Topic 14 Bonding (HL)
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Transcript of Topic 14 Bonding (HL)
Topic 14 Bonding (HL) • Shapes of molecules and ions• Hybridisation• Delocalisation of electrons
14.1 Shapes of molecules and ions
• Valence Shell Electron Pair Repulsion-VSEPR 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 bondAxial: 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 shape- XeF4
• 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 C2H6
single 348 154 1 s
EtheneC2H4
double 612 134 1 s + 1p
EthyneC2H2
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 s-orbitals, with s-bonds:
The green sp2-orbital is used to s –bond, and the red p-orbital 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 sp-orbitals (trigonal planar shape). Two p-orbitals is left over (red)
• Two carbons with sp-orbitals now bind 2 hydrogen s-orbitals, 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
• 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
Benzene
Resonance • a and b are resonance structures of benzene
• c is a resonance hybrid- the most stable form. By delocalisation of the electrons the molecule gain resonance energy
• d resonance in pyridined
Why has phenol acidic properties?
Low pKa- strong acid (HCl -7)High pKa (1-14)- weak acidpKa > 14 no acid
Phenol pKa 8
Why has acetic acid acidic properties but not ethanol?
C2H5OH + H2O ↔ C2H5O- + H3O+ pKa= 16
CH3COOH + H2O ↔ 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 resonancestructures of :
NO3-
NO2-
CO32-
O3
Draw Lewis/resonance structures of: NO3
-, NO2-, CO3
2-, O3