Molecular Solid

So I threw a lot of definitions at you at once, so let's break this down piece by piece. Let's first talk about the crystalline lattice. Below is what a crystalline lattice looks like:

Fig.1 Example of a crystal lattice versus amorphous solid

A crystal lattice is essentially just a repeating pattern where each atom/molecule is being held together by some force. These forces also determine how the lattice will be structured.

For molecular solids, these forces are van der Waals forces. There are three main van der Waals forces, and each of the three types of molecular solids corresponds to one of these forces. The main “point” of van der Waals forces is that they are caused by the attraction of a partial negative charge (δ-) to a partial positive one (δ+). How these charges occur is dependent on the type of force, which we get into just a bit.

Types of Molecular Solids

As I just mentioned, there are three types of molecular solids that correspond to the three types of van der Waals forces. These are:

• Non-polar solids.
• Polar solids.
• Hydrogen-bonded solids.

Non-polar Solids

First up are non-polar solids. When a species is non-polar, it is for one of two reasons:

1. The molecule is symmetrical, canceling out any polarity.
2. The difference in electronegativity is less than 0.5.

Non-polar solids are held together by London dispersion forces

Below is what this process looks like:

Now that the species has an instantaneous dipole, it can induce a dipole in a nearby non-polar species. The electrons in the non-polar species are being pulled towards the partially positive end of the instantaneous dipole. This causes the electrons to become unevenly distributed and form a dipole.

Here is what this process looks like:

The attraction between the instantaneous dipole and the induced dipole is what we consider to be London dispersion forces. Induced dipoles are considered temporary since they will disappear when moved away from a molecule with a dipole.

Polar Solids

Next, we have polar solids, which are held together by dipole-dipole interactions.

Hydrogen Bonded Solids

Hydrogen-bonded solids are, you guessed it, held together by hydrogen bonds.

Hydrogen bonding is a special type of dipole-dipole interaction. It is actually much stronger due to the large difference in electronegativity between H and N, O, or F. Below is an example of a hydrogen bond:

Fig.5-Hydrogen bonding in ice

Molecular Solid Examples

Now that we've covered each type, let's see what these solids look like!

First, we have a non-polar molecule solid. Below is the structure for iodine (I₂). It is a non-polar molecule, so it is being held by London dispersion forces

Fig.6-Iodine is an example of a non-polar molecular solid

Iodine's crystalline structure is a cube. There is an iodine molecule on each corner and the center of each face. Unlike with our other examples, there is no line drawn for the intermolecular forces. This is because there is no permanent dipole on any of the molecules, so the direction of attraction will change as instantaneous dipole are formed and destroyed (and therefore the induced dipoles of nearby molecules).

Let's zoom in an look at what these intermolelcular forces would look like:

Fig.7-Iodine dispersion forces

One iodine molecule forms a temporary dipole, which induces a dipole in a nearby molecule. Once the dipole is destroyed a new one can be formed.

These interactions are strong enough to hold the iodine structure together, but since they are weaker than the covalent bond within each individual iodine molecule, each molecule is spaced farther apart. Iodine is shaped like a cube since the strength of these interactions is equal.

Next, let's look at a polar solid. Our example below is HCl:

Fig.8 HCl is a polar solid

Its solid structure is a zigzag, where each molecule is being held together by dipole-dipole interactions.

The most common hydrogen-bonded solid is ice, shown below:

Fig.9-The crystalline structure of ice

Calling back to the introduction, this is why snowflakes are hexagonal. Snowflakes are unique because of the different hydrogen-bonding patterns, but will always have that core hexagonal shape.

Properties of Molecular Solids

Molecular solids get their properties from the weak van der Waals forces that hold them together. Here are some of the common properties of molecular solids:

• Soft
  • Weak forces --> easy to deform
• Low melting point
  • Weak forces --> easy to overcome
• Low density
  • Density is a measure of mass per volume. The intermolecular “bonds” are long, so the space between molecules is great
• Poor conductors of electricity
  • The structure prevents electron movement
• Poor thermal conductors
  • The structure is too far apart

Depending on the strength of the van der Waals forces, these properties are more or less accurate. For example, water has a relatively high melting point since it has stronger hydrogen bonding

From weakest to strongest:

• Weakest: London dispersion forces.
• Dipole-dipole.
• Strongest: Hydrogen bonding.

Molecular Solid Conductivity

Let's dive into why molecular solids are poor conductors (and therefore good insulators).

To understand why molecular solids are bad electrical conductors, let's look at what makes a good conductor.

Conductors are basically “electron highways”. They allow electrons to flow freely through them. Molecules that are good conductors allow their electrons to move freely.

Molecular solids are composed of neutral molecules, so they have no free electrons. This means that electrons cannot flow freely, so they are poor conductors.

They are also poor conductors of heat. Heat conductivity is a “passing” of heat energy from one particle to another. Think of it like passing a ball across a line of people. If the people are shoulder-to-shoulder, it doesn't take much time to pass the ball along, so many balls can be passed along quickly. If the people are standing a few feet apart, it takes a lot more time, so it is much less efficient.

Molecular solids have weak van der waals forces holding them together, so they are relatively far apart. This means the passage of heat takes much longer.

Below is an example:

Fig.10-Difference in conductivity based on intermolecular forces

Ionic solids like sodium chloride (NaCl) have strong electrostatic forces holding them together, while molecular solids like iodine have weak van der waals forces.

These weaker forces hold molecules farther apart, which is why they are poor conductors.