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