Topic 14 Bonding (HL)
• Shapes of molecules and ions
• Hybridisation
• Delocalisation of electrons
14.1 Shapes of molecules and ions
• Valence Shell Electron Pair RepulsionVSEPR for 5- and 6-negatively charged centre
=> Shapes are based on trigonal bipyramid and
octahedron
Expanded valence shells
• Sometimes the octet rule doesn’t hold
• The atom have 8 or 10 electrons
• The PCl5 molecule has 5 bonding electron pairs -a
symmetrical trigonal bipyramidal shape.
• 5 negative centres!
Trigonal bipyramid- PCl5
• Two types of electron rich regions:
– Equatorial: 3 bonds with 120o between.
– Axial: 2 bonds with 180o between
• Equatorial to Axial: 90o.
http://www.chem.ufl.edu/~myers/chm2045/shapes.htm
Trigonal bipyramid- SF4 and ClF3
• Non-bonding orbitals always occupy equatorial positions
• SF4 Equatorial:
2 bonds with 104o (<120o),
Axial:
2 bonds with 177o (<180o)
• ClF3 Equatorial:
1 bond
Axial: 2 bonds with 87,5*2= 175o
(<180o)
http://www.chem.ufl.edu/~myers/chm2045/shapes.htm
Octahedron
• All positions are equal- 90o
between all positions- SF6
• If two non-bonding orbitals: they take place
opposite each other => plan square shapeXeF4
• Many of the compounds that are forming
trigonal bipyramids and octahedrons are
fluorides because only high electronegative
ions can increase the number of valence
electrons
• Fluoride is also quite small (bigger ions
doesn’t have space enough).
14.2 Hybridisation
• So far we have talked about s-p-d-f-orbitals.
They only exist in single atoms in the gaseous
state
• When atom binds to each other the orbitals
will change their shape; they will undergo a
Hybridisation
(mathematics: linear combination)
s -bonds and p -bonds
http://ibchem.com/IB/ibnotes/full/bon_htm/14.2.htm
The bonds between carbon atoms
Bond type
Bond
energy
(kJ/mol)
Bond
length
(pm)
Hybrid
orbitals
Ethane
C 2H 6
single
348
154
1s
Ethene
C 2H 4
double
612
134
1 s + 1p
Ethyne
C 2H 2
triple
837
120
1s + 2 p
http://www.chemguide.co.uk/basicorg/bonding/methane.html
Single bonds in ethane C2H6
• Hybridisation: One s-orbital and three p-orbitals
=> Four sp3-orbitals (tetrahedral shape)
• Two carbons with sp3-orbitals now bind 3
hydrogen s-orbitals, with s-bonds:
The last orbital is used to s -bond to the next
carbon:
Single- and double bonds in ethene C2H4
• Hybridisation: One s-orbital and two p-orbitals =>
Three sp2-orbitals (trigonal planar shape). One p-orbital
is left over (red)
• Two carbons with sp2-orbitals now bind 4 hydrogen sorbitals, with s-bonds:
The green sp2-orbital is used to s –bond, and the red porbital is used to p -bond to the next carbon:
Double bond, cont
• Consist of one s -bond and one p -bond
• The p -bonding to the next carbon is at a right
angle, 90o, to the next carbon
http://www.groveridgeconsulting.com/?page_id=546
Single- and double bonds in ethyne C2H2
• Hybridisation: One s-orbital and one p-orbital=> Twoo sporbitals (trigonal planar shape). Two p-orbitals is left over
(red)
• Two carbons with sp-orbitals now bind 2 hydrogen sorbitals, with s-bonds:
The green sp-orbital is used to s –bond, and the red p-orbitals
are used to p -bond to the next carbon:
Triple bond
• Consist of one s -bond and two p -bonds
• The sp-orbitals give a linear shape
• The two p -bonding to the next carbon is at a
right angle to the next carbon and at right
angle to each other
Molecular shape and types of
hybridisation
• The shape of the hybrids corresponds to the
structure given by VSEPR / Lewis structure.
=> Determine the hybridisation by studies
of the shape of the molecule.
• Ethane : Ethene : Ethyne
sp3 : sp2 : sp
• Ammonia: sp3
• Water: sp3
14.3 Delocalisation of electrons
• Electrons that are not located at a certain
atom (c.f. metallic bond)
• Or in a certain bond between two atoms
• Often gives rise to a stronger (shorter) bond
• The delocalised electrons absorb light in the
UV- or visible region
Benzene
• 6 sp2-hybridised carbons, 6 p-orbitals
• The p-orbitals can overlap both to the right and
to the left- a system of delocalised p-electrons
are formed. The electrons are said to be
delocalised, often shown as a circle
• Single bond 154 pm, double 134 pm, benzene 140 pm
Resonance
d
• a and b are resonance
structures of benzene
• c is a resonance hybridthe most stable form. By
delocalisation of the
electrons the molecule
gain resonance energy
• d resonance in pyridine
Why has phenol acidic properties?
Low pKa- strong acid (HCl -7)
High pKa (1-14)- weak acid
pKa > 14 no acid
Phenol pKa 8
Why has acetic acid acidic properties
but not ethanol?
C2H5OH + H2O
↔
CH3COOH + H2O ↔
C2H5O- + H3O+
pKa= 16
CH3COO- + H3O+ pKa= 4.75
(Bond length C=O 124 pm, C-O 143 pm, but in acetate ion C´-O 127 pm)
Draw the resonance
structures of :
NO3NO2CO32O3
Draw Lewis/resonance structures of:
NO3-, NO2-, CO32-, O3