Topic 7: Covalency and Bond Breaking

Electron density maps

-What are they?

Electron density maps show just how many electrons there are in a certain area of the molecule. Contours are used to establish that the largest electron density is closest to the nucleus. This is, obviously, because of the nuclear attraction experienced between the two. More importantly, electron density maps have been used to differentiate between covalent and ionic compounds, as well as establishing molecular structure.

- How are they obtained?

X-ray diffraction. When X-rays of a particular frequency (and wavelength) are targeted on a crystalline solid, the rays diffract and give a pattern. This pattern is related to the pattern of the electrons in the crystals. When processed, the familiar electron density map is produced. Further analysis may reveal the atoms and their positions.

- How can they be used to investigate ionic and covalent bonds?

For example, the electron density map of sodium chloride, an ionic compound, will look something like fig 7.1. We can conclude the following:

-          There are NO electrons between the ions

-          Electron density is smooth around the ions, and most concentrated near the nucleus.

Therefore, ionic compounds have NO electrons between their ions. This is the main feature of an electron density map of an ionic compound.

As for a covalent molecule, fig 7.2:

-          There is a considerable electron density between the atoms.

-          The electron density isn’t smooth around the atoms.

Therefore, electrons are SHARED in covalent molecules. This property distorts the electron densities around the atoms slightly.

- How can you tell which atom is which?

Since different atoms have different radii and electron densities, you can differentiate between them. For example, in fig 7.2, Oxygen is larger than the carbon, and has more electron density ‘rings’.

Hydrogen atoms are always missing from the electron density maps. This is because they have a very weak electron density (Only one electron on the outer shell.) which cannot be presented in X-ray diffraction.

Thus, ionic bonds do not consist of shared electrons, while molecular bonds (covalent bonds) consist of shared electrons.

Shapes of molecules

Although you may see display formulae suggesting that molecules are flat, the fact is, it isn’t usually the case.

The correct way to display the shape of a molecule is to use the following convention:

But what is it that determines the way these molecules are shaped? We can investigate this using the following table:

 Name Number of lone pairs Bond Angle Carbon dioxide 0 180° Boron trichloride 0 120° Methane 0 109.5° Ammonia 1 107° Water 2 104.5°

Why does Boron Trichloride have such a shape? We can see that Boron trichloride has three bonds. Each of those bonds contains electrons (bond pair). Since electrons have the same charge, they repel each other. Since all the electrons are in bonds, they repel each other by the same amount. This means they repel each other till each bond is furthest away from the other. This means there is a bond angle of 120°(360°/3).

But, why doesn’t Methane do the same? The only way for four bonds to be as far as possible is to form a tetrahedral shape. This gives a bond angle of 109.5

This all seems to fit together, until we get to Ammonia. Ammonia only has three bonds, so why doesn’t it do what Boron Trichloride does?

The answer lies in the presence of a lone pair. A lone pair, as the name suggests, is an un-bonded pair of electrons. Since it is not bonded, it is situated closer to the nucleus. This means it has a greater repulsion effect on the other bonds. This means, in addition to the three bonds repelling each other, the lone pair exerts an even stronger repulsion. This gives us a trigonal pyramidal shape, and a bond angle of 107°.

We can conclude that:

Lone pair – Lone pair > Lone pair – bond pair > bond pair – bond pair

Have a look at the last three. For every additional lone pair in a structure with 4 electron centres, there is a decrease of 2.5° in the bond angle.

Now, a more interesting issue to tackle. We should expect, from the electron repulsion theory explained above that Hydrogen Sulphide has the same bond angle as Hydrogen Oxide (Water). Both have 2 lone pairs, and two bond pairs, but when we look at the bond angle, one is 104.5° (Water) and the other is 98°!! Email me if you think you know the answer, and I’ll tell you if you are correct.

Covalent giant structures (Macromolecular structures)

Both diamond and graphite are of the same element: carbon. Then what causes the major differences between the two?

Graphite is constructed of a carbon atoms covalently bonded to three other carbon atoms. This forms sheets of carbon hexagons, in which the sheets are on top of each other to form layers. Because each carbon atom is bonded to three others, then there is a single lone electron left. This electron delocalises, and with all the other delocalised electrons, forms a sea of electrons which is shared by all the atoms. This is a kind of lubricating layer, which enables the layers of carbon to slide over each other. The delocalised electrons also allow it to conduct electricity.

Diamond has a tetrahedral shape. Each atom is bonded covalently to 4 other atoms. Since all the electrons are taken up in the bonds, there are no delocalised electrons. Thus diamond doesn’t conduct electricity.

Both have a high melting point, because of the numerous covalent bonds that keep the structure together.

Dative Covalence

Not all covalent bonds involve the sharing of electrons. The most common example used (even in our textbook) is ammonium.

The ammonia molecule bonds with a positive hydrogen ion (a proton) to form the positive polyatomic ion ammonium. If we try and demonstrate this using our normal bonding rules, we discover that the hydrogen has no electrons to share! So, how does this bonding occur?

The ammonia molecule has a lone pair of electrons. All the rest of the nitrogen’s electrons are bonded to hydrogen. A covalent bond is formed between the lone pair and the hydrogen ion. This means, although they’re being shared, the electrons all come from the nitrogen. This is dative covalent bonding.

I could have drawn diagrams to show this, but the textbook has enough of those. As you can see, the dative covalent bond is demonstrated using an arrow sign, pointing from the donating atom (nitrogen) to the receiving atom (hydrogen).

ELECTRONEGATIVITY

The more electronegative an atom is, the greater its ability to attract electrons to itself. In a diatomic covalent molecule, there are shared electrons between both atoms. Both atoms have positive nuclei. The nucleus with more protons attracts the shared electrons more, bringing the shared electrons closer to it.

This is where, almost everything you’ve been told about simple molecules in GCSE, was a lie…

The fact these electrons shift towards one of the atoms induces a negative charge to that atom (there is an excess of negative charge on that side). This also means there is a lack of electrons on the other atom, which induces a positive charge on it. We are then left with a molecule with two opposite charges, just like an ionic bond! The stronger the electronegativity of the atom, the stronger the charges, the more obvious the ionic properties. This introduces some ionic properties to the molecule. It is this property that allows the covalent molecule HCl to dissolve in water.

The covalent bond formed with charges is called a polar bond. A permanent dipole is formed. The elements F, O, Cl, N are highly electronegative. On the other hand, elements like caesium and francium are least electronegative.

The rest of this topic will be published soon…