9.2: Introduction to Intermolecular Forces

London Dispersion Forces

Dispersion Forces (also called London Forces) are the weakest intermolecular forces. It results from the instantaneous dipole and induced dipole of the molecules. For nonpolar molecules, the constant shifting and distortion of electron density leads to a short-lived dipole at a given moment, which is called an instantaneous dipole. Such temporary dipoles will induce the electrons in a neighboring molecule to get distorted as well, and to develop a corresponding transient dipole of its own, which is the induced dipole. These rapidly fluctuating, temporary dipoles thus result in a relatively weak electrostatic attraction between the species—a so-called dispersion force like that illustrated in Figure \(\PageIndex{3}\).

Ultimately, all nonpolar molecules are attracted to each other; however, considering that the dipole is in place for a short time, the interaction is weak. Since it applies to all types of molecules (it is the only intermolecular force for nonpolar molecules), dispersion forces are also the most fundamental intermolecular force.

Non-polar Molecules with non-polar bonds experience no other type of intermolecular interaction than dispersion forces. For Molecules with polar bonds, where the covalent bonds are oriented in space in such a way that the effect of the polar bonds in the molecules cancel each other out, the individual bonds are polar, but due to molecular symmetry, the overall molecule is nonpolar. Such molecules only experience London Forces. Carbon dioxide (CO₂) and carbon tetrachloride (CCl₄) are examples of such molecules (Figure \(\PageIndex{4}\)).

The Magnitude of Dispersion Forces

depends on two factors:

• Larger and heavier atoms and molecules. The relative polarizability of electrons. The simple understanding of polarizability is how easily the electrons get distorted. For larger atoms, there are more electrons in a larger space, therefore the electrons are more loosely held and more easily polarized, so the dispersion force is stronger. Generally, the larger the molar mass of the molecule, the stronger the dispersion force.
• The relative surface area of the molecule. Molecules with longer, flatter, or cylindrical shapes have a greater surface area compared to the bulky, branched molecules, and therefore have a stronger dispersion force. Taking the two constitutional isomers of C₄H₁₀, butane, and isobutane, as an example, the dispersion force of butane is stronger than that of isobutane. Figure \(\PageIndex{5}\).

Substances that experience only dispersion forces are typically soft in the solid phase and have relatively low melting points. Examples include waxes, which are long hydrocarbon chains that are solids at room temperature because the molecules have a high number of electrons.

Exercise \(\PageIndex{1}\)

Order the following hydrocarbons from weakest to strongest total intermolecular force: C₂H₆, C₃H₈, and C₄H₁₀.

Answer

All of these compounds are nonpolar and exhibit only London dispersion forces; the larger the molecule, the stronger the dispersion forces. The ordering from weakest to strongest total intermolecular force is therefore

C₂H₆ < C₃H₈ < C₄H₁₀.

Dipole-Dipole Forces

A covalent bond that has an unequal sharing of electrons is called a polar covalent bond. A covalent bond that has an equal sharing of electrons is called a nonpolar covalent bond. The charge separation in a polar covalent bond is not as extreme as is found in ionic compounds, but there is a related result: oppositely charged ends of different molecules will attract each other. This type of intermolecular interaction is called a dipole-dipole interaction. Most molecules with polar covalent bonds experience dipole-dipole interactions. HF is an example of a polar molecule, also known as a dipole.

Polar molecules have a partial positive charge on one side and a partial negative charge on the other side of the molecule—a separation of charge called a dipole. Consider a polar molecule such as hydrogen chloride, HCl. In the HCl molecule, the more electronegative Cl atom bears the partial negative charge, whereas the less electronegative H atom bears the partial positive charge. 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 \(\PageIndex{7}\).

Hydrogen Bonding Intermolecular Forces

Molecules with hydrogen atoms bonded to electronegative atoms such as O, N, and F tend to exhibit unusually strong intermolecular interactions due to a particularly strong type of dipole-dipole attraction called hydrogen bonding. The large difference in electronegativity between the H atom (2.1) and the atom to which it is bonded (4.0 for an F atom, 3.5 for an O atom, or 3.0 for a N atom), 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. Because the hydrogen atom does not have any electrons other than the ones in the covalent bond, its positively charged nucleus is almost completely exposed, allowing strong attractions to other nearby lone pairs of electrons.

Despite use of the word “bond,” keep in mind that hydrogen bonds are intermolecular attractive forces, not intramolecular attractive forces (covalent bonds). Hydrogen bonds are much weaker than covalent bonds, but are generally much stronger than other dipole-dipole attractions.

Examples of hydrogen bonds include HF⋯HF, H₂O⋯HOH, and H₃N⋯HNH₂, in which the hydrogen bonds are denoted by dots. Figure \(\PageIndex{9}\) illustrates hydrogen bonding between water molecules. The physical properties of water, which has two O–H bonds, are strongly affected by the presence of hydrogen bonding between water molecules. Most molecular compounds that have a mass similar to water are gases at room temperature. However, due to the strong hydrogen bonds, water molecules are able to remain condensed in the liquid state.

A hydrogen bond is an intermolecular attractive force in which a hydrogen atom, that is covalently bonded to a small, highly electronegative atom, is attracted to a lone pair of electrons on an atom in a neighboring molecule. Figure \(\PageIndex{10}\) shows how methanol (CH₃OH) molecules experience hydrogen bonding. Methanol contains both a hydrogen atom attached to O; methanol can thus form hydrogen bonds by acting as either a hydrogen bond donor or a hydrogen bond acceptor (lone pair). The hydrogen-bonded structure of methanol is as follows:

Ion-Dipole

Ion-dipole force is not a common intermolecular force, however it is a type of important force that is responsible for the interaction between ions and other polar substance. A simple example is the dissolving of an ionic solid, or salt, in water. When table salt (NaCl) is dissolved in water, the interactions between the ions and water molecules are strong enough to overcome the ionic bond that holds the ions in the crystal lattice. As a result, the cations and anions are separated apart completely, and each ion is surrounded by a cluster of water molecules. This is called a solvation process. The solvation occurs through the strong ion-dipole force. Lots salts, or ionic compounds, are soluble in water because of such interactions.

Ionic Bond

The strongest force between any two particles is the ionic bond, in which two ions of opposing charge are attracted to each other. Thus, ionic interactions between particles are another type of intermolecular interaction. Substances that contain ionic interactions are relatively strongly held together, so these substances typically have high melting and boiling points. Sodium chloride (Figure \(\PageIndex{12}\)) is an example of a substance whose particles experience ionic interactions (Table \(\PageIndex{1}\)). These attractive forces are sometimes referred to as ion-ion interactions.

Folding Proteins

These interactions are common in proteins and contribute to the protein's folding. Ion-ion interactions are called salt bridges.

Figure \(\PageIndex{13}\): A protein strand in blue folds into a compact structure via ionic interactions called salt bridges.

Hydrogen Bonding and DNA

Deoxyribonucleic acid (DNA) is found in every living organism and contains the genetic information that determines the organism’s characteristics, provides the blueprint for making the proteins necessary for life, and serves as a template to pass this information on to the organism’s offspring. A DNA molecule consists of two anti-parallel chains of repeating nucleotides, which form its well-known double helical structure, as shown in Figure \(\PageIndex{14}\).

Each nucleotide contains a (deoxyribose) sugar bound to a phosphate group on one side, and one of four nitrogenous bases on the other. Two of the bases, cytosine (C) and thymine (T), are single-ringed structures known as pyrimidines. The other two, adenine (A) and guanine (G), are double-ringed structures called purines. These bases form complementary base pairs consisting of one purine and one pyrimidine, with adenine pairing with thymine, and cytosine with guanine. Each base pair is held together by hydrogen bonding. A and T share two hydrogen bonds, C and G share three, and both pairings have a similar shape and structure Figure \(\PageIndex{15}\)

The cumulative effect of millions of hydrogen bonds effectively holds the two strands of DNA together. Importantly, the two strands of DNA can relatively easily “unzip” down the middle since hydrogen bonds are relatively weak compared to the covalent bonds that hold the atoms of the individual DNA molecules together. This allows both strands to serve as templates for replication.

Applications: Geckos and Intermolecular Forces

Geckos have an amazing ability to adhere to most surfaces. They can quickly run up smooth walls and across ceilings that have no toe-holds, and they do this without having suction cups or a sticky substance on their toes. And while a gecko can lift its feet easily as it walks along a surface, if you attempt to pick it up, it sticks to the surface. How are geckos (as well as spiders and some other insects) able to do this? Although this phenomenon has been investigated for hundreds of years, scientists only recently uncovered the details of the process that allows geckos’ feet to behave this way.

Geckos’ toes are covered with hundreds of thousands of tiny hairs known as setae, with each seta, in turn, branching into hundreds of tiny, flat, triangular tips called spatulae. The huge numbers of spatulae on its setae provide a gecko, shown in Figure \(\PageIndex{16}\), with a large total surface area for sticking to a surface. In 2000, Kellar Autumn, who leads a multi-institutional gecko research team, found that geckos adhered equally well to both polar silicon dioxide and nonpolar gallium arsenide. This proved that geckos stick to surfaces because of dispersion forces—weak intermolecular attractions arising from temporary, synchronized charge distributions between adjacent molecules. Although dispersion forces are very weak, the total attraction over millions of spatulae is large enough to support many times the gecko’s weight.

In 2014, two scientists developed a model to explain how geckos can rapidly transition from “sticky” to “non-sticky.” Alex Greaney and Congcong Hu at Oregon State University described how geckos can achieve this by changing the angle between their spatulae and the surface. Geckos’ feet, which are normally nonsticky, become sticky when a small shear force is applied. By curling and uncurling their toes, geckos can alternate between sticking and unsticking from a surface, and thus easily move across it. Further investigations may eventually lead to the development of better adhesives and other applications.