed that surrounds the molecule &
corresponds to the area where the shared
electrons are more likely to be found
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Formation of a
Molecular Orbital
In a molecular orbital, no more than two
electrons are accommodated (each with
opposite spins like in atomic orbitals); the
bonding electrons are mainly located at a
higher density between the two nuclei
The bonding interaction is achieved because
the energy of the molecular orbital is lower
(more favorable) than the energy of the atomic
orbital
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Formation of a
Molecular Orbital
Molecular orbitals that are symmetrical about
the axis of the bond (head to head overlap of
atomic orbitals) are called sigma (σ)
molecular orbitals
atomic
orbital 1s
H•
+
atomic
orbital 1s
σ molecular
•H
H• •H
orbital
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Formation of a
Molecular Orbital
Molecular orbital theory actually predicts that the
combination of two atomic orbitals leads to the
formation of two molecular orbitals named
bonding molecular orbital & anti-bonding
molecular orbital
Ø bonding molecular orbitals are at lower energy than
the atomic orbitals they were formed from
Ø antibonding molecular orbitals are at higher energy
than the atomic orbitals from which they were formed
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Formation of a
Molecular Orbital
As a results the electrons participating in the bond
will only occupy the bonding molecular orbital
(here σ) & the antibonding molecular orbital
(here σ*) will remain empty
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Other Types of σ
Molecular Orbitals
σ molecular orbitals can be formed also between
s & p atomic orbital as well as between p atomic
orbitals
atomic
orbital 1s
atomic orbital 2p
atomic orbital 2p
atomic orbital 2p
σ molecular orbital
σ molecular orbital
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Detailed Formation of
a Double or Triple Bond
Example: double bond between 2 oxygen atoms
••
••
••
••
O=O
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Detailed Formation of
a Double or Triple Bond
The first molecular orbital σ forms through
the linear combination of 2p atomic orbitals
along the inter-nuclear axis (head to head
overlap of atomic orbitals)
atomic orbital 2p
atomic orbital 2p
molecular orbital σ
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Detailed Formation of
a Double or Triple Bond
The second molecular orbital forms through
the lateral combination (sideways overlap)
of the remaining 2p orbitals
This molecular orbital is named π (pi)
atomic
orbital 2p
atomic
orbital 2p
molecular orbital π
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Detailed Formation of
a Double or Triple Bond
The maximum electron density in the π
orbital is above & below the inter-nuclear
axis
Attention: there is only one orbital, although the
bonding electrons are accommodated in two
apparently unconnected regions
The π orbital occupies the empty space, not
used by the σ orbital
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Detailed Formation of
a Double or Triple Bond
Note
stability of the
π orbital
(energy required
to break π)
<
stability of the
σ orbital
<
(energy required
to break σ)
As a result, in case of bond breakage, the first
bond broken is the one corresponding to the π
orbital, because of its lower stability
36
Detailed Formation of
a Double or Triple Bond
Example:
triple bond between 2 nitrogen atoms
The first molecular orbital
formed is the σ orbital
(head to head overlap of
atomic orbitals)
••
••
N≡N
molecular orbital σ
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Detailed Formation of
a Double or Triple Bond
Then two π orbitals form through the lateral
combination of the remaining two 2p atomic orbitals
(sideways overlap)
v
v
one π above & below
the inter-nuclear axis
one π in front &
behind
the
internuclear axis
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Detailed Formation of
a Double or Triple Bond
••
••
N≡N
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Hybrid Orbitals explained (video)
https://www.youtube.com/watch?v=vHXViZTxLXo (12 min)
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Bibliographic & Online Resources
Nuclear Chemistry (2015) in Bettelheim. F. A., et al. Introduction to
general, organic and biochemistry (Chapter 9 - 11th ed.). US: Cengage
Learning.
— Glencoe Virtual Lab: How can you tell which elements form chemical
bonds?
http://www.glencoe.com/sites/common_assets/science/virtual_labs/E20/E
20.html
— General Chemistry Online – Glossary: Atoms, elements and ions
htmlhttp://antoine.frostburg.edu/chem/senese/101/atoms/glossary.shtml
• How to draw Lewis Structures : five easy steps:
https://www.youtube.com/watch?v=1ZlnzyHahvo
• UC Davis Chemwiki: Bonding and anti-bonding orbitals
http://chemwiki.ucdavis.edu/Core/Theoretical_Chemistry/Chemical_Bonding
/Pictorial_Molecular_Orbital_Theory/Bonding_and_antibonding_orbitals
—
• PBS Learning Media (Covalent bonding)
http://www.pbslearningmedia.org/asset/lsps07_int_covalentbond/
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