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Module: BIOM - 1007

Lecturer: Dr Bhambra

Date: 11/10/16

The Octet Rule, Electronegativity, Bonding and Molecular Orbitals
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The bonds in compounds are formed during the redistribution of valence electrons
 Protons and neutrons never participate in bonding
 The most stable atoms have full valence shells
 Only a few elements exist with full shells, the noble gases
 This is why they are highly unreactive

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A compound has a much more stable distribution of atoms compared to separate atomic components
 In oxygen for example, 2 electrons are shared from each oxygen atom’s valence shell to fill both, as
seen here:

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We
full s

only
consider
and p
orbitals
terms of a

in
full valence shell, because of this, only 8 electrons are needed for any given shell
 The Octet Rule states that an atom seeks to gain a full valence shell with 8 electrons
 Exceptions to this are H (1e-) and He (2e-)
o

A valence shell can be represented using Lewis Dot Symbols
 This is a series of dots that surround the chemical symbol
 The dots are placed on the 4 sides of the symbol to show the 4 orbitals (1s and 3p)
 When adding dots to show a bond, the free spaces are filled up first, then the space next to it
 The following is a diagram of some Lewis dot symbols:

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Covalent bonds are when one or more pairs of electrons are shared equally between atoms whereas
an ionic bond is when one or more electrons is transferred to another atom

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The type of bonding that occurs is dependent upon the electronegativity () of elements, this is
an indication of how strongly an atom of an element can attract an electron
experiences periodicity and increases as you move across a period, left to right
 If the difference in is less than 1
...
7 then the bond will be ionic
 This is because the pull from the strong electronegative is too much for the weak
electronegative and so the electron will be completely transferred
 The electron configuration changes in an ionic bond, an example is sodium
(0
...
0):

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Valence shells
may contain
orbitals with
unpaired electrons
or that are fully
occupied (Pauli
exclusion
principal)
deemed non These are
bonding
or lone pairs of
electron
 These electrons still do play a role in showing the structure of the molecule however

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There are instances that, during a covalent bond, both electrons come from the same atom, this is
termed a dative/coordinate covalent bond
 One example is an ammonium ion, Nitrogen contributes a complete valence pair, for
sharing with the H+ ion
...
This holds the 4 sub units together

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The sharing of valence electrons between
atoms occurs through an overlap of atomic orbitals containing these electrons: These are now known as
molecular orbitals

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

More electrons must be present in bonding than anti-bonding orbitals
 Non-bonding orbitals also exist but don’t directly affect the formation of covalent bonds,
they are occupied by non-bonding electrons
The following image is what occurs:

o

There are two types of molecular orbital:
 Sigma (σ)
 Pi (π)

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Both orbitals have a distinctive shape and are
involved in forming covalent bonds
 The image below shows the shape of sigma
and pi orbitals:



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The sharing of one pair of electrons occupy the sigma orbital
A double bond, two pairs of electrons between two atoms, is shared, one occupies a sigma
orbital and one occupy a pi orbital
Finally, a triple bond is shared: one occupies a sigma orbital and 2 pairs occupy 2 pi orbitals



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