10.2 Types of Intermolecular Forces

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Figure 10.4

The attractions between molecules are not nearly as strong as the intramolecular attractions (bonds) that hold compounds together. Many physical properties reflect intermolecular forces, like boiling points, melting points, viscosity, surface tension, and capillary action.

1. Weakest to strongest forces: dispersion forces (or London dispersion forces)
2. Dipole–dipole forces
3. Hydrogen bonding (a special dipole–dipole force)
4. ion–dipole forces

Note: The first two types are also referred to collectively as van der Waals forces.

1. Dispersion Forces

The weakest type of intermolecular force is called London dispersion forces. London dispersion forces occur between all atoms and molecules, but they are so weak, they are only considered when there is no other intermolecular forces. For example, London dispersion forces exist between water molecules, but water molecules also have a permanent polar attraction so much stronger than the London dispersion forces that the London dispersion force is insignificant and not mentioned.

The cause of London dispersion forces is not obvious. Although we usually assume that the electrons of an atom are uniformly distributed around the nucleus, this is not true at every instance. As the electrons move around the nucleus, at a given instance, more electrons may be on one side of the nucleus than the other. This momentary nonsymmetrical electron distribution can produce a temporary dipolar arrangement of charge. This temporary dipole can induce a similar dipole in a neighboring atom and produce a weak, short-lived attraction.

The cases where London dispersion forces would be considered as the only intermolecular force of attraction would be for the noble gases and non-polar molecules such as helium, neon, argon, krypton, xenon, hydrogen, oxygen, methane, carbon dioxide, and so forth. Since non-polar molecules do not have a permanent dipole and no further bonding capacity, their only means of attracting each other is through London dispersion forces. Some of the substances whose intermolecular forces of attraction are London dispersion forces are held in the liquid state so weakly, they have the lowest melting points of all substances (see Table below for examples).

The temporary dipoles that cause London dispersion forces are affected by the molar mass of the particle. The greater the molar mass of the particle, the greater the force of attraction caused by London dispersion forces. The molar masses of H2, N2, and O2 are 2, 28, and 32 g/mol, respectively, and their boiling points increase in similar fashion; −253∘C for H2, −196∘C for N2, and −183∘C for O2. For molecules with a high molecular weight, the London dispersion forces become strong enough that the substance will be a liquid or solid even at room temperature. Carbon tetrachloride, molar mass 154 g/mol, and bromine, molar mass 160 g/mol, boil at +77∘C and +59∘C, respectively. Many long carbon chain, non-polar substances such as gasoline and oil remain liquids at common temperatures.

The figure below shows how a nonpolar particle
(in this case a helium atom) can be temporarily polarized to allow dispersion force to form.

Figure 10.5

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Factors affecting the degree of dispersion force in a molecule?

• number of electrons in an atom (more electrons, more dispersion force)
• size of atom or molecule/molecular weight
• shape of molecules with similar masses (more compact, less dispersion force)

neopentane, bp= 282.7K

n-pentane, bp=309.4 K

Another factor is polarizability. The tendency of an electron cloud to distort is
called its polarizability. If something is easier to polarize, it has a lower boiling point. Remember: This means less intermolecular force (smaller molecule: lower molecular weight, fewer electrons).

Figure 10.6

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Figure 10.7

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The fact that noble gas elements and completely non-polar molecules such as H₂ and N₂ can be condensed to liquids or solids tells us that there must be yet another source of attraction between particles that does not depend on the existence of permanent dipole moments in either particle. To understand the origin of this effect, it is necessary to realize that when we say a molecule is “nonpolar”, we really mean that the time-averaged dipole moment is zero. This is the same kind of averaging we do when we draw a picture of an orbital, which represents all the locations in space in which an electron can be found with a certain minimum probability. On a very short time scale, however, the electron must be increasingly localized; not even quantum mechanics allows it to be in more than one place at any given instant. As a consequence, there is no guarantee that the distribution of negative charge around the center of an atom will be perfectly symmetrical at every instant; every atom therefore has a weak, fluctuating dipole moment that is continually disappearing and reappearing in another direction.

Dispersion or London forces can be considered to be “induced dipole – induced dipole” interactions.

Although these extremely short-lived fluctuations quickly average out to zero, they can still induce new dipoles in a neighboring atom or molecule, which helps sustain the original dipole and gives rise to a weak attractive force known as the dispersion or London force (after F. London, who explained this effect in 1930).

Although dispersion forces are the weakest of all the intermolecular attractions, they are universally present. Their strength depends to a large measure on the number of electrons in a molecule. This can clearly be seen by looking at the noble gas elements, whose ability to condense to liquids and freeze to solids is entirely dependent on dispersion forces.

Table 10.3

It is important to note that dispersion forces are additive; if two elongated molecules find themselves side by side, dispersion force attractions will exist all along the regions where the two molecules are close. This can produce quite strong attractions between large polymeric molecules even in the absence of any stronger attractive forces.

2) Dipole-dipole Interactions

Non-bonding electrostatic interactions: Dipoles and dipole moments

According to Coulomb’s law, the electrostatic force between an ion and an uncharged particle having Q = 0 should be zero. Bear in mind, however, that this formula assumes that the two particles are point charges having zero radii. A real particle such as an atom or a molecule occupies a certain volume of space. Even if the electric charges of the protons and electrons cancel out (as they will in any neutral atom or molecule), it is possible that the spatial distribution of the electron cloud representing the most loosely-bound [valence] electrons might be asymmetrical, giving rise to an electric dipole moment. There are two kinds of dipole moments:

Permanent electric dipole moments can arise when bonding occurs between elements of differing electronegativities.

Induced (temporary) dipole moments are created when an external electric field distorts the electron cloud of a neutral molecule.

In physics, an electric dipole refers to a separation of electric charge. An idealized electric dipole consists of two point charges of magnitude +q and –q separated by a distance r. Even though the overall system is electrically neutral, the charge separation gives rise to an electrostatic effect whose strength is expressed by the electric dipole moment given by

μ = q r

Dipole moments possess both magnitude and direction, and are thus vectorial quantities; they are conventionally represented by arrows whose heads are at the negative end.

Permanent dipole moments

See our CHEM 1211 eBook of the lesson on molecular structure for more on dipole moments.

In many molecules that are otherwise electrically neutral, the differing electronegativities of the various atoms give rise to an uneven distribution of negative charge. Such molecules are said to be polar, and to possess a permanent dipole moment (the word “permanent” is commonly dropped in ordinary discourse.)

Figure 10.8

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:The most well-known molecule having a dipole moment is ordinary water. The charge imbalance arises because oxygen, with its nuclear charge of 8, pulls the electron cloud that comprises each O–H bond toward itself. These two “bond momens” add vectorially to produce the permanent dipole moment denoted by the red arrow. Note the use of the δ (Greek delta) symbol to denote the positive and negative ends of the dipoles.

When an electric dipole is subjected to an external electric field, it will tend to orient itself so as to minimize the potential energy; that is, its negative end will tend to point toward the higher (more positive) electric potential. In liquids, thermal motions will act to disrupt this ordering, so the overall effect depends on the temperature. In condensed phases the local fields due to nearby ions or dipoles in a substance play an important role in determining the physical properties of the substance, and it is in this context that dipolar interactions are of interest to us here. We will discuss each kind of interaction in order of decreasing strength.

Induced dipoles

Even if a molecule is electrically neutral and possesses no permanent dipole moment, it can still be affected by an external electric field. Because all atoms and molecules are composed of charged particles (nuclei and electrons), the electric field of a nearby ion will cause the centers of positive and negative charges to shift in opposite directions. This effect, which is called polarization, results in the creation of a temporary, or induced dipole moment. The induced dipole then interacts with the species that produced it, resulting in a net attraction between the two particles.

The larger an atom or ion, the more loosely held are its outer electrons, and the more readily will the electron cloud by distorted by an external field. A quantity known as the polarizability expresses the magnitude of the temporary dipole that can be induced in it by a nearby charge.

Ion-Dipole interactions

Figure 10.9

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A dipole that is close to a positive or negative ion will orient itself so that the end whose partial charge is opposite to the ion charge will point toward the ion. This kind of interaction is very important in aqueous solutions of ionic substances; H₂O is a highly polar molecule, so that in a solution of sodium chloride, for example, the Na⁺ ions will be enveloped by a shell of water molecules with their oxygen-ends pointing toward these ions, while H₂O molecules surrounding the Cl^– ions will have their hydrogen ends directed inward. As a consequence of ion-dipole interactions, all ionic species in aqueous solution are hydrated; this is what is denoted by the suffix in formulas such as K⁺(aq), etc.

The strength of ion-dipole attraction depends on the magnitude of the dipole moment and on the charge density of the ion. This latter quantity is just the charge of the ion divided by its volume. Owing to their smaller sizes, positive ions tend to have larger charge densities than negative ions, and they should be more strongly hydrated in aqueous solution. The hydrogen ion, being nothing more than a bare proton of extremely small volume, has the highest charge density of any ion; it is for this reason that it exists entirely in its hydrated form H₃O⁺ in water.

Figure 10.10

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When covalent bonds form between identical atoms, such as in H2, N2, O2, and so on, the electrons shared in the bonds are shared equally. The two atoms have the same electronegativity and therefore the same pull on the shared electrons (as illustrated in the figure below).

Figure 10.11

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The center of negative charge for the entire molecule will be in the exact center of the molecule. This will coincide with the center of positive charge for the molecule. When the center of negative charge and the center of positive charge coincide, there is no charge separation and no dipole.

In the case of a symmetrical molecule with polar bonds, like the one shown below, the symmetry of the electron displacements will also keep the center of negative charge in the center of the molecule, which coincides with the center of positive charge. As a result, no dipole will occur.

Figure 10.12

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If the two atoms sharing the bonding pair of electrons are not of the same element, the atom with the greater electronegativity will pull the shared electrons closer to it. Because of the resulting uneven distribution of electrons, the center of negative charge will not coincide with the center of positive charge and a dipole is created on the molecule. When the centers of positive and negative charge do not coincide, a charge separation exists and a dipole is present. For example, in the CO2 molecule above, both carbon-oxygen bonds are polar and the bonding electrons are shifted toward the oxygen.

The end of the molecule with the more electronegative atom will have a partial negative charge, and the end of the molecule with the more electropositive atom will have a slight positive charge. The symbols δ+ and δ−, as illustrated in the figure below, are used because these are not full positive and negative charges.

Figure 10.13

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This polarity is much less than the charge separation in an ionic bond. These charges are only fractions of the full “+1” and “−1” charges. How much polarity a bond will experience depends on the difference in the electronegativities of the atoms.

For molecules that have a permanent dipole, the attraction between oppositely charged ends of adjacent molecules are the dominant intermolecular force of attraction. The figure below represents a polar solid; a polar liquid would look similar except the molecules would be less organized. On average, these polar attractions are stronger than London dispersion forces, so polar molecules in general have higher boiling points than London dispersion liquids. There is significant overlap, however, between the boiling points of the stronger London dispersion molecules and the weaker polar molecules.

Figure 10.14

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For molecules of approximately equal mass and size, the more polar the molecule, the higher its boiling point.

Figure 10.15-10.18

neo-pentane: MW: 44 amu, bp= 231 K

dimethyl ether: MW: 46 amu

Bp= 248 K

Acetaldehyde: MW: 44 amu

Bp=  294 K

Acetonitrile: MW: 41 amu

Bp= 355 K                                                                        

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Which Have a Greater Effect: Dipole–Dipole Interactions or Dispersion Forces?

We may say that if two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

Combinations of interactions: van der Waals forces

Although nonpolar molecules are by no means uncommon, many kinds of molecules possess permanent dipole moments, so liquids and solids composed of these species will be held together by a combination of dipole-dipole, dipole-induced dipole, and dispersion forces. These weaker forces (that is, those other than coulombic attractions) are known collectively as van der Waals forces.

The following table shows some estimates of the contributions of the various types of van der Waals forces that act between several different types of molecules. Note particularly how important dispersion forces are in all of these examples, and how this, in turn, depends on the polarizability.

Table 10.4

• Hydrogen Bond:

Figure 10.19

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1. The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds.
2. A hydrogen bond is an attraction between a hydrogen atom attached to a highly electronegative atom and a nearby small electronegative atom in another molecule or chemical group.
3. Attractive forces on the molecular level are divided into two categories – the forces inside a molecule holding atoms together, and the forces between molecules holding molecules together. In the image below, the forces holding the oxygen and hydrogen atoms together are the intramolecular forces. Intramolecular forces are the forces inside the molecule and consist of ionic and covalent bonds. The forces that hold the H₂O molecules together as either liquid water or ice are the intermolecular forces. Intermolecular forces are the forces between molecules and are responsible for holding molecules in the solid or liquid state when no covalent or ionic bonds are possible. You can relate these two types of forces to bus systems where intra-city buses move people around inside one city and inter-city buses move people from one city to another
4. Source: https://www.chem1.com/acad/webtext/virtualtextbook.html
6.  

Fig. 10.20

• When a substance in the solid phase is heated sufficiently, the molecular motions increase to the point that they overcome the forces holding the solid together, so the solid melts to liquid form. If the liquid is then heated sufficiently, the molecular motion increases to the point that it completely overcomes the attractive forces, so the liquid will change to the gaseous phase. The reverse of this process occurs when the substance is cooled. The conversion of solid to liquid, liquid to gas, gas to liquid, and liquid to solid are called phase changes.
• The intermolecular forces must be overcome during a phase change. Therefore, the stronger the intermolecular forces of attraction, the greater the molecular motion (temperature) required to overcome them. The solids and liquids with the strongest intermolecular forces of attraction will have the highest melting and boiling points, and vice versa.

Fig 10.21

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Dipoles of Hydrogen Bonds:

Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.

• These atoms interact with a nearly bare nucleus (which contains one proton).

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Fig. 10.22

Water: the molecule

In water, each hydrogen nucleus is covalently bound to the central oxygen atom by a pair of electrons that are shared between them. In H₂O, only two of the six outer-shell electrons of oxygen are used for this purpose, leaving four electrons which are organized into two non-bonding pairs. The four electron pairs surrounding the oxygen tend to arrange themselves as far from each other as possible to minimize repulsions between these clouds of negative charge. This would ordinarily result in a tetrahedral geometry in which the angle between electron pairs (and therefore the H-O-H bond angle) is 109.5°. However, because the two non-bonding pairs remain closer to the oxygen atom, these exert a stronger repulsion against the two covalent bonding pairs, effectively pushing the two hydrogen atoms closer together. The result is a distorted tetrahedral arrangement in which the H—O—H angle is 104.5°.

Water’s large dipole moment leads to hydrogen bonding

Fig. 10.23

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The H₂O molecule is electrically neutral, but the positive and negative charges are not distributed uniformly. This is illustrated by the gradation in color in the schematic diagram here. The electronic (negative) charge is concentrated at the oxygen end of the molecule, owing partly to the nonbonding electrons (solid blue circles), and to oxygen’s high nuclear charge which exerts stronger attractions on the electrons. This charge displacement constitutes an electric dipole, represented by the arrow at the bottom; you can think of this dipole as the electrical “image” of a water molecule.

Opposite charges attract, so it is not surprising that the negative end of one water molecule will tend to orient itself so as to be close to the positive end of another molecule that happens to be nearby. The strength of this dipole-dipole attraction is less than that of a normal chemical bond, and so it is completely overwhelmed by ordinary thermal motions in the gas phase.

Fig. 10.24

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Hydrogen bonding in water

But when the H₂O molecules are crowded together in the liquid, these attractive forces exert a very noticeable effect, which we call (somewhat misleadingly) hydrogen bonding. And at temperatures low enough to turn off the disruptive effects of thermal motions, water freezes into ice in which the hydrogen bonds form a rigid and stable network.

Notice that the hydrogen bond (shown by the dashed green line) is somewhat longer than the covalent O—H bond. It is also much weaker, about 23 kJ mol^–1 compared to the O–H covalent bond strength of 492 kJ mol^–1.

Fig. 10.25

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One factor that makes water unique even among other hydrogen-bonded liquids is its very small mass in relation to the large number of hydrogen bonds it can form. Owing to disruptions of these weak attractions by thermal motions, the lifetime of any single hydrogen bond is very short — on the order of a picosecond. At any instant, the average H₂O molecule is bound to somewhat fewer than four neighbors — estimates vary from 2.4 to 3.6.

Why water’s boiling point is so high?

The most apparent pecularity of water is its very high boiling point for such a light molecule. Liquid methane CH₄ (molecular weight 16) boils at –161°C,

As you can see from this diagram, extrapolation of the boiling points of the various Group 16 hydrogen compounds to H₂O suggests that this substance should be a gas under normal conditions.

 Fig. 10.26

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Why ice floats on water?

The most energetically favorable configuration of H₂O molecules is one in which each molecule is hydrogen-bonded to four neighboring molecules. Owing to the thermal motions described above, this ideal is never achieved in the liquid, but when water freezes to ice, the molecules settle into exactly this kind of an arrangement in the ice crystal. This arrangement requires that the molecules be somewhat farther apart then would otherwise be the case; as a consequence, ice, in which hydrogen bonding is at its maximum, has a more open structure, and thus a lower density than water.

Fig. 10.27 [image source]

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Here are three-dimensional views of a typical local structure of water (left) and ice (right.) Notice the greater openness of the ice structure which is necessary to ensure the strongest degree of hydrogen bonding in a uniform, extended crystal lattice. The more crowded and jumbled arrangement in liquid water can be sustained only by the greater amount of thermal energy available above the freezing point. 

Fig. 10.28

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When ice melts, the more vigorous thermal motion disrupts much of the hydrogen-bonded structure, allowing the molecules to pack more closely. Water is thus one of the very few substances whose solid form has a lower density than the liquid at the freezing point. Localized clusters of hydrogen bonds still remain, however; these are continually breaking and reforming as the thermal motions jiggle and shove the individual molecules. As the temperature of the water is raised above freezing, the extent and lifetimes of these clusters diminish, so the density of the water increases.

At higher temperatures, another effect, common to all substances, begins to dominate: as the temperature increases, so does the amplitude of thermal motions. This more vigorous jostling causes the average distance between the molecules to increase, reducing the density of the liquid; this is ordinary thermal expansion.

Because the two competing effects (hydrogen bonding at low temperatures and thermal expansion at higher temperatures) both lead to a decrease in density, it follows that there must be some temperature at which the density of water passes through a maximum. This temperature is 4° C; this is the temperature of the water you will find at the bottom of an ice-covered lake in which this most dense of all water has displaced the colder water and pushed it nearer to the surface.

Hydrogen bonding in biopolymers

Hydrogen bonding plays an essential role in natural polymers of biological origin in two ways: Hydrogen bonding between adjacent polymer chains (intermolecular bonding); Hydrogen bonding between different parts of the same chain (intramolecular bonding;

Hydrogen bonding of water molecules to –OH groups on the polymer chain (“bound water”) that helps maintain the shape of the polymer.

The examples that follow are representative of several types of biopolymers.

Cellulose

Fig. 10.29

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Cellulose is a linear polymer of glucose (see above), containing 300 to over 10,000 units, depending on the source. As the principal structural component of plants (along with lignin in trees), cellulose is the most abundant organic substance on the earth. The role of hydrogen bonding is to cross-link individual molecules to build up sheets as shown here. These sheets than stack up in a staggered array held together by van der Waals forces. Further hydrogen-bonding of adjacent stacks bundles them together into a stronger and more rigid structure.

Proteins

These polymers made from amino acids R—CH(NH₂)COOH depend on intramolecular hydrogen bonding to maintain their shape (secondary and tertiary structure) which is essential for their important function as biological catalysts (enzymes). Hydrogen-bonded water molecules embedded in the protein are also important for their structural integrity.

The principal hydrogen bonding in proteins is between the -N—H groups of the “amino” parts with the -C=O groups of the “acid” parts. These interactions give rise to the two major types of the secondary structure which refers to the arrangement of the amino acid polymer chain:

Fig. 10.30

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Although carbon is not usually considered particularly electronegative, C—H—-X hydrogen bonds are also now known to be significant in proteins.

DNA (Deoxyribonucleic acid)

Who you are is totally dependent on hydrogen bonds! DNA, as you probably know, is the most famous of the biopolymers owing to its central role in defining the structure and function of all living organisms.

Fig. 10.31

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[image]

Each strand of DNA is built from a squence of four different nuclotide monomers consisting of a deoxyribose sugar, phosphate groups, and a nitrogenous base conventionally identified by the letters A,T, C and G. DNA itself consists of two of these polynuclotide chains that are coiled around a common axis in a configuration something like the protein alpha helix depicted above. The sugar-and-phosphate backbones are on the outiside so that the nucleotide bases are on the inside and facing each other. The two strands are held together by hydrogen bonds that link a nitrogen atom of a nucleotide in one chain with a nitrogen or oxygen on the nucleotide that is across from it on the other chain.

Efficient hydrogen bonding within this configuration can only occur between the pairs A-T and C-G, so these two complementary pairs constitute the “alphabet” that encodes the genetic information that gets transcribed whenever new protein molecules are built.

Water molecules, hydrogen-bonded to the outer parts of the DNA helix, help stabilize its structure:

Adapted from the Chem1 Virtual Textbook by Stephen Lower. http://www.chem1.com/acad/webtext/virtualtextbook.html