Here we will consider how the type of bonding and the forces between atoms, ions and molecules give rise to different physical properties.
Properties of simple molecules
Many substances, especially organic substances, consist of individual molecules with strong covalent bonds holding the atoms together within the molecule, and weak intermolecular forces holding one molecule to another.
To change the state of such substances from solid to liquid to gas we have to overcome the intermolecular forces that hold the molecules to one another, but we do not break any of the covalent bonds – the molecules are identical regardless of their physical state.
|Characteristic||Simple Molecular Structure|
|Boiling and melting points||Low Reason: only weak intermolecular forces need to be overcome, which requires relatively little energy.|
|Electrical conductivity||Not in any state Reason: No ions or free electrons to act as charge carriers|
|Solubility||Often dissolve in non-polar solvents (e.g. hydrocarbons)|
Properties of giant structures
We need to be able to distinguish between the giant structures (and the types of bonding in them), and their characteristic properties. To change the state of a giant structure from solid to liquid to gas, the strong bonds between the atoms or ions throughout the structure have to be broken.
|Giant Ionic Lattice||Giant Covalent Lattice||Giant Metallic Lattice|
|Boiling and melting points||High
Reason: strong ionic bonds throughout the lattice have to be broken, which requires a lot of energy.
Reason: strong covalent bonds throughout the lattice have to be broken, which requires a lot of energy.
Reason: strong metallic bonds throughout the lattice have to be broken, which requires a lot of energy.
|Electrical conductivity||Solid: No
Molten or in solution: Yes
Reason: Ions in solid are fixed in lattice, but ions in solution/liquid are mobile and can therefore act as charge carriers
|Not in any state
Reason: No ions or free electrons to act as charge carriers
Graphite has one delocalized electron per C atom which is free to act as charge carrier
Sea of delocalized electrons are mobile and act as charge carriers
|Solubility||Often dissolve in polar solvents (e.g. water)||Insoluble
Case Studies in Bonding and Structure
i) Diamond – a typical giant covalent lattice
In diamond each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement.
Since all four outer shell electrons in carbon are used for bonding there are no charge carriers in diamond, so it does not conduct electricity.
Diamond sublimes at an extremely high temperature of 5,530°C. This, the extreme strength of the substance, and its lack of solubility in any solvents are all due to the very strong covalent bonds throughout the lattice. These must be broken, requiring a lot of energy, before the structure will melt, break or dissolve.
The second element in Group 4, silicon, has the same giant 3D structure. Silicon has high melting and boiling points of 1414°C and 3265°C respectively, for the same reasons as diamond. It is strong strong and very brittle, and is a semiconductor, having electrical conductivity somewhere between that of metals and non-metals.
ii) Graphite – an unusual giant covalent lattice
In graphite each carbon atom is covalently bonded to three others forming a hexagonal layer structure with 120° bond angles. Each layer has further layers above and below.
The one remaining electron in the outer shell of each carbon atom is not used in bonding. This becomes delocalized and can move throughout the lattice carrying electrical charge, so graphite conducts electricity like a metal even though it is not a metal.
Graphite is a soft substance. The layers are held one to another by weak London forces, which means that the layers can easily slide over each other making graphite slippery and good for use as a lubricant. Pencils leave their marks on paper because the graphite ‘lead’ in the pencil leaves layers of graphite stuck to the paper.
Graphite has a very high melting point (and boiling point) and is insoluble. To melt or dissolve graphite would require the layers to be broken down, and this means large numbers of very strong covalent bonds have to be broken requiring lots of energy.
Graphene is a two-dimensional material. Each carbon atom in graphene is covalently bonded to 3 neighbouring atoms. Each atom has four outer shell electrons. Three of these are involved in single covalent bonds to adjacent carbon atoms, and a fourth electron from each atom is delocalised and free to move throughout a conjugated network that extends throughout the lattice. This arrangement means that carbon-carbon bonds in graphene are extremely strong.
Thanks to this structure, graphene has some unusual properties:
- Huge surface area to weight ratio, up to 2630 m2/g
- High electrical conductivity
- Exceptionally stiff: Young’s modulus about 1 TPa
- Strong – fracture strength approximately 130 GPa
- The highest thermal conductivity of all known materials
These unique properties mean that graphene could prove valuable in many application areas. Some examples include:
- Transparent conductive films, e.g. for touchscreen displays
- Chemical, biological and mechanical sensors, e.g. for detecting nerve agents, diseases and strain
- Additives for composite materials, e.g., for aircraft, car tyres and prosthetics
- Inert coatings for protecting materials from harsh environments
The first thorough investigations of graphene properties were reported in 2004, based on material produced by micromechanical exfoliation, also known as the ‘sticky tape’ method. Despite producing high-quality material, this approach is laborious and is only capable of providing small-area samples. There are now many scalable methods for graphene production e.g. chemical vapour deposition (CVD) that have been developed for the industrial production of graphene.
iv) Sodium Chloride – a typical giant ionic lattice
Sodium chloride has a regular 3D lattice structure in which sodium ions and chloride ions alternate. The electrostatic attraction between the ions is non-directional so every sodium ion attracts chloride ions all around it – above, below and on each side. Similarly each chloride ion attracts six sodium ions.
The formula NaCl does not indicate that sodium and chloride ions “go around in pairs”, but that there is a ratio of one sodium ion for every chloride ion in the lattice – it is an empirical formula.
Because the electrostatic attractions between the oppositely-charged ions run throughout the lattice, sodium chloride is hard and brittle. These attractions are strong, and must be overcome before the lattice can be broken down to form a liquid, which requires a lot of energy. Sodium chloride thus has a high melting and boiling point.
Sodium chloride dissolves in polar solvents such as water. The energy needed to break down the lattice so the ions can be in solution is offset by the energy released when bonds are formed between the ions and the water molecules.
Once in solution (or melted), the sodium and chloride ions are free to move around. They are charge carriers, so in these states sodium chloride is able to conduct electricity. When solid the ions are locked in place in the lattice and although they are charged, they are unable to move and carry the electrical charge. This means that solid sodium chloride does not conduct electricity.
v) Iodine – a simple molecular lattice
Iodine has a simple molecular structure. Iodine atoms are bonded in pairs by covalent bonds, so the formula is I2.
In the solid state, the weak intermolecular forces between the iodine molecules hold them together in a 3D lattice structure. This is a simple molecular lattice. There are no delocalized electrons in iodine, and no ions, so there are no charge carriers and iodine cannot conduct electricity.
Iodine has the unusual property that when heated it goes directly from being solid to being a gas. This property is called sublimation. It does this at a low temperature (114°C) because only weak intermolecular forces hold the iodine molecules together and these take little energy to overcome. Iodine vapour still consists of I2 molecules – no bonds have been broken in turning iodine into a gas.
Being non-polar iodine dissolves readily in most non-polar organic solvents such as hexane or trichloromethane. It is almost completely insoluble in water because to dissolve in water the hydrogen bonds between the water molecules have to be broken, and there are no new bonds formed between iodine and the water so the energy cannot be offset by any released energy.
vi) Water and Ice
In water there are two lone pairs on the δ- O atom, and two δ+ H atoms so each water molecule can form two hydrogen bonds making networks and lattices of water molecules are possible.
Ice consists water molecules held in a 3D lattice by a network of hydrogen bonds. The two lone pairs on the oxygen atom in each water molecule form hydrogen bonds to the hydrogen atoms of two other water molecules forming a 3D network of hydrogen bonds.
The hydrogen bonds do not take anywhere as much energy to overcome as covalent bonds would, so the temperature required to break down the simple molecular lattice in ice is much lower than the temperature required to melt giant covalent lattices.
Water has a number of unusual properties which arise from the hydrogen bonding between water molecules:
- anomalously high boiling point and melting point compared to similar simple molecules without hydrogen bonding
- ice floats (solids are usually more dense than same substance in liquid state)
- surface tension (needles float, pond skaters walk on it)
The strength of the hydrogen bonds which need to be overcome in ice before it can be melted give rise to the unusually high melting point. In the liquid state water molecules collect in hydrogen-bonded groups (clumps). These must be broken up and separated before water can be boiled and water vapour formed. The energy required to do this gives rise to the unusually high boiling point.
In ice, each hydrogen atom is covalently bonded to one oxygen atom and hydrogen bonded to another. The hydrogen bonds (0.159nm) are LONGER than the covalent bonds (0.096nm), which spaces out the molecules in the ice lattice, making them typically further apart than in water – ice has a more open structure than water, so has a lower density (which is unusual for a solid). This is why ice floats, and why it expands when it freezes.
At the surface, the water molecules form a 2-D network of hydrogen bonded water molecules – this is the origin of surface tension, explaining why a water droplet can stand up on a flat surface – it is difficult to penetrate the surface because that requires breaking the network of hydrogen bonds. Also explains why pond-skaters can walk on water.