Are Giant Molecular Structures Soluble In Organic Solvents

[Theodora Andy]

ImportantThere's not much point in reading this page unless you are reasonably happy about the origin of hydrogen bonding and van der Waals forces. Follow these links first if you aren't sure about these.

Molecules are made of fixed numbers of atoms joined together by covalent bonds, and can range from the very small (even down to single atoms, as in the noble gases) to the very large (as in polymers, proteins or even DNA).

The covalent bonds holding the molecules together are very strong, but these are largely irrelevant to the physical properties of the substance. Physical properties are governed by the intermolecular forces - forces attracting one molecule to its neighbours - van der Waals attractions or hydrogen bonds.

Molecular substances tend to be gases, liquids or low melting point solids, because the intermolecular forces of attraction are comparatively weak. You don't have to break any covalent bonds in order to melt or boil a molecular substance.

Note: This is really important! You can make yourself look extremely stupid if you imply in an exam that boiling water, for example, splits it into hydrogen and oxygen by breaking covalent bonds. Exactly the same water molecules are present in ice, water and steam.

The size of the melting or boiling point will depend on the strength of the intermolecular forces. The presence of hydrogen bonding will lift the melting and boiling points. The larger the molecule the more van der Waals attractions are possible - and those will also need more energy to break.

The problem is the hydrogen bonds between the water molecules. If methane were to dissolve, it would have to force its way between water molecules and so break hydrogen bonds. That costs a reasonable amount of energy.

The only attractions possible between methane and water molecules are the much weaker van der Waals forces - and not much energy is released when these are set up. It simply isn't energetically profitable for the methane and water to mix.

The reversible arrows show that the reaction doesn't go to completion. At any one time only about 1% of the ammonia has actually reacted to form ammonium ions. The solubility of ammonia is mainly due to the hydrogen bonding and not the reaction.

Molecular substances are often soluble in organic solvents - which are themselves molecular. Both the solute (the substance which is dissolving) and the solvent are likely to have molecules attracted to each other by van der Waals forces. Although these attractions will be disrupted when they mix, they are replaced by similar ones between the two different sorts of molecules.

Molecular substances won't conduct electricity. Even in cases where electrons may be delocalised within a particular molecule, there isn't sufficient contact between the molecules to allow the electrons to move through the whole solid or liquid.

Iodine is a dark grey crystalline solid with a purple vapour. M.Pt: 114C. B.Pt: 184C. It is very, very slightly soluble in water, but dissolves freely in organic solvents.Iodine is therefore a low melting point solid. The crystallinity suggests a regular packing of the molecules.

The orientation of the iodine molecules within this structure is quite difficult to draw (let alone remember!). If your syllabus and past exam papers suggests that you need to remember it, look carefully at the next sequence of diagrams showing the layers.

All these diagrams show an "exploded" view of the crystal. The iodine molecules are, of course, touching each other. Measurements of the distances between the centres of the atoms in the crystal show two different values:

The iodine atoms within each molecule are pulled closely together by the covalent bond. The van der Waals attraction between the molecules is much weaker, and you can think of the atoms in two separate molecules as just loosely touching each other.

There are lots of different ways that the water molecules can be arranged in ice. This is one of them, but NOT the common one - I can't draw that in any way that makes sense! The one below is known as "cubic ice", or "ice Ic". It is based on the water molecules arranged in a diamond structure.

This is just a small part of a structure which extends over huge numbers of molecules in three dimensions. In the diagram, the lines represent hydrogen bonds. The lone pairs that the hydrogen atoms are attracted to are left out for clarity.

Note: Don't worry about this problem. If asked to draw ice in an exam at this level (16 - 18 year olds), don't try to be too clever. It is probably best not to go beyond the top five molecules in the above diagram. This will show the essential features of the bonding in the structure without getting bogged down in stuff which is far beyond this level.

If you are interested in following this up, try a Google search using the search term ice structure hexagonal cubic (or something similar). This will throw up lots of information together with an assortment of fairly dreadful diagrams which I for one don't have the visual imagination to unscramble!

The hydrogen bonding forces a rather open structure on the ice - if you made a model of it, you would find a significant amount of wasted space. When ice melts, the structure breaks down and the molecules tend to fill up this wasted space.

Remnants of the rigid hydrogen bonded structure are still present in very cold liquid water, and don't finally disappear until 4C. From 0C to 4C, the density of water increases as the molecules free themselves from the open structure and take up less space. After 4C, the thermal motion of the molecules causes them to move apart and the density falls. That's the normal behaviour with liquids on heating.

Polymers like poly(ethene) - commonly called polythene - consist of very long molecules. Poly(ethene) molecules are made by joining up lots of ethene molecules into chains of covalently bound carbon atoms with hydrogens attached. There may be short branches along the main chain, also consisting of carbon chains with attached hydrogens. The molecules are attracted to each other in the solid by van der Waals dispersion forces.

Low density polythene has lots of short branches along the chain. These branches prevent the chains from lying close together in a tidy arrangement. As a result dispersion forces are less and the plastic is weaker and has a lower melting point. Its density is lower, of course, because of the wasted space within the unevenly packed structure.

This page describes the structures of giant covalent substances like diamond, graphite and silicon dioxide (silicon(IV) oxide), and relates those structures to the physical properties of the substances.

This is a giant covalent structure - it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable - depending on the size of the crystal.

Note: We quoted the electronic structure of carbon as 2,4. That simple view is perfectly adequate to explain the bonding in diamond. If you are interested in a more modern view, you could read the page on bonding in methane and ethane in the organic section of this site. In the case of diamond, each carbon is bonded to 4 other carbons rather than hydrogens, but that makes no essential difference.

is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.

Graphite has a layer structure which is quite difficult to draw convincingly in three dimensions. The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced.

You might argue that carbon has to form 4 bonds because of its 4 unpaired electrons, whereas in this diagram it only seems to be forming 3 bonds to the neighbouring carbons. This diagram is something of a simplification, and shows the arrangement of atoms rather than the bonding.

Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon atom become delocalised over the whole of the sheet of atoms in one layer. They are no longer associated directly with any particular atom or pair of atoms, but are free to wander throughout the whole sheet.

If you are interested (beyond A'level): The bonding in graphite is like a vastly extended version of the bonding in benzene. Each carbon atom undergoes sp2 hybridisation, and then the unhybridised p orbitals on each carbon atom overlap sideways to give a massive pi system above and below the plane of the sheet of atoms.

The important thing is that the delocalised electrons are free to move anywhere within the sheet - each electron is no longer fixed to a particular carbon atom. There is, however, no direct contact between the delocalised electrons in one sheet and those in the neighbouring sheets.

The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalised electrons. So what holds the sheets together?

In graphite you have the ultimate example of van der Waals dispersion forces. As the delocalised electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below - and so on throughout the whole graphite crystal.

has a high melting point, similar to that of diamond. In order to melt graphite, it isn't enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure.

has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. You can think of graphite rather like a pack of cards - each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper.

3a8082e126