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Structure (Topic 2B)
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Chemicals be classified as simple molecular covalent structures or giant lattice structures.
In simple molecular covalent structures the exact number and type of atoms within a molecule are defined. Examples of chemicals which exhibit simple molecular covalent structures are water (ice), sulfur iodine and nanoparticles such as A lattice of simple molecules is held together by intermolecular forces.
Giant structures can be ionic metallic or covalent They have fixed ratios of elements but the overall structural size is not set. In a giant lattice there are strong bonds throughout the structure.
Sodium chloride is a common example of a giant ionic lattice.
In this structure, each sodium ion is surrounded by six chloride ions , and each chloride ion is surrounded by six sodium ions.
and ions alternate in a three-dimensional lattice structure.
Metallic bonding arises from the electrostatic forces of attraction between positive metal ions and a sea of delocalised electrons.
The bonding in magnesium is metallic bonding.
Lattice of positive ions: Magnesium atoms lose their two outermost electrons to form ions, which are arranged in a regular lattice structure.
Electrostatic attraction: Strong electrostatic forces exist between the positively charged ions and the delocalised electrons, holding the structure together.
Delocalised electrons: The lost electrons become free-moving (delocalised) within the metallic lattice, creating a ‘sea of electrons’.
Diamond has a giant covalent lattice structure.
Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement, forming a rigid 3D network.
Iodine molecules, , consist of two iodine atoms joined by a strong covalent bond. The bonds hold the molecule together but do not hold the lattice together in the solid.
In the solid state, molecules arrange into a molecular lattice held together by London dispersion forces.
These weak forces arise due to temporary dipoles induced in the molecules.
This results in a crystalline solid that can easily sublimate, converting directly from solid to gas upon heating.
Ice is an example of a molecular lattice stabilised by hydrogen bonding.
Each molecule forms hydrogen bonds with four other molecules, creating an extensive network.
The regular, open lattice structure of ice is responsible for its crystalline nature in the solid state. The spacing of the molecules in the lattice means that ice has a lower density than liquid water, which is why ice floats.
Diamond has a giant covalent lattice structure.
Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement, forming a rigid 3D network.
Graphite has a giant covalent structure.
- It consists of layers of carbon atoms arranged in a hexagonal lattice.
- Within each layer, each carbon atom is covalently bonded to three other carbon atoms, forming planar sheets.
- The layers are held together by weak London dispersion forces, allowing them to slide over one another.
- Covalent bonds within the layers are strong, involving shared electron pairs.
- Each carbon atom has one delocalised electron, which is free to move throughout the layers.
Graphene is a single layer of graphite and, therefore, also consists of carbon atoms bonded to three other carbon atoms arranged in hexagons.
As graphene consists of one layer of graphite, it is only one atom thick and is considered two dimensional.
One electron per carbon atom is unpaired and not involved in covalent bonding. These electrons become delocalised and free to move along the plane of the structure.
Where electrons are free to move there is electrical conductivity.
The physical properties of substances can be used to identify their structure and bonding.
‘High’ melting and boiling points are generally over 200 ℃ but can be significantly higher. For example, diamond exists in a solid state until above 3500 ℃.
Giant covalent lattices always have high melting points and boiling points due to the many strong covalent bonds between atoms in the lattice, which require lots of energy to break.
Most giant covalent lattices are unable to conduct electricity as all the outer shell electrons are held in position within covalent bonds.
There are exceptions: graphite and graphene only bond to three other carbon atoms, leaving one electron to become delocalised per carbon atom.
These electrons are free to move throughout the whole structure and therefore current can flow.
Giant covalent lattices are insoluble in all solvents.
The covalent bonds in a giant covalent lattice are significantly stronger than any interaction between the atoms in the lattice and a solvent.
It is not energetically favourable to break lattice bonds in order to interact with solvent molecules.
Giant covalent lattices are generally strong, but strength can vary depending on their structure.
Diamond is hard because of its rigid tetrahedral arrangement of carbon atoms.
Graphite is soft because it requires little energy to overcome the weak intermolecular forces between layers.
The physical properties of diamond can be explained using its structure:
- Hardness: The extensive network of strong covalent bonds makes diamond extremely hard.
- High melting point: A large amount of energy is required to break the strong covalent bonds throughout the structure.
- Electrical insulation: Diamond does not conduct electricity because there are no delocalised electrons or free ions.
The physical properties of graphite can be explained using its structure:
- Soft and slippery: The weak London dispersion forces/weak Van der Waals forces between layers allow the layers to slide, making graphite a good lubricant.
- Conductivity: The delocalised electrons can move freely throughout the layers, allowing graphite to conduct electricity.
- High melting point: Strong covalent bonds within the layers require a significant amount of energy to break.
- Insolubility: Graphite does not dissolve in solvents due to the strength of the covalent bonds within the layers.
Graphene is an excellent conductor of electricity.
Combined with its thin structure and its exceptional strength, graphene is a critical component in small-scale electronics.
Graphene has exceptional strength.
The individual bonds in graphene are shorter and stronger than those in diamond.
Combined with the regular structure across the layer, graphene has exceptional tensile strength.
Giant metallic lattices usually have high melting points and boiling points.
This is due to the strong electrostatic forces of attraction between the positive metal ions and delocalised electrons, which require a lot of energy to overcome.
The melting point of a giant metallic lattice will vary with the nature of the cation formed.
Larger cations with lower charges will exhibit weaker electrostatic attraction and a lower boiling point.
Metals can conduct electricity, even when solid, due to the delocalised electrons, which are free to move throughout the giant metallic lattice and carry charge.
Metals are insoluble.
There are interactions between polar solvents and metals. However, this leads to chemical reactions rather than dissolution.
Ionic compounds typically have high melting and boiling points.
This is due to the strong electrostatic forces of attraction between the oppositely charged ions in the giant ionic lattice.
A significant amount of energy is required to overcome these forces and change the compound from solid to liquid (melting) or from liquid to gas (boiling).
The ability of ionic compounds to conduct electricity relies on the ability of the ions to move and carry charge, which depends on the state they are in:
- In the solid state, ionic compounds do not conduct electricity because the ions are fixed in the lattice and cannot move freely, and therefore cannot carry charge through the giant ionic lattice.
- When molten, ionic compounds become good conductors of electricity. The ions move freely in the liquid state, allowing them to carry charge.
- Ionic compounds also conduct electricity when dissolved in water. The ions dissociate and move freely in the solution, allowing them to carry charge.