7 Intermolecular Forces

Learning Objectives

  • Use strength of intermolecular forces in organic molecules to explain differences in physical properties such as boiling point and solubility.
    | Intermolecular Forces |

| Key Concepts and Summary | Glossary |

Intermolecular Forces

Intermolecular forces (IMF) are the various forces of attraction that exist between atoms and/or molecules due to electrostatic properties. We will often use values such as boiling or melting points as indicators of the relative strengths of IMFs present within different substances. The stronger the IMF, the higher the melting and boiling points, and the lower the vapor pressure. The IMF (starting with the strongest) are:

  1. Ion-Ion: the force between two ions (example: NaCl(s))
  2. Ion-Dipole: the force between an ion and a polar molecule (example: Na+(aq) and H2O(l))
  3. Dipole-Dipole: the force between two polar molecules.
  4. Hydrogen Bonding: certain polar molecules containing F, O, or N also exhibit this strong IMF (example: H2O(l) with itself, H2O(l) and NH3(l))
  5. Dipole-Induced Dipole: the force between a polar molecular and a non-polar molecule (example: H2O(l) and C6H12O6(aq), table sugar)
  6. London Dispersion Forces, also known as Induced Dipole-Induced Dipole, van der Waal’s, or just Dispersion: the force between any two molecules that each have at least one electron.

Dispersion Forces

The electrons of an atom or molecule are in constant motion, so at any moment in time, an atom or molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically. The presence of this dipole can, in turn, distort the electrons of a neighboring atom or molecule, producing an induced dipole. These two rapidly fluctuating, temporary dipoles result in a relatively weak electrostatic attraction between the molecules, a so-called dispersion force, like that illustrated in Figure 1.

Two pairs of molecules are shown where each molecule has one larger blue side labeled “delta sign, negative sign” and a smaller red side labeled “delta sign, positive sign.” Toward the middle of the both molecules, but still on each distinct side, is a black dot. Between the two images is a dotted line labeled, “Attractive force.” In the first image, the red and blue sides are labeled, “Unequal distribution of electrons.” Below both images are brackets. The brackets are labeled, “Temporary dipoles.”
Figure 1. Dispersion forces result from the formation of temporary dipoles, as illustrated here for two nonpolar diatomic molecules.

Dispersion forces that develop between atoms in different molecules can attract the two molecules to each other. The forces are relatively weak and only become significant when the molecules are very close. Larger and heavier atoms and molecules exhibit stronger dispersion forces than those found in smaller and lighter atoms and molecules. F2 and Cl2 are gases at room temperature (reflecting weaker dispersion forces); Br2 is a liquid, and I2 is a solid (reflecting stronger dispersion forces).

The increase in melting and boiling points with increasing atomic/molecular size may be rationalized by considering how the strength of dispersion forces is affected by the electronic structure of the substance. In a larger atom, the valence electrons are, on average, farther from the nuclei than in a smaller atom. Thus, they are less tightly held and can more easily form the temporary dipoles that produce the attraction. The measure of how easy or difficult it is for another electrostatic charge (for example, a nearby ion or polar molecule) to distort a molecule’s charge distribution (its electron cloud) is known as polarizability. A molecule that has a charge cloud that is easily distorted is said to be very polarizable and will have large dispersion forces; one with a charge cloud that is difficult to distort is not very polarizable and will have small dispersion forces.

The shapes of molecules also affect the magnitudes of the dispersion forces between them. For example, boiling points for the isomers n-pentane, isopentane, and neopentane (shown in Figure 2) are 36 °C, 27 °C, and 9.5 °C, respectively. Even though these compounds all have the same empirical formula, C5H12, the difference in boiling points suggests that dispersion forces in the liquid phase for each are different, the greatest seen in n-pentane and the weakest in neopentane. The elongated shape of n-pentane provides a greater surface area available for contact between molecules, resulting in correspondingly stronger dispersion forces. The more compact shape of isopentane offers a smaller surface area available for intermolecular contact and, therefore, results in weaker dispersion forces. Neopentane molecules are the most compact of the three, offering the least available surface area for intermolecular contact and, hence, the weakest dispersion forces. This behavior is analogous to the connections that may be formed between strips of VELCRO brand fasteners: the greater the area of the strip’s contact, the stronger the connection.

Three images of molecules are shown. The first shows a cluster of large, gray spheres each bonded together and to several smaller, white spheres. There is a gray, jagged line and then the mirror image of the first cluster of spheres is shown. Above these two clusters is the label, “Small contact area, weakest attraction,” and below is the label, “neopentane boiling point: 9.5 degrees C.” The second shows a chain of three gray spheres bonded by the middle sphere to a fourth gray sphere. Each gray sphere is bonded to several smaller, white spheres. There is a jagged, gray line and then the mirror image of the first chain appears. Above these two chains is the label, “Less surface area, less attraction,” and below is the label, “isopentane boiling point: 27 degrees C.” The third image shows a chain of five gray spheres bonded together and to several smaller, white spheres. There is a jagged gray line and then the mirror image of the first chain appears. Above these chains is the label, “Large contact area, strong attraction,” and below is the label, “n-pentane boiling point 36 degrees C.”
Figure 2. The strength of the dispersion forces increases with the contact area between molecules, as demonstrated by the boiling points of these pentane isomers.

Dipole-Dipole Forces

Recall from earlier in this section that polar molecules have a partial positive charge on one side of the molecule and a partial negative charge on the other side—a separation of charge called a dipole. Consider a polar molecule such as hydrogen chloride, HCl. In the HCl molecule, the less electronegative H atom bears the partial positive charge (shown in red in Figure 12) and the more electronegative Cl atom bears the partial negative charge (shown in blue). An attractive force between HCl molecules results from the attraction between the positive end of one HCl molecule and the negative end of another. This attractive force is called a dipole-dipole attraction—the electrostatic force between the partially positive end of one polar molecule and the partially negative end of another, as illustrated in Figure 3.

Two pairs of molecules are shown where each molecule has one larger blue side labeled “delta sign, negative sign” and a smaller red side labeled “delta sign, positive sign. In the first pair, the red sides of the two molecules both face to the left and the blue side to the right. A horizontal dotted line lies in between the two. In the second pair, the molecules face up and down, with the red and blue ends aligning. A horizontal dotted line lies between the red and blue ends facing upward and another lies between the red and blue ends facing downward.
Figure 3. This image shows two arrangements of polar molecules, such as HCl, that allow an attraction between the partial negative end of one molecule and the partial positive end of another.

The effect of a dipole-dipole attraction is apparent when we compare the properties of HCl molecules to nonpolar F2 molecules. Both HCl and F2 consist of the same number of atoms and have approximately the same molecular mass. However, the dipole-dipole attractions between HCl molecules are sufficient to cause them to “stick together” to form a liquid, whereas the relatively weaker dispersion forces between nonpolar F2 molecules are not, and so F2 is gaseous at this temperature. The higher boiling point of HCl (188 K) compared to F2 (85 K) is a reflection of the greater strength of dipole-dipole attractions between HCl molecules, compared to the attractions between nonpolar F2 molecules.

Hydrogen-Bonding Forces

Nitrosyl fluoride (ONF, molecular mass 49 g/mol) is a gas (weak IMF) at room temperature. Water (H2O, molecular mass 18 g/mol) is a liquid (stronger IMF), even though it has a lower molecular mass. We clearly cannot attribute this difference of physical state between the two compounds to dispersion forces. Both molecules have about the same shape and ONF is the heavier and larger molecule. It is, therefore, expected to experience more significant dispersion forces. Additionally, we cannot attribute this difference in boiling points to differences in the dipole moments of the molecules. Both molecules are polar and exhibit comparable dipole moments.

The large difference between the boiling points is due to a particularly strong dipole-dipole attraction that may occur when a molecule contains a hydrogen atom bonded to a fluorine, oxygen, or nitrogen atom (the three most electronegative elements). The very large difference in electronegativity between the H atom and the atom to which it is bonded (F, O, or N), combined with the very small size of a H atom and the relatively small sizes of F, O, or N atoms, leads to highly concentrated partial charges with these atoms. Molecules with F-H, O-H, or N-H bonds are very strongly attracted to lone pairs on F, O, or N in nearby molecules, a particularly strong type of dipole-dipole attraction called hydrogen bonding. Examples of hydrogen bonds include HF⋯HF, H2O⋯HOH, and H3N⋯HNH2, in which the hydrogen bonds are denoted by dots. Figure 4 illustrates hydrogen bonding between water molecules.

Five water molecules are shown near one another, but not touching. A dotted line lies between many of the hydrogen atoms on one molecule and the oxygen atom on another molecule.
Figure 4. Water molecules participate in multiple hydrogen-bonding interactions with nearby water molecules.

Despite use of the word “bond,” keep in mind that hydrogen bonds are intermolecular attractive forces, not intramolecular (covalent bonds). Hydrogen bonds are much weaker than covalent bonds, only about 5 to 10% as strong, but are generally much stronger than other dipole-dipole attractions and dispersion forces.

Demonstration: Immiscibility of Molecules based on Chain Length

Set up. In this demonstration, water with green food coloring is added to four graduated cylinders, each containing a different liquid: ethanol, 1-propanol, 1-butanol, and 1-hexanol.

Prediction.For each liquid, predict whether it will be miscible or immiscible with the added water.

Explanation. In order for two molecules to be miscible, they must have similar intermolecular forces.  Water is a polar molecule and is only miscible with other molecules that are also polar. Although each of the four molecules that were mixed with water were alcohols, as the carbon-chain length increases, the majority of the molecule becomes non-polar. This is why ethanol and propanol are miscible with water, but alcohols with any longer carbon-chain are immiscible. This phenomenon is known as “like dissolves like”.

Key Concepts and Summary

After knowing both the molecular geometry and polarity, we can determine which intermolecular forces (IMF) each molecule interacts with. The stronger the IMF, the higher the melting and boiling points will be. Knowing which IMF exist within a chemical system will be extremely important for the rest of the semester! The weakest IMF are dispersion forces, which are present between all molecules with an electron, but are the only IMF between nonpolar molecules. The next strongest IMF are dipole-induced dipole forces. This IMF requires one molecule to be polar, while the other molecule is either polar or non-polar. The presence of the permanent dipole on one molecule induces a dipole (or a stronger/weaker dipole) in another molecule. The next IMF are dipole-dipole forces, which require both molecules to be polar. There is a special subset of dipole-dipole force called hydrogen bonding, which requires one molecule to have a hydrogen bonded to an F, O, or N and the other molecule to have a lone pair on F, O, or N. The strong but transient hydrogen bond forms between the hydrogen on on molecule and the lone pair on the other. IMF stronger than dipole-dipole require an ion to be present.

Glossary

dipole-dipole attraction
the attractive force between two polar molecules, with the attraction between the partially positive end of one molecule and the partially negative end of another

hydrogen bonding
a particularly strong type of dipole-dipole attraction, where the hydrogen of a molecule with F-H, O-H, or N-H is very strongly attracted to lone pairs on F, O or N in nearby molecules

induced dipole
a temporary dipole created by the presence of a dipole in a nearby atom or molecule

instantaneous dipole
a temporary dipole created by the constant motion of electrons in an atom or molecule

intermolecular forces (IMF)
the various forces of attraction that exist between atoms and/or molecules due to electrostatic properties

molecular geometry
geometry only based o the location of atoms, not lone electron pairs

nonpolar molecule
a molecule without a dipole moment

polar molecule
a molecule with a dipole moment

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