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Organic chemistry
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A summary of the lecture on organic chemistry - describing the various forms of carbon bonding, their properties and orbitals
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How can carbon form 4 covalent Molecular orbital theoryOrganic chemistry – the chemistry of
bonds? In H2 – s orbital single bond
compounds of carbon (excluding
Paired orbitals cannot form 2 singly occupied s orbitals on each H
those traditionally dealt with in
covalent bonds. And for orbitals to combine to form two molecular
inorganic chemistry e.g. CO, CO2 -
mix, they must have the same orbitals. (note that the number of
usually 1 carbon molecules)
symmetry. orbitals will always stay the same).
Examples: methane, penicillin,
Orbital hybridisation- the mixing of You get a sigma anti-bonding orbital
buckminsterfullerene (allotrope) and
orbitals on an atom. It’s a and a sigma bonding orbital (lower
DNA.
mathematical device (used to energy, more favourable).
Organic compounds can come in
explain observations), atoms don’t Two half-filled atomic orbitals have
many physical states and structures,
actually do this. combined to form a single/sigma –
which underpins their properties
Carbon can form 4 covalent bonds bond, consisting of 2 molecular
and hence the chemistry of life.
as the two 2p2 and 2s2 electrons orbitals. Most likely to find the
combine to form four 2sp3 hybrid electrons between the 2 nuclei.
Carbon orbitals -> energetically
Can bond strongly to other elements favourable.
and itself. It can also catenate (form Carbon-carbon single bonds
long chains of carbon atoms) e.g in ethane, we have three sp3
C-C = 346kJ/mol orbitals mix with an s orbital of
C=C = 598 KJ/mol three H atoms to form 3σ bonds.
C≡C = 835 KJ/mol The remaining sp3 orbital then
Electronic configuration: 1s2 , 2s2, mixed with the sp3 orbital from a Double bond between p orbitals
2p2. It therefore requires 4 second carbon atom giving a 1st pair of electrons (in p orbitals) –
electrons to be stable (octet rule) carbon-carbon σ bond. Look like two px orbitals overlap end-on to
the 2px overlap shown below. form two sigma orbitals.
Carbon-carbon triple bonds Gives a tetrahedral shape!
Simplest example is ethyne, C2H2
Linear molecule (180°). One p Organic
orbital is hybridised with an s
Carbon-carbon double bonds
orbital. This give us two sp orbitals chemistry
e.g in ethene
and two unhybridized p orbitals.
each carbon bonds to 2 H atoms via
The sp orbitals are linear and at
σ bonds and double bonds to
right angles to the unhybridized
another carbon atom. The molecule
orbitals. 2nd pair of electrons in p orbitals – two
is planar (107°) and trigonal.
E.g in ethyne, 1 sp orbital mixes py orbitals overlap side-on to form
Two p orbitals mix with the s
with the 1s orbital of a H atom, the
orbital, giving 3 sp2 orbitals and one two π orbitals
other sp orbital forms a σ bond, A double bond consists of a σ bond
unhybridized orbital.
and the two unhybridised p orbitals
Sp2 orbitals point away from each and a π bond
on each carbon overlap. This forms
other and the unhybridised orbital
a triple bond between the two
stick up, above and below the plane
carbons (1 σ bond + 2 π bonds).
of the molecule.
1st π bond = 252KJ/mol
In ethene, each C has σ bonds to 2 H
2nd π bond = 237 KJ/mol
atoms, and one σ bond to the
second C. There is also then a π
Orbital hybridisation is important
bond between the C and the second
in telling us the shape of the
C through the overlapping of the
molecule. This is vital to its
unhybridized p orbitals.
function e.g ethene is used in the
Double bond: one π bond and one Dot & cross/Lewis structure
commercial ripening of fruit, but
σ bond. (simplest & most limited theory)
ethyne and ethane don’t fit the
receptor. Shows arrangement of atoms but
In DNA the bases contain does explain how the bonds form,
alternating double bonds -> planar how electrons are shared between
so they can act as rungs atoms or the geometry of the
On the backbone – single bonds molecules.
give a tetrahedral arrangement.
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