molecular solids e.g. sugar (sucrose), carbon dioxide, ammonia, iodine
covalent network e.g. diamond (3D), silica/silicondioxide (3D), and graphite (2D)
ionic e.g. sodium chloride, magnesium chloride
metallic e.g. copper, magnesium

We will learn more about each type of solid.

Watch the Crashcourse video below for an introduction to solids.

Molecular Solids

Molecular solids are made of separate molecules packed together. The molecules are usually made up of non-metal atoms bonded together. Exceptions include chlorophyll and haemoglobin molecules which contain magnesium atoms and iron atom respectively. Molecular solids can be molecules of elements e.g. iodine, or can be molecules of compounds e.g. carbon dioxide, sugar (sucrose).
Molecules can be either polar or non-polar. See the previous page for information about this.
Molecular substances which are usually gases or liquids at room temperature and that have been cooled to become solids are called molecular solids. For example ice is a molecular solid made up of frozen water molecules, and dry ice is a molecular solid made up of cooled/frozen carbon dioxid gas.
Separate molecules are linked by weak forces of attraction between molecules called intermolecular forces. These are sometimes referred to as van der Waals forces. This is shown in the diagrams below.
Within the molecules the atoms are held together by strong covalent bonds. These are called intramolecular forces/bonds. This is shown in the diagrams below.
Due to the weak intermolecular bonds, molecular solids have a low melting point, low boiling point and are generally soft. However, the larger the molecule, the larger the van der Waals forces and then the higher the melting and boiling points.
Some molecular solids are made up of polar molecules, such as hydrogen chloride. These solids have a slightly higher melting and boiling point due to the electrostatic attraction between the molecules due to dipoles, as well as van der Waals forces.
If the molecular solid contains the more electronegative nitrogen, oxygen or fluorine (NOF atoms) bonded to a hydrogen atom then the bond formed is called a hydrogen bond, Hydrogen bonds are the strongest type of intermolecular force of attraction.
Molecular solids do not conduct electricity since they do not contain any free electrons or ions.
As we discussed on the previous page, the solubility of molecular solids depends on the polarity of the molecular solid and the polarity of the solvent. Remember, like dissolves like. Non-polar molecular solids will dissolve in non-polar solvents. Whereas, polar molecular solids will dissolve in polar solvents, such as water.

Molecular Solids Questions To Answer

1) 'Dry ice' is solid carbon dioxide. 'Dry ice' can be used to produce artificial 'smoke' in theatrical stage productions because it sublimes easily (turns directly from solid to gas). Using your knowledge of the structure and bonding in solid carbon dioxide explain why it sublimes easily.

2) The masses of the two molecules S8 (sulfur molecule) and I2 (iodine molecule) are very similar. What would you expect to see about the melting points of these two solids? Which one would have the higher melting point? Explain your answer.

3) Both bromine, Br2 and chlorine, Cl2 form diatomic molecules (molecules made up of two atoms). When they are both in the solid state which one will have the highest melting point? Explain your answer.

4) This question relates to hydrogen forming compounds with the group 16 elements oxygen, sulfur, selenium and tellurium. In the series H2O, H2S, H2Se, and H2Te, the molecular masses of the molecules increase as you go from H2O to H2Te. However, water has the highest melting point of all four. Explain why this is.

(Answers are below)

SciPAD questions

Now do the questions on pages 70-75 in your Level 2 Chemistry External SciPAD Workbook.

Molecular Solids Questions To Answer

Ionic Solids

Ionic compounds contain positively charged cations and negatively charged anions. For example sodium chloride contains Na⁺ ions and Cl^- ions.
The ions are arranged in a lattice structure. Each positive ion is surrounded by a fixed number of negatively charged ions. In the case of sodium chloride each sodium ion, Na⁺, is surrounded by six negatively charged chloride ions, Cl^-, and each chloride ion is surrounded by sodium ions. So you can see there is a 1:1 ratio of ions in the case of sodium chloride.
The positive and negative ions are not the same size. In the case of sodium chloride, chloride ions are larger than sodium ions.
Strong electrostatic forces of attraction called ionic bonds exist between the positive and negative ions. These hold the ionic solid together.
Since the electrostatic forces of attraction (ionic bonds) holding the ions together in a rigid lattice are strong, ionic solids are strong.
Despite their strength, ionic solids are brittle. The lattice structure becomes brittle and fractures if the lattice is deformed. This is because the ionic bonds (electrostatic forces of attraction) are directional. If a sideways force is applied and a section of the lattice slides, then ions of the same charge may come in close contact with one another and repel. The ionic solid will then break.
The ions in an ionic solid are held firmly in position and are not free to move (there are no mobile charge carriers). Therefore, ionic solids do not conduct electricity. However, if the ionic solid is molten (melted) or if it is dissolved in water, the ions are free to move and a molten or aqueous solution of an ionic solid will conduct electricity.
Many ionic solids dissolve in polar solvents, with water being the most effective polar solvent. Dissolving occurs when small numbers of polar water molecules align themselves with the positive and negative ions, pull them out of their crystal lattice and surround them (see previous page).
Some ionic substances are insoluble in water because the ions in the solid lattice are held together so firmly that even a group of water molecules can't exert enough attractuion to pull the ions out of the lattice.

Ionic Solids Questions to Answer

1) Why do ionic solids conduct electricity when molten but not when solid?

2) Many ionic solids are white crystalline solids at room temperature. Which ionic solid, magnesium oxide or potassium chloride would have the higher melting point? Explain your answer.

3) Which of the two ionic solids in the question above would you expect to be most soluble in water? Explain you answer.

Metallic Solids

Metallic solids are made up of one type of metal atom. For example, copper is made up of copper atoms, aluminium is made up of aluminium atoms.
Metals are usually solid, dense, hard, shiny, malleable, ductile, good conductors of electricity and heat, with relatively high melting points and boiling points. There are some exceptions to this:
Mercury is the only liquid metal at room temperature
Group 1 metals (lithium, sodium, potassium) have relatively low densities and melting points. They can also be cut with a knife and are soft.
Metallic bonds exist between metal atoms. These are strong non-directional attractive forces between the positive ions and the negative de-localised electrons.
The presence of strong metallic bonds means that metallic solids generally have high melting points. Iron has a melting point of 1 535 degrees Celsius. However, group 1 alkali metals, mercury and gallium have much lower melting points.
Metallic solids have a regular arrangement consisting of positively charge metal cations whose valence electrons are free to move through the structure. The electrons are de-localised. You might see this described as "positive ions surrounded by a 'sea' of mobile electrons".
Metallic solids are a lattice of positively charged cations surrounded by delocalised electrons so there are mobile charge carriers available. The delocalised electrons readily move through the metallic lattice when an electric potential is applied and can therefore conduct electricity in a solid state, or when molten.
The positive cations in the metallic structure repel each other but the 'free' or mobile electrons hold them together.

Metallic solids are insoluble in polar and non-polar solvents. Some metals, e.g. sodium and other Group 1 alkali metals react with water to produce a hydroxide and hydrogen gas.
Since the metallic bonds are non-directional, if pressure is applied to the metallic solid, then layers of the solid can 'slide' over each other. Since the strong metallic bonds are non-directional the metallic particles still remain strongly bonded. This gives the metallic solid the property of being ductile (being drawn into wires) and malleable (being moulded/hammered into sheets).

Video - What Are Metallic Bonds?

Giant Covalent Network Solids

The three covalent network solids you need to really know about and understand are:- diamond, graphite and silica.
In covalent network solids the atoms occupy regular lattice positions. Every atom is bonded to a number of other atoms.
A great deal of energy is needed to overcome the strong covalent bonds within the solid and separate the atoms from their lattice position.
Each crystal in diamond and silica (quartz) is like one big molecule with continuous bonding throughout the structure.
These solids are insoluble since water molecules cannot exert enough attraction to remove atoms from their lattices and surround them.
The specific structures of the three covalent network solid account for the range of properties they have.

Diamond

Diamond is made up of only carbon atoms.
Each carbon atom is joined to four other carbon atoms by strong covalent bonds giving a continuous three dimensional crystalline structure.

Diamond can cut things eg paper and other solids. This is because of its very strong three dimensional lattice arrangement of carbon atoms, each joined to four other carbon atoms by strong covlent bonds.
Diamond is more dense than graphite because each carbon atom in the diamond lattice is covently bonded to four other carbon atoms in a compact, 3D array. The distance between the bonded carbon atoms is small and this means that diamond is quite dense.
Diamond is also a good abrasive due to its very strong 3D lattice of carbon atoms.
Diamond is an electrical insulator and does not conduct electricity. Since every carbon is bonded to three other carbon atoms then there are no 'free' valence electrons, so no mobile charge carriers. Hence diamond does not conduct electricity.
Diamond has a very high melting point since it has a very strong 3D lattice bonded by strong covalent bonds. Much heat energy is needed to break the bonds in the crystal lattice and so the melting point is high.

Graphite

Graphite is also only made up of carbon atoms.
In graphite each carbon atom is held together by strong covalent bonds to three other carbon atoms. It forms two dimensional layers of hexagonal rings of carbon atoms. The layers are held together by weak intermolecular forces. Delocalised electrons that are free to move are found between the layers.

The structure of graphite accounts for a number of its properties.

Graphite is used in pencils (as pencil 'lead') since it leaves a mark on paper. This is because it is made of two-dimensional layers of carbon atoms, arranged in hexagonal rings, which can easily slide over each other. This is made possible because there are weak intermolecular forces between the layers of hexagonal rings of carbon. This also explains why graphite is used as a dry lubricant.
Graphite is a good electrical conductor. In graphite, each carbon atom only uses three of its four available electrons to form bonds with other carbon atoms. This leaves one electron delocalised or 'free' to move and carry electrical charge when an electric potential is applied.
Graphite is less dense than diamond because in graphite every carbon atom is only bonded to three other carbon atoms in flat sheets. The distance between these sheets of hexagonal rings is relatively large compared to the distance between the bonded carbon atoms due to the weaker intermolecular forces. This means graphite is less compact than diamond and therefore, less dense.

Silica

Silica (quartz) is made of silicon and oxygen atoms arranged in a lattice. The ratio of silcon atoms to oxygen atoms is one to two, giving an empirical formula of SiO₂. Each silicon atom is covalently bonded to four oxygen atoms and each oxygen atom is covalently bonded to two silicon atoms.