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MED109 - Lecture 02a - Chemical Bond Intro. Covalent bonds

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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
7
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
8
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
9
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)
10
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)
11
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)
12
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
13
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
14
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
15
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
16
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
17
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
18
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)
19
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•
••
20
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
21
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 •
•
22
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:
••
•
••
23
Lewis dot structures
for molecules
example 3:
N2
O=O
••
••
N≡N
••
O2
••
example 2:
•• ••
24
How to draw Lewis structures
https://www.youtube.com/watch?v=1ZlnzyHahvo
25
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
26
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
27
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
28
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
29
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
30
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
31
Detailed Formation of
a Double or Triple Bond
Example: double bond between 2 oxygen atoms
••
••
••
••
O=O
32
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
33
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
34
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
35
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
37
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
38
Detailed Formation of
a Double or Triple Bond
••
••
N≡N
39
Hybrid Orbitals explained (video)
https://www.youtube.com/watch?v=vHXViZTxLXo (12 min)
40
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/
41
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