4.3    Shapes of  Molecules

Valence shell electron pair repulsion(VSEPR) theory

• All electrons are negatively-charged, so they will repel each other when they are close together
• So, a pair of electrons in the bonds surrounding the central atom in a molecule will repel the other electron pairs. This repulsion forces the pairs of electrons apart until the repulsive forces are minimised
• The amount of repulsion is as follow:

 

• General steps to determine the shape of a molecule:
  1. determine the number of valence electrons in the central atom
  2. find the total number of electrons surrounding the central atom by adding the number of  shared electrons to it. (Dot-and-cross diagram might be necessary)
  3. find the number of electron pairs by dividing the total number of electrons by two
  4. determine how many pairs is/are bond pairs and lone pairs. (A double bond or triple bond is counted as one bond pair)
  5. refer to the table to obtain the shape of the molecule

Effect of  lone pair on bond angle

• For methane, ammonia and water, the electron pair geometries are tetrahedral. However, the molecular geometries are different
• In methane, all the bonds are identical, repulsion between the bonds is the same. Thus, methane has a perfect tetrahedral structure with bond angle 109.5°.
• In ammonia, the repulsion between the lone pair and the bond pairs is stronger than in This forces the bond angle to decrease slightly to 107°.
• In water, there are two lone pairs and thus the repulsion is the greatest, the two bond pairs are pushed closer to one another and the bond angle is reduced to 104.5°.

 

Effect of electronegativity on bond angle

• Water and hydrogen sulfide have the same general shape with the same number of bond pairs and lone pairs. However, their bond angles are different
• This is because oxygen has a higher electronegativity than The bond pairs of electrons are closer to the oxygen atom compared to the sulfur atom.
• This results in greater repulsion in the O-H bonds than in the S-H bonds. Therefore, the bond angle increases from 92.5° to 104.5°.

 

Sigma(σ) bond and pi(π) bond

• A sigma bond is formed by orbitals from two atoms overlapping end-to-end.
• In a sigma bond, the electron density is concentrated between the two nuclei

• A pi bond is formed by the p orbitals from two atoms overlapping sideways.
• In a pi bond, there are two regions of high electron density alongside the nuclei.
• A pi bond is weaker than a sigma bond because the overlapping of charge clouds is less than in a sigma bond
• In covalent molecules, single bonds are sigma bonds(σ), a double bond consists of one sigma bond and one pi bond(1σ, 1π), and a triple bond consists of one sigma bond and two pi bonds(1σ, 2π).

Hybridisation

• Hybridisation is the mixing of atomic orbitals to produce a new set of hybrid orbitals of equivalent energies. This is a theory used to account for the discrepancies in explaining the shapes of covalent molecules
• There is a problem with simple view on methane, CH4. Methane has two unpaired electrons only in its outer shell to share with the hydrogen atoms, but the formula of methane is not CH2.
• The concept of hybridisation is used to account for this discrepency
• General steps in hybridisation:
  1. promotion of electron
  2. mixing of orbitals to produce a new set of hybrid orbitals of equivalent energies(sp, sp² or sp³ hybrid orbitals)
  3. forming of a new molecular orbital

sp³ hybridisation

• An example of compound which undergoes sp³ hybridisation is methane, CH4.
• The carbon atom uses some energy to promote one of its electron from 2s to empty 2p orbital so that there are four unpaired electrons for covalent bonding
• The carbon now is said to be in an excited state(C*).
• The orbitals then ‘mix’ or hybridise to produce four hybrid orbitals of equivalent energies. The new orbitals are called sp³ hybrid orbitals because they are made from one s orbital and three p orbitals.
• Each hybrid orbital has one big lobe and one small They rearrange themselves so that they are as far as possible to form a tetrahedral geometry. The hybrid orbitals are 109.5° apart.
• The s orbitals from the hydrogen atoms then overlap with the four hybrid orbitals to form four sigma bonds because the overlapping is end-to-end. All the bonds are identical
• Another example is ethane, C2H6. The two carbon atoms undergo sp³ hybridisation to form four hybrid orbitals. The two carbon atoms are bonded by overlapping one of their hybrid orbitals. The remaining ones then overlap with the s orbitals of the hydrogen atoms.
• The bond angle is approximately 109.5°. This is an approximation because all the bonds are not

sp² hybridisation

• An example of compound which undergoes sp² hybridisation is ethene, C2H4.
• The same thing happens as in sp³ hybridisation, except that this time the carbon atoms ‘mix’ or hybridise three of  the four orbitals only because the carbon atom is bonding with three other atoms only.
• This produces three sp² hybrid orbitals because they are made from one s orbital and two p Another p orbital remains unchanged and it is perpendicular to the plane containing the hybrid orbitals.
• The hybrid orbitals rearrange themselves so that they are as far as possible, that is, a trigonal planar arrangement, the hybrid orbitals are 120° apart.
• The hybrid orbitals then overlap with s orbitals from the hydrogen atoms and another hybrid orbital from the other carbon atom to form ftve sigma bonds. The remaining p orbitals overlap sideways to form a pi bond. A double bond is formed between the two carbon atoms
• Another example is boron trichloride, BCl3. The boron atom undergoes sp² hybridisation to produce three sp² hybrid orbitals. The hybrid orbitals rearrange themselves to form a trigonal planar The p orbitals from chlorine atoms then overlap with the hybrid orbitals to form three sigma bonds.

 

sp hybridisation

• An example of compound which undergoes sp hybridisation is ethyne, C2H2.
• The same thing happens as in sp³ and sp² hybridisation, except that this time the carbon atoms ‘mix’ or hybridise two of  the four orbitals only because the carbon atom is bonding with two other atoms only.
• This produces two sp hybrid orbitals because they are made from one s orbital and one p orbital. The other two p orbitals remain unchanged and they are perpendicular to each other and to the two hybrid orbitals
• The hybrid orbitals rearrange themselves so that they are as far as possible, that is, a linear arrangement, the hybrid orbitals are 180° apart
• The hybrid orbitals overlap with the s orbitals from the hydrogen atoms and to the hybrid orbital from the other carbon atom to form three sigma bonds. The remaining p orbitals overlap sideways to form two pi bonds. A triple bond is formed between the two carbon atoms

Example of covalent molecule with multiple hybridisations

1) In carbon dioxide, CO2, the carbon atom undergoes sp hybridisation while the oxygen atoms undergo sp² hybridisation. The overlapping of the hybrid and p orbitals are shown in the diagram.