LSS_States of Matter Lesson
States of Matter
The states of matter differ fundamentally in how much space is between the particles as shown in the diagram below.
There are some other basic properties of solids, liquids, and gases with which you are probably already familiar. Gases expand to fit their container, solids and liquids do not. Solids and liquids maintain their volume. Solids keep their shape while liquids and gases transform to the shape of the container. Gases are easily compressed while solids and liquids are not.
In this module, we will study the other properties of liquids and solids but will study gases in a separate module. There is a specific reason for this. The particles that make up liquids and solids are so close that they interact with each other quite often. In gases, the molecules are so far apart that the attractive forces are almost negligible, so any differences between the attractive forces are so small that they hardly matter at all. As a result, chemical composition has little effect on the properties of a gas. But in a liquid or a solid, the molecules are close together and the attractions are strong. Differences among these attractions caused by differences in chemical makeup are greatly amplified, so the properties of liquids and solids depend quite heavily on chemical composition whereas they do not for gases.
Intramolecular Versus Intermolecular
In the last module we studied bonding. Bonding could also be called intramolecular attraction. To see what this means, think of the prefix "intra." For example, intramural sports take place among two teams within the same school. Intramolecular attractions refer to the attraction within the same molecule (in other words what holds the molecule together in the first place). The reason for referring to bonding by this term is to now introduce a new term, intermolecular attraction. Be VERY careful reading these terms as they look so much alike!
To understand what intermolecular attraction means, think of the prefix "inter." For example, international means between nations. The interstate is a highway between two states. Intermolecular attractions (abbreviated IMA) refer to the attractions between molecules. This is something we have alluded to but never discussed.
Bonding, or intramolecular attraction, describes the attraction of the atoms within the molecule. Intermolecular attraction describes the attraction between molecules or compounds once they are already formed. Intermolecular attractions are always much weaker than the intramolecular forces.
For example, in HCl the hydrogen and chlorine atoms are held together by a strong intramolecular attraction, a covalent bond. The strength of this bond affects the chemical properties of HCl. When one HCl molecule moves near another HCl molecule, there will be a brief attraction between the neighboring molecules. These brief attractions are the intermolecular attractions. They are what determine the physical properties of liquids and solids.
Types of Intermolecular Attractions
Collectively, intermolecular attractions are called Van der Waals forces. There are several types of these attractions. We will discuss them in order of strength, starting with the strongest.
Dipole - Dipole Attractions
The strongest of the Van der Waals forces are dipole-dipole attractions. You learned in the last module that polar bonds can create polar molecules if they are shaped in a way that creates a dipole moment. Recall that this means a polar molecule has a partial positive side and partial negative side. So, molecules with dipole moments can attract each other electrostatically by lining up so that the + and - ends of neighboring molecules are close to each other.
The strongest of the Van der Waals forces are dipole-dipole attractions. You learned in the last module that polar bonds can create polar molecules if they are shaped in a way that creates a dipole moment. Recall that this means a polar molecule has a partial positive side and partial negative side. So, molecules with dipole moments can attract each other electrostatically by lining up so that the + and - ends of neighboring molecules are close to each other.
Although a dipole-dipole attraction is the strongest of the intermolecular attractions, it is still a very weak force. As the molecules move farther apart from each other, this weak force becomes even weaker. For molecules of approximately equal mass and size, the strengths of dipole-dipole attractions increase with increasing polarity.
Hydrogen Bonds
Hydrogen bonding is a specific type of dipole-dipole attraction. Do not be fooled by the name. This is not another type of bond. It is an intermolecular attraction, not a bond. In fact, you should think of this "hydrogen attraction" instead of hydrogen bonding. We will continue to use the term hydrogen bond, as that is the term using in the scientific community. However, every time you see it, stop and say to yourself, "This is not a bond. This is an IMA." H bonds occur among molecules in which hydrogen is bonded to fluorine, oxygen, and nitrogen, small, very electronegative atoms. This occurs for two reasons. The difference in electronegativity is very great, and the partial charges that occur are very concentrated into a small molecule, allowing the charged ends of the molecule to be very close to the oppositely charged end of another molecule. Hydrogen bonds are anywhere from 5-10 times stronger than other dipole-dipole attractions.
London Forces
Attractions can also occur between nonpolar molecules. These are called London forces (or London dispersion forces). This is a bit more complicated to understand because these particles are neutral therefore you would expect no attraction. Since all molecular particles move, the movement of the electrons in each individual molecule or atom can influence the movement of a neighboring molecule or atom because electrons will repel each other. Molecules or atoms can develop a momentary non-symmetrical electron distribution that produces a temporary charge called an instantaneous dipole.
When an instantaneous dipole is created in one molecule, it causes the electrons to respond in a neighboring molecule. This is called an induced dipole. Both instantaneous and induced dipoles are very short lived because the electrons are constantly moving, creating new instantaneous and induced dipoles.
The strength of London forces are determined by:
- Size of the electron cloud
- Number of atoms in a molecule
- Molecular shape
Size of the electron cloud
Electron clouds that are large are more easily polarizable and therefore have stronger IMA. The outer electrons are further from the nucleus of the atom, which results in being held more loosely to the nucleus. This allows the electron cloud to change shapes easily, making the London forces stronger than those of atoms with smaller electron clouds.
Number of atoms in a molecule
Molecules that have the same elements in them experience increased London forces when there are more atoms in the molecule than a molecule made of the same elements, but with fewer atoms. For example, hydrocarbons (organic compound composed of hydrogen and carbon) with fewer hydrogens and carbons experience less London forces than those with more hydrogens and carbons.
Molecular shape
The last factor affecting London forces is molecular shape. Molecules that have the exact same molecular formulas show differences in London forces, due to the differences in their shape.
All molecules attract to each other, but there are different patterns and strengths of attraction that occur.
Ions are also affected by London forces. Ions can be attracted to polar molecules. This is called ion-dipole attraction. Ions are also able to create dipoles in molecules. This results in ion-induced dipole attractions. Because the charge of an ion is permanent, ion-induced dipole attractions can be stronger than regular London forces that are created by instantaneous-dipoles.
Remember to work on the module practice problems as you complete each section of content.
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