Expansion of the Octet (HL) (HL IB Chemistry)

Expansion of the octet

• Elements in period 3 and above have the possibility of having more than eight electrons in their valence shell
• This is because there is a d-subshell present which can accommodate additional pairs of electrons
• This is known as the expansion of the octet
• The concept explains why structures such as PCl₅ and SF₆ exist, which have 5 and 6 bonding pairs of electrons respectively, around the central atom

Five electron pairs

Phosphorus pentachloride, PCl₅

• An example of a molecule with five bonding electron pairs is phosphorus pentachloride, PCl₅
• The total number of valence electrons is = P + 5Cl = 5 + (5 x 7) = 40
• The number of bonding pairs is 5, which accounts for 10 electrons
• The remaining 30 electrons would be 15 lone pairs, so that each Cl has 3 lone pairs
• The completed Lewis formula looks like this:

Lewis formula of PCl₅ 

The octet of the central phosphorous atom has been expanded to hold 10 electrons

Sulfur tetrafluoride, SF₄

• The total number of valence electrons is = S + 4F = 6 + (4 x 7) = 34
• The number of bonding pairs is 4, which accounts for 8 electrons
• The remaining 26 electrons would be 13 lone pairs
• Fluorine cannot expand the octet so each fluorine would accommodate 3 lone pairs, accounting for 24 electrons, leaving 1 lone pair on the sulfur (sulfur has expanded the octet)
• The completed Lewis formula looks like this:

Lewis formula of SF₄ 

The octet of the central sulfur atom has been expanded to hold 10 electrons

Chlorine trifluoride, ClF₃

• The total number of valence electrons is = Cl + 3F = 7 + (3 x 7) = 28
• The number of bonding pairs is 3, which accounts for 6 electrons
• The remaining 22 electrons would be 11 lone pairs
• Fluorine cannot expand the octet so each fluorine would accommodate 3 lone pairs, accounting for 18 electrons, leaving 2 lone pairs on the chlorine
• The completed Lewis formula looks like this:

Lewis formula of ClF₃ 

The octet of the central chlorine atom has been expanded to hold 10 electrons

Triiodide ion, l₃^-

• The total number of valence electrons is = 3I + the negative charge = (3 x 7) + 1 = 22
• The number of bonding pairs is 2, which accounts for 4 electrons
• The remaining 18 electrons would be 9 lone pairs
• Iodine would accommodate 3 lone pairs, accounting for 12 electrons, leaving 3 lone pairs on the central iodine
• The completed Lewis formula looks like this:

Lewis formula of I₃^– 

The octet of the central iodine atom has been expanded to hold 10 electrons

Six electron pairs

Sulfur hexafluoride, SF₆

• An example of a molecule with six bonding electron pairs is sulfur hexafluoride, SF₆
• The total number of valence electrons is = S + 6F = 6 + (6 x 7) = 48
• The number of bonding pairs is 6, which accounts for 12 electrons
• The remaining 36 electrons would be 18 lone pairs, so that each F has 3 lone pairs, accounting for all electrons and no lone pairs
• The completed Lewis formula looks like this:

Lewis formula of SF₆ 

The octet of the central sulfur atom has been expanded to hold 12 electrons

Bromine pentafluoride, BrF₅

• The total number of valence electrons is = Br + 5F = 7 + (5 x 7) = 42
• The number of bonding pairs is 5, which accounts for 10 electrons
• The remaining 32 electrons would be 16 lone pairs
• Fluorine cannot expand the octet so each fluorine would accommodate 3 lone pairs, accounting for 30 electrons, leaving 1 lone pair on the bromine
• The completed Lewis formula looks like this:

Lewis formula of BrF₅ 

The octet of the central bromine atom has been expanded to hold 12 electrons

Xenon tetrafluoride, XeF₄

• The total number of valence electrons is = Xe + 4F = 8 + (4 x 7) = 36
• The number of bonding pairs is 4, which accounts for 8 electrons
• The remaining 28 electrons would be 14 lone pairs
• Each fluorine would accommodate 3 lone pairs, accounting for 24 electrons, leaving 2 lone pairs on the xenon
• The completed Lewis formula looks like this:

Lewis formula of XeF₄ 

The octet of the central xenon atom has been expanded to hold 12 electrons

Revisiting Valence Shell Electron Pair Repulsion Theory (VSEPR)

• Molecular shapes and the angles between bonds can be predicted using the three basic rules associated with valence shell electron pair repulsion theory known by the abbreviation VSEPR theory
• VSEPR theory consists of three basic rules:
  • All electron pairs and all lone pairs arrange themselves as far apart in space as possible.
  • Lone pairs repel more strongly than bonding pairs
  • Multiple bonds behave like single bonds
• For more information about valence shell electron pair repulsion theory, see our revision note on Shapes of Molecules

Molecular geometry versus electron domain geometry

• It is important to distinguish between molecular geometry and electron domain geometry in exam questions
  • Molecular geometry refers to the shape of the molecules based on the relative orientation of the atoms
  • Electron domain geometry refers to the relative orientation of all the bonding and lone pairs of electrons
• The Lewis formula for water enables us to see that there are four electron pairs around the oxygen so the electron domain geometry is tetrahedral
• However, the molecular geometry shows us there are two angled bonds so the shape is bent, angular, bent linear or V-shaped (when viewed upside down)

Lewis formula and molecular shape of water

The Lewis formula of water and molecular shape can be used to determine the electron domain and molecular geometries

Five electron domains

Table of molecular geometries associated with five electron domains

Electron domain geometry	Bonding pairs	Lone pairs	Molecular geometry	Shape example
trigonal bipyramidal	5	0	trigonal bipyramidal
trigonal bipyramidal	4	1	seesaw
trigonal bipyramidal	3	2	T-shape
trigonal bipyramidal	2	3	Linear

 

• PCl₅  is a symmetrical molecule so the electron cloud charge is evenly spread
  • This means that it will be a non-polar molecule as any dipoles from the P–Cl bonds would be cancelled out
• SF₄ and ClF₃ are asymmetrical molecules having one or two lone pairs on one side of the central axis making the overall molecule polar

Six electron domains

Table of molecular geometries associated with six electron domains

Electron domain geometry	Bonding pairs	Lone pairs	Molecular geometry	Shape example
octahedral	6	0	octahedral
octahedral	5	1	square based pyramid
octahedral	4	2	square planar

 

• SF₆ is a symmetrical molecule so the electron cloud charge is evenly spread with 90^o between the bonds
  • This means that it will be a non-polar molecule as any dipoles from the S–F bonds would be cancelled out
• XeF₄ is also non-polar despite having two lone pairs
  • The bonding pairs are at 90^o to the plane and the lone pairs are at 180^o
  • The lone pairs are arranged above and below the square plane resulting in an even distribution of electron cloud charge
• BrF₅ is asymmetrical having a lone pair at the base of the pyramid making the overall molecule polar