IFAP - Intermolecular Forces (Lesson)
Intermolecular Forces
Introduction
Intermolecular forces explain the physical properties of a molecular compound.
Intermolecular forces, as the term implies, are attractions that exist between different molecules. These intermolecular forces (IMFs) are those that exist beyond the forces that hold atoms together in a molecule. They are the attractions that attract molecules together and are responsible for many chemical properties that will be discussed throughout this module.
What You Already Know
In the previous module, you explored the VSEPR model of molecular structure and the unequal sharing of electron pairs in polar covalent bonds. These two concepts account for polar molecules where the electronegativity differences of the covalent bonds cause the molecule as a whole to have regions of high electron density and regions with low electron density creating a polar molecule. Depending on the manner in which polar chemical bonds are arranged in three-dimensional space it is possible to have molecules that are comprised of polar chemical bonds only yet remain nonpolar as a whole.
How Do Polar and Nonpolar Molecules Account For the Physical Properties of a Substance?
Consider a group of HCl molecules. These are polar molecules with the dipole shown in the diagram below:
The figure above shows two HCl molecules that are in close proximity to one another. Due to the large electronegativity difference between the hydrogen and the chlorine, there is an uneven distribution of electrons causing the chlorine to have a partial negative charge while the hydrogen is partially positive. Due to the uneven distribution of charge between the atoms, the molecule is said to be polar in nature.
When one HCl molecule approaches another, there will be an attraction between these neighboring molecules due to this dipole. These attractions are the intermolecular attractions. When the hydrogen of one HCl molecule approaches the chlorine of another, the partial charge across the bond from the unbalanced sharing of electrons between H-Cl inside the molecule creates forces between these molecules. These forces between molecules are an example of an intermolecular force and help determine the physical properties of liquids and solids.
Intermolecular Forces
There are four categories of intermolecular forces:
- ion-ion forces
- ion-dipole forces
- dipole-dipole forces
- H-bonding forces
- dispersion forces
Ion-ion Forces
Ion-ion forces are essentially the forces of attraction that are present in ionic compounds. In these substances, there are fully positive charges (cations) and fully negative charges (anions). The Coulombic force of attraction between these ions is simply referred to as ion-ion forces. They are the strongest type of intermolecular force due to the fact that the ions are fully charged (as opposed to being only partially charged as in a polar covalent bond).
Ion-Dipole Forces
These forces typically only exist in solutions of ionic compounds dissolved in a polar solvent. One of the most common solvents, water, is a great example. Let's consider a solution of sodium chloride (NaCl) dissolved in water. As illustrated in the figure below, there is a permanent dipole in a water molecule resulting in partially positive and partially negative portions of the molecule. The cations and anions from sodium chloride will interact in different ways. A force of attraction exists between the positively charged sodium (ion) and the partially negative portion of the water molecule (dipole). Hence the name ion-dipole force. Likewise the negatively charged chloride will interact in a similar fashion with the partially positive portion of the water molecule.
Dipole-Dipole Forces
A third type of intermolecular force is dipole-dipole forces. This type of force is described above using HCl as an example. With dipole-dipole forces, there are no fully positive or negative ions, only partially charged atoms due to uneven sharing of electrons. Because of this, the dipole-dipole force is typically, though not always, weaker than ion-ion or ion-dipole forces.
Hydrogen Bonding
Hydrogen bonding is actually a special case of a dipole-dipole force that is particularly strong in nature. In specific cases of O−H, N
−H, or F
−H, the dipole created between these atoms is so strong that it warrants their description as a separate type of force altogether. It must be stated that only these three types of bonds can participate in hydrogen bonding. Other hydrogen containing bonds such as C
−H, S
−H, etc. are not polar enough to engage in hydrogen bonding.
Water is a classical example of a substance that exhibits hydrogen bonding. The strong attraction between its molecules gives water high boiling and melting points compared to other nonpolar molecules of similar size. This strong IMF is why water has a large heat capacity and arranges itself into crystals in the solid phase. Ammonia (NH3) and hydrogen fluoride (HF) are two other molecules that experience Hydrogen IMFs.
Dispersion Forces
The term dispersion forces is a bit of a misnomer. The idea of things "dispersing" implies that they are being repelled by each other. Since this is not the case and dispersion forces are actually attractive forces like all other IMFs, what exactly is being "dispersed"? Consider an atom of Ne with 10 electrons. In isolation an atom of Ne would be perfectly nonpolar; however, when two Ne atoms come in close contact with one another as shown below, there is a temporary dipole that is created. This instantaneous dipole is caused when the electrons from one atom of Ne repel the electrons in the second Ne causing very slight, very weak dipoles in the atom. These dipoles can then attract one another in a similar fashion as dipole-dipole forces described above. It is referred to as an instantaneous dipole because it only remains in existence for a brief instant. When the two atoms move away from each other the electrons spread out again and the dipole is destroyed. It is important to note that every substance exhibits this type of IMF.
Summary of Intermolecular Forces
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