DUNCANRIG SECONDARYADVANCED HIGHER CHEMISTRY
This topic concerns the 3 dimensional shape of molecules. Molecular shape is very important when chemical reactions are considered. This is especially true in Biochemical reactions. This aspect of molecular shape will be dealt with in the medicinal topic of the Organic Chemistry unit. For now we will limit the discussion to simple molecules with relatively few atoms.
You are already familiar with the basic shape of methane and other hydrocarbons.
The methane molecule has a tetrahedral geometry and the angle between each of the bonds is 109.5 degrees
The question is - why does a methane molecule adopt this particular shape?
The basic answer to this question is due to the electrons in the covalent bonds holding the atoms together.
Each covalent bond is a shared pair of electrons. As electrons are negatively charged, the electrons in a bond will repel the electrons in other bonds.
This diagram shows the covalent bonds in methane.
Each bonded pair of electrons will repel the other three bonded pairs of electrons.
The theory known as valence shell electron pair repulsion (VSEPR) is used to explain why molecules adopt a particular shape.
The basic principle behind the theory states:
Consider the molecule shown below. It contains atom X which has three covalent bonds to the atoms of Y.
In this diagram two of the bonds angles are 90 degrees andone is 180 degrees. This is an unfavourable geometry as the repulsion of the bonded electrons is unequal – there will be greater repulsion between the electron pairs which are 90 degrees apart than between the electron pairs which are 180 degrees apart.
In this diagram all the bond angles are all 120 degrees and so the electron pairs are as far apart from each other as possible. This will be the lowest energy situation for this molecule as the repulsive forces are pushed as far apart as possible.
In this example atom X has three bonded pairs of electrons –atom X is bonded to three other atoms. This is not always the case. Molecules can have more or less bonded pairs than the molecule shown (e.g. methane has 4 bonded pairs). The number of pairs of electrons a molecule has will dictate the shape adopted by the molecule. In addition some molecules have non – bonded pairs of electrons called LONE PAIRS which also influence molecular shape.
When all the pairs of electrons are bonding pairs the repulsive force generated by each pair will be very similar. This results in the FIVE basic molecular geometries (shapes) shown below
Knowing the number of electron pairs around the central atom allows the molecular geometry to be determined.
To determine the number of electrons pair around the central atom in a molecule or ion we can use the following formula
If the particle is an ion the number of electrons lost or gained (the value of the charge) must be added (- ion) or subtracted (+ ion) from the total number of electrons before dividing by two.
1. Determine the number of electron pairs in the following molecules or ions.
a.H2Ob.NH3c.SF6d.ICl4-
e.BCl3f.NF3g.SiCl4h.PF5
i.PH4+j.BeF2k.IF5l.SiCl62-
Another important factor in determining molecular geometry is the nature of the electron pairs.
Consider the bonding in the ammonia molecule, NH3
The electronic configuration of nitrogen is
1s2 2s2 2p3
Nitrogen has three unpaired electrons in its valenceshell . It is these electrons which form the three bonds with the hydrogen atoms in ammonia. However, nitrogen also has a pair of electrons in its valence shell – the 2 electrons in the 2s sub-shell. These are usuallynon – bonding electrons.
Ammonia has 4 pairs of electrons. Three are bonding pairs – hence three bonds, and one lone pair {non bonding pair}.
The geometry of the ELECTRONS is TETRAHEDRAL but the geometry of the ammonia molecule is a TRIGONAL PYRAMID.
Study the examples below which show how to determine the number of bonding and non-bonding pairs in a molecule or ion.
This molecule has (8+4)/2 = 6 electron pairs.
This indicates the shape the electron pairs will be octahedral.
As the molecule has four F atoms connected to the central Xe atom it has four bonding pairs of electrons and so it must also have two non- bonding pairs.
The shape the atoms adopt is based on the octahedral geometry of the electron pairs but as there are only four atoms on the central xenon atom the shape of the molecule is SQUARE PLANAR.
This molecule has (6+2)/2 = 4 electrons pairs
This indicates the shape the electron pairs will be tetrahedral.
As there are only two hydrogen atoms connected to the central oxygen it has two bonding pairs of electrons and so it must have two non – bonding pairs of electrons.
The shape the atoms adopt is based on the tetrahedral geometry of the electron pairs but as there are only two atoms on the central oxygen atom the shape of the molecule is described as NON-LINEAR or ANGULAR.
1. Go back to the question on page 5 and decide how many bonding and how many non - bonding electron pairs each substance has.
In the exam you may be asked to draw a diagram of a molecule or ion which clearly shows the shape of the molecule or the shape that the electron pairs adopt.Your drawing must not look “FLAT” and must give an idea of the three dimensional arrangement of the electron pairs or of the atoms.
The diagrams below show one way to do this.
Match the electron pair geometry of the molecules below to the shapes above.
a. BrF5 b. PCl4+ c AlBr3 d. XeF2e. H3O+
Look back at page three to see the bond angles formed when all the electron pairs around the central atom are bonding pairs.
If a molecule, or ion, has lone pairs of electronson the central atom, the shapes are slightly distorted away from the regular shapes. This is because of the extra repulsion caused by the lone pairs.
As a result of the extra repulsion, bond angles tend to be slightly less as the bonds are“squeezed” together.
The three molecules above showdecrease in bond angle as the numberof NON-BONDED electron pairs exert a greater repulsive effect on the bonded pairs. The same effect is seen in other molecules.
Remember that lone pairs of electrons cause bond angles to vary from normal.
A list of “learning outcomes” for the topic is shown below. When the topic is complete you should review each learning outcome.
Your teacher will collect your completed notes, mark them, and then decide if any revision work is necessary.
- State that the shapes of simple molecules or ions can be explained by the VESPR theory.
- Be able to determine the number of bonding and the number of non-bonding electron pairs around a central atom.
- Be able to relate the shape of a molecule or ion to the number of electron pairs around the central atom,
- Be able to use draw perspective diagrams of the following shapes. Trigonal planar, trigonal pyramid, tetrahedral, trigonalbipyramid, square planar and octahedral.
- State the bond angles the molecular shapes listed above.
- State that lone pairs of electrons exert more repulsive force than bonded pairs of electrons and that this results in molecules or ions with distorted shapes.
I have discussed the learning outcomes with my teacher.
My work has been marked by my teacher.
Teacher Comments.
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