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Chemistry I · original Chapter 6 · combined Chapter 6
Section 4
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In the previous section, we learned what were the driving forces that could hold molecules together. But now that we have a better understanding of that, it is just as important to see the possible applications behind intermolecular forces and why some materials behave in certain states.
Solids
There are many ways to characterize solids, but two of the most classic classifications include two categories. Crystalline solids are solids whose atoms have a structured arrangement, while amorphous solids have a very disordered structure.
Examples of solids that have crystalline structures are those beautiful colored crystals such as amethyst. The “crystal” look that it has is only achieved because of the way the atoms/molecules are arranged in three dimensional space. The position of these components are usually represented by what is known as a lattice, which is a three dimensional rendering of points designating the center of the component (usually an atom or molecule) and also showing the repetitious nature of that component. The smallest unit of a lattice is known as a unit cell. It shows what a typical repeating component in a lattice structure looks like which is shown in the next page:
Gallery 1.6: Unit Cell and Lattice Structure
Here is what a typical unit cell looks like of a crystalline solid
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Image 1
Here is what a typical unit cell looks like of a crystalline solid
Image 2
A look at what a lattice looks like. As you can see it is a combination of multiple unit cells.
Solids retain shape for different structural reasons. Molecular solids have discrete molecules held together by intermolecular attractions, including dispersion and sometimes dipole interactions or hydrogen bonding. Ionic solids contain extended ion lattices; metals have extended metallic bonding; covalent-network solids have extended covalent bonds. A unit cell describes the repeating structure, not a general collection of molecules held by London forces. Mass alone does not explain every solid’s shape or strength.
Related meniscus note
Water on clean glass can have a concave meniscus when adhesion to the surface is strong relative to its cohesive interactions. The behavior depends on both liquid and surface; intermolecular forces do not give every liquid the same meniscus shape.
Liquids
Liquids behave very differently from solids be it that they can’t really retain a shape besides the container that they hold. It is important to consider though the intermolecular forces that are associated with liquids. For example, water is held very closely together by hydrogen bonding, which is why water has such a high boiling point when compared to other liquids. The water molecules can arrange themselves to the point where the IMF forces are strong enough to hold the molecules together.
Liquids can behave in interesting ways, one of the most common applications of intermolecular forces in liquids can be seen in your very classroom perhaps; when you pour water into a graduated cylinder. You notice that when you do this, the water tends to form some sort of concave shape and you are taught that this is usually where you measure the water level. The reason that this happens is because of intermolecular forces. If you think about it, the water at the surface is only affected by the water molecules surrounding it on it’s sides and bottom as shown in the figure.
Interactive 1.7: Intermolecular Forces at the Surface
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Hydrogen bonding
The strong intermolecular forces cause the water molecules on the surface to be pulled towards the inside of the liquid
Surface tension can support small organisms at an air–water interface. A gecko clinging to a dry wall uses a different mechanism: its many tiny foot contacts bring surfaces close enough for dispersion attractions to contribute to adhesion. Do not use wall climbing as an example of standing on a stretched liquid surface.
Gases
You can imagine then that gases have very weak intermolecular forces as they do not really have any large attractions to each other to be able to form tangible things that you can see and touch. This is why, it is easy for gas molecules to move around and is also why you don’t see them with the human eye. We will cover gases later on in more depth in a later chapter but for now that’s all you’ll really need to know.
Changes of State
Since we’ve already covered solids, liquids and gases at this point and how they are different in terms of intermolecular forces, it is important to consider how exactly do they change state and how that relates to intermolecular forces. This is mainly done by looking at a heating diagram. This diagram outlines how an object can change states at various temperatures as a function of time. Look at the figure on the next page. It’ll help explain things a little more.
Gallery 1.7 Carbon Nanotubes
This is what they typically look like. Very small tubes, with about a sixth of the weight of a steel wire
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Image 1
This is what they typically look like. Very small tubes, with about a sixth of the weight of a steel wire
Image 2
Nanotubing is made up of carbon atoms that are interconnected as chains of a six-membered ring as shown
The flat lines are important to consider as well. They tell us that over a given period of time, the object enters a sort of transition phase where it is both a solid and a liquid or a liquid and a gas. At these points, the energy is not going in to raising the temperature, but instead going into breaking the intermolecular forces that are held up between each molecule.
Interactive 1.8: Heating Diagram for Water
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Solid to Liquid
Flat line indicates a change in state from solid to liquid. Energy is being added to break the intermolecular forces of the solid to turn it into a liquid
Notice the Plateaus
Flat lines indicate that energy is being added into the water to change it from water to steam
Within a carbon nanotube, carbon atoms form a covalently bonded network. Its ring pattern does not make those connections intermolecular forces. Separate tubes can attract one another through dispersion. Distinguish bonding inside each structure from attractions between separate structures when explaining strength or bulk properties.
A big problem with carbon nanotubes however is that when they are manufactured, they come in a variety of sizes which makes it very difficult to sort.
Overall it is important to realize the role that intermolecular forces and atomic interactions play a role in everyday applications. They are the driving forces for why objects behave the way they do and why we have different states of matter. It’s going to get cheesy for a minute but it is the reason life is so...well complicated!
Glossary and deeper reading · 4 refreshers
These editorial refreshers supplement the reading. Definitions and examples are added guidance; original source wording and recorded answers remain in their original context.
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Intermolecular force
Intermolecular forces are interactions between molecules. They affect properties such as boiling point and solubility and are distinct from the bonds within each molecule. For example, vaporizing water separates molecules while largely preserving their O–H bonds.
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Crystal lattice
A crystal lattice describes the repeating arrangement of structural sites in a crystalline solid. The occupants can be ions, atoms, or molecules depending on the material. For example, an ionic lattice places opposite charges near one another throughout a sodium chloride crystal.
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Hydrogen bond
Hydrogen bonding is a directional interaction involving hydrogen bonded to a sufficiently electronegative atom and an appropriate electron-rich acceptor. In the introductory model, N, O, and F commonly provide the relevant partners. For example, water molecules hydrogen bond when one molecule's O–H group interacts with another molecule's oxygen lone pair.
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Ion
An ion is an atom or group of atoms with a net electric charge. Ordinary ion formation changes electron count without changing the element's proton count. For example, a sodium atom losing one electron becomes a sodium ion with charge +1.