14.1 Shapes of molecules and ions (HL)

● The octet is the most common electron arrangement because of its stability.

● Exceptions:a) Fewer electrons (incomplete octet) if the central atom is a small atoms, e.g. Be and B

b) More than eight electrons (expanded octet) if the central atom is a 3rd row element or below, e.g. P and S

Species with five negative charge centres

● The shape of a molecule or ion can be predicted by the valence shell electron pair repulsion theory (VSEPR).

● Pairs of electrons (=negative charge centres) arrange themselves around the central atom so that they are as far apart from each other as possible.

● If a molecule has five charge centres and they all are bonding electrons, the shape is triangular bipyramidal.

● If one or more of these five negative charge centres is a non-bonding pair, this will influence the final shape of the molecule.

● One: Tetrahedron

● Two: T-shaped ClF3

● Three: Linear I3

-

Species with six negative charge centres

● Molecules with six charged centres that are all bonding have an octahedral shape, e.g. SF

6.

● One non-bonding pair: square pyramidal BrF5

● Two non-bonding pairs: square planar XeF4

Rivi 1 Rivi 2 Rivi 3 Rivi 40

2

4

6

8

10

12

● Ex. Predict the shape and bond angles of :

PF5, PF

6

-

● Homework: p.129 Ex.14, 15,16● p.155 Ex. 11, 16 a) and b), 17 d)

14.2 Hybridization

● The Lewis structure is a useful model, but it makes one false assumption:

– It assumes that all eight electrons are equal.

● The energy of the electrons are not equal, since some of them exist in s sub-levels and other in p sub-levels.

● A more advanced model of bonding is called the molecular orbital theory.

Molecular orbital theory

● When a bond is formed, atomic orbitals overlap to form new molecular orbitals that are lower in energy.

Sigma (σ) bonds

● A sigma bond is formed when two atomic orbitals on different atoms overlap along a line drawn through the two nuclei (”head on”).

pi (π) bonds

● A pi bond is formed when two p orbitals overlap ”sideways on”.

● The electron density is concentrated in two regions, above and below the plane of the bond axis.

sp3-hybridization

● Methane contains four equal C-H bonds.

● When the carbon bonds to hydrogens, the one 2s and the three 2p orbitals hybridize to form four new energetically equal hybrid orbitals.

● These four sp3-orbitals arrange themselves tetrahedrally, bond angle 109,5º, and four equal σ bonds are formed with the hydrogen.

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom5s2_6.swf

● Hybridization is not just restricted to carbon compounds.

● Tetrahedral compounds (i.e. 4 negative charge

centers) have sp3-hybridization:

NH4

+, NH3, PCl

3 , H

2O

Other compounds with sp3 hybridization

sp2-hybridization

● In ethene, the one 2s orbital hybridizes with two 2p orbitals.The remaining 2p orbital stays as it is.

● These orbitals form three sigma bonds with carbon and 2 hydrogens, bond angle 120o.

● The remaining p orbitals of carbon form a pi bond.

● Therefore, the double bond consists of one sigma and one pi bond.

Other compounds with sp2 hybridization

● Planar triangular compounds (i.e. 3 negative charge

centers),e.g.

BF3 , SO

3, SO

2, propanone

sp hybridization

● In ethyne, the 2s orbital hybridizes with one 2p orbital. The remaining two 2p orbitals stay as they are.

● These orbitals form two sigma bonds with carbon and hydrogen, bond angle 180o.

● The remaining 2 p orbitals of carbon form two pi bonds.

● Therefore, the triple bond consists of one sigma and two pi bonds.

Other compounds with sp hybridization

● Linear compounds (i.e. 2 negative charge

centers), e.g. N2, HCN

14.3 Delocalization of electrons

● http://www.youtube.com/watch?v=eDiDl-ByUd4

Delocalization of electrons in benzene● Delocalization of electrons can occur whenever alternate

double and single bonds occur between carbon atoms.

● The carbon atoms in benzene are sp2-hybridized and each carbon has a p orbital containing one electron.

● Instead of forming double and single bonds, the electrons are delocalized over all six carbon atoms.

Resonance structures

● When writing the Lewis structure for some molecules, it is possible to write more than one structure.

● For example benzene can be written:

● These two structures are known as resonance hybrides. The true structure lies somewhere in between the two.

Other common compounds with resonance structures

● Resonance occurs when more than one valid Lewis structure can be drawn for a molecule.

● NO3

-

● NO2

-

● CO3

-

● O3

● RCOO-

Properties of species with delocalized electrons

1. intermediate bond lengths and strengths

● All affected bonds have equal bond strengths and the bond lengths are intermediate between those of single and double bonds.

● Bond length is dependend on the number of bonds between the atoms: triple < double < single bond

● Bond order: number of shared electron pairs

number of bonding positions

● The higher the bond order, the greater the electron density (=the shorter the bond).

2. Stability

● Delocalization spreads the electrons as far apart as possible and therefore minimizes the repulsion between them.

● Ex. 1 The delocalization makes the benzene molecule more stable by ca. 150 kJ/mol. This is called the delocalization enthalpy or resonance energy.

● This makes the molecule less chemically reactive, since this extra energy has to be put in to break the bonds.

● Ex 2. The relative stability of R-O- ions depends on to what extent the negative charge is delocalized between the two bonds:

● The more stable the ion, the more likely is it formed in a reaction.

3. Electrical conductivity

● Both metals and graphite have delocalized electrons spread out through the entire structure and thus conduct electricity.