M10Q1: An Introduction to Intermolecular Forces

Introduction

In Modules 8 and 9, the attractive forces between atoms within an individual molecule (covalent bonds) has been examined. This section explores intermolecular forces, a collective term for attractive forces that occur between molecules, ions and non bonded atoms in bulk samples. The different types of intermolecular forces that exist are introduced, as well as how to determine what types of intermolecular force are important in a given substance or system.

Learning Objectives for Intermolecular Forces

| Key Concepts and Summary | Glossary | End of Section Exercises |

As was the case for gaseous substances, the kinetic molecular theory may be used to explain the behavior of solids and liquids. In the following description, the term particle will be used as a collective term to refer to an atom, molecule, or ion. Note that we will use the popular phrase “intermolecular attraction” to refer to attractive forces between the particles of a substance, regardless of whether these particles are molecules, atoms, or ions.

Consider these two aspects of the molecular-level environments in solid, liquid, and gaseous matter:

  • Particles in a solid are tightly packed together and often arranged in a regular pattern; in a liquid, they are close together with no regular arrangement; in a gas, they are far apart with no regular arrangement.
  • Particles in a solid vibrate about fixed positions and do not generally move in relation to one another; in a liquid, they move past each other but remain in essentially constant contact; in a gas, they move independently of one another except when they collide.

The differences in the properties of a solid, liquid, or gas reflect the strengths of the attractive forces between the atoms, molecules, or ions that make up each phase. The phase in which a substance exists depends on the relative extents of its intermolecular forces (IMFs) and the kinetic energies (KE) of its molecules. IMFs are the various forces of attraction that may exist between the atoms and molecules of a substance due to electrostatic phenomena, as will be detailed in this module. In this section, we will use the phase of a substance as an approximation of its IMFs, with solids having the strongest IMFs and gases having the weakest. These forces serve to hold particles close together, whereas the particles’ KE provides the energy required to overcome the attractive forces and thus increase the distance between particles. Figure 1 illustrates how changes in physical state may be induced by changing the temperature, hence the average KE, of a given substance.

Three sealed flasks are labeled, “Crystalline solid,” “Liquid,” and “Gas,” from left to right. The first flask holds a cube composed of small spheres sitting on the bottom while the second flask shows a lot of small spheres in the bottom that are spaced a small distance apart from one another and have lines around them to indicate motion. The third flask shows a few spheres spread far from one another with larger lines to indicate motion. There is a right-facing arrow that spans the top of all three flasks. The arrow is labeled, “Increasing K E ( temperature ).”
Figure 1. Transitions between solid, liquid, and gaseous states of a substance occur when conditions of temperature or pressure favor the associated changes in intermolecular forces. (Note: The space between particles in the gas phase is much greater than shown.)

As an illustration of a process depicted in this figure, consider a sample of water vapor. When gaseous water is cooled sufficiently, by coming into contact with a cold surface for example, the attractions between H2O molecules will be capable of holding the molecules together when they come into contact with each other; the gas condenses, forming liquid H2O. For example, liquid water forms on the outside of a cold glass as the water vapor in the air is cooled by the cold glass, as seen in Figure 2.

Image a shows a brown colored beverage in a glass with condensation on the outside. Image b shows a body of water with fog hovering above the surface of the water.
Figure 2. Condensation forms when water vapor in the air is cooled enough to form liquid water, such as (a) on the outside of a cold beverage glass or (b) in the form of fog. (credit a: modification of work by Jenny Downing; credit b: modification of work by Cory Zanker)

Finally, if the temperature of a liquid becomes sufficiently low, or the pressure on the liquid becomes sufficiently high, the molecules of the liquid no longer have enough KE to overcome the IMF between them, and a solid forms.

Intramolecular Forces vs. Intermolecular Forces

Intermolecular attractive forces are present in all substances as evidenced by the fact that everything will eventually solidify under appropriate conditions. They may be very weak, as is the case with the noble gases, but eventually even these atomic substances will liquify and then freeze if the temperature drops low enough. This is a result of the intermolecular forces as opposed to intramolecular forces. Intramolecular forces are those within a molecule that keep the atoms bonded together as an integral molecular unit. Intermolecular forces are the attractions between molecules or non bonded atoms, which determine many of the physical properties of a substance. Figure 3 illustrates the distinction between these two.

An image is shown in which two molecules composed of a green sphere labeled “C l” connected on the right to a white sphere labeled “H” are near one another with a dotted line labeled “Intermolecular force ( weak )” drawn between them. A line connects the two spheres in each molecule and the line is labeled “Intramolecular force ( strong ).”
Figure 3. Intramolecular forces keep a molecule intact. Intermolecular forces hold multiple molecules together and determine many of a substance’s properties.

The strengths of IMFs vary widely, though usually the IMFs between small molecules are weak compared to the intramolecular forces that bond atoms together within a molecule. For example, to overcome the IMFs in one mole of liquid HCl(l) and convert it into gaseous HCl(g) (a physical change) requires only about 17 kilojoules. However, to break the covalent bonds between the hydrogen and chlorine atoms (a chemical change) in one mole of HCl requires about 25 times more energy—430 kilojoules. Attractive forces between neutral atoms and molecules are known collectively as van der Waals forces, although they are usually referred to more informally as intermolecular attraction. We will consider the various types of IMFs in the next sections of this module.

Dispersion Forces

One of the van der Waals forces is present in all condensed phases, regardless of the nature of the atoms or molecules composing the substance. This attractive force is called the London dispersion force in honor of German-born American physicist Fritz London who, in 1928, first explained it. This force is often referred to as simply the dispersion force or London force. Because the electrons of an atom or molecule are in constant motion (or, alternatively, the electron’s location is subject to quantum-mechanical variability), at any moment in time, an atom or molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically, as shown in Figure 4.

This gif shows an oval that represents a molecule. As the dipole fluctuates, one side of the oval becomes larger to represent more electron density and the other side becomes smaller to represent a lack of electron density. Then the electron density changes and the opposite side of the oval becomes larger.
Figure 4. Fluctuating dipole in a non-polar molecule. The pink region is a region of high electron density.

The presence of this transient dipole can, in turn, distort the electrons of a neighboring atom or molecule in the sample, producing an induced dipole in the other molecule. Another name for this type of force is instantaneous dipole-induced dipole, though this term is less commonly used in our course. These two 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 5.

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 5. 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, however, and become significant only when the molecules are very close. Larger and heavier atoms and molecules exhibit stronger dispersion forces than do smaller and lighter atoms and molecules. F2 (38 g/mole) and Cl2 (71 g/mole) are gases at room temperature (reflecting weaker attractive forces); Br2 (160 g/mole) is a liquid, and I2 (254 g/mole) is a solid (reflecting stronger attractive forces). We will discuss this more in the next section on boiling points and melting points..

In a larger atom, the valence electrons are, on average, farther from the nuclei than in a smaller atom. Thus, the electrons are less tightly held and can more easily form the temporary dipoles that produce this 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 a bulk sample will have relatively strong dispersion forces; one with a charge cloud that is difficult to distort is not very polarizable and a bulk sample will have relatively weak dispersion forces.

The shapes of molecules also affect the magnitudes of the dispersion forces between them. For example, consider the isomers n-pentane, isopentane, and neopentane (shown in Figure 6). Even though these compounds are composed of molecules with the same chemical formula, C5H12, the dispersion forces in the liquid phase are different, being greatest for n-pentane and least for 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, 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” 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” 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”
Figure 6. The strength of the dispersion forces increases with the contact area between molecules.

Chemistry in Real Life: 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 7, 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.

Three figures are shown. The first is a photo of the bottom of a gecko’s foot. The second is bigger version which shows the setae. The third is a bigger version of the setae and shows the spatulae.
Figure 7. Geckos’ toes contain large numbers of tiny hairs (setae), which branch into many triangular tips (spatulae). Geckos adhere to surfaces because of van der Waals attractions between the surface and a gecko’s millions of spatulae. By changing how the spatulae contact the surface, geckos can turn their stickiness “on” and “off.” (credit photo: modification of work by “JC*+A!”/Flickr)

Dipole-Dipole Attractions

Recall from the module on Molecule Polarity that 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 8.

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 8. 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 to nonpolar F2. Both HCl and F2 are diatomic molecules with approximately the same molar mass and therefore bulk samples of these substances would have comparable London forces. However, HCl(l) has a much higher boiling point (188 K) compared to F2(l) (85 K), an indication that HCl(l) has stronger IMFs as a result of dipole-dipole attractions in addition to the London forces.

Hydrogen Bonding

Nitrosyl fluoride (ONF, molecular mass 49 amu) is a gas at room temperature. Water (H2O, molecular mass 18 amu) is a liquid, even though it has a lower molecular mass. We clearly cannot attribute this difference 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 to differences in the dipole moments of the molecules. Both molecules are polar and exhibit comparable dipole moments. The difference 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 (three of the most electronegative elements). The very large difference in electronegativity between the H atom (2.2) and the atom to which it is bonded (4.0 for an F atom, 3.5 for an O atom, or 3.1 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. Molecules with F-H, O-H, or N-H groups are very strongly attracted to similar groups in nearby molecules, and a particularly strong type of dipole-dipole attraction called hydrogen bonding is formed. This hydrogen bond forms from the interaction between the lone pair of electrons on an N, O, or F atom (the donor) and the antibonding orbital in a highly polarized sigma bond (the acceptor). Examples of hydrogen bonds include HF⋯HF, H2O⋯HOH, and H3N⋯HNH2, in which the hydrogen bonds are denoted by dots.

Figure 9 illustrates hydrogen bonding between water molecules. Notice that the hydrogen bond forms between a hydrogen that is attached to an oxygen atom on one molecule and an oxygen atom on an adjacent molecule. This example highlights the two requirements for forming a hydrogen bond: (1) One molecule must have an H atom bonded to an N, O, or F atom; and (2) another molecule must have an N, O, or F atom with a lone pair of electrons.

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 9. Water molecules participate in multiple hydrogen-bonding interactions with nearby water molecules. Click on the image above to be taken to WebMO to be able to rotate the water molecule.

Despite use of the word “bond,” keep in mind that 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. For example, hydrogen bonds have a pronounced effect on the properties of condensed phases (liquids and solids).

So far we have emphasized that IMFs are between separate molecules, but this is not always the case. In very large molecules such as proteins or DNA, these IMFs can occur within the same molecule and help stabilize their three-dimensional structure, which is essential for their mode of operation. The hydrogen bonding in DNA is described below.

Chemistry in Real Life: 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 10.

Two images are shown. The first lies on the left side of the page and shows a helical structure like a twisted ladder where the rungs of the ladder, labeled “Base pair” are red, yellow, green and blue paired bars. The red and yellow bars, which are always paired together, are labeled in the legend, which is titled “Nitrogenous bases” as “adenine” and “thymine,” respectively. The blue and green bars, which are always paired together, are labeled in the legend as “guanine” and “cytosine,” respectively. At the top of the helical structure, the left-hand side rail, or “Sugar, dash, phosphate backbone,” is labeled as “3, prime” while the right is labeled as “5, prime.” These labels are reversed at the bottom of the helix. To the right of the page is a large Lewis structure. The top left corner of this structure, labeled “5, prime,” shows a phosphorus atom single bonded to three oxygen atoms, one of which has a superscripted negative charge, and double bonded to a fourth oxygen atom. One of the single bonded oxygen atoms is single bonded to the left corner of a five-membered ring with an oxygen atom at its top point and which is single bonded to an oxygen atom on the bottom left. This oxygen atom is single bonded to a phosphorus atom that is single bonded to two other hydrogen atoms and double bonded to a fourth oxygen atom. The lower left of these oxygen atoms is single bonded to another oxygen atom that is single bonded to a five-membered ring with an oxygen in the upper bonding site. The bottom left of this ring has a hydroxyl group attached to it while the upper right carbon is single bonded to a nitrogen atom that is part of a five-membered ring bonded to a six-membered ring. Both of these rings have points of unsaturation and nitrogen atoms bonded into their structures. On the right side of the six-membered ring are two single bonded amine groups and a double bonded oxygen. Three separate dotted lines extend from these sites to corresponding sites on a second six-membered ring. This ring has points of unsaturation and a nitrogen atom in the bottom right bonding position that is single bonded to a five-membered ring on the right side of the image. This ring is single bonded to a carbon that is single bonded to an oxygen that is single bonded to a phosphorus. The phosphorus is single bonded to two other oxygen atoms and double bonded to a fourth oxygen atom. This group is labeled “5, prime.” The five-membered ring is also bonded on the top side to an oxygen that is bonded to a phosphorus single bonded to two other oxygen atoms and double bonded to a fourth oxygen atom. The upper left oxygen of this group is single bonded to a carbon that is single bonded to a five-membered ring with an oxygen in the bottom bonding position. This ring has a hydroxyl group on its upper right side that is labeled “3, prime” and is bonded on the left side to a nitrogen that is a member of a five-membered ring. This ring is bonded to a six-membered ring and both have points of unsaturation. This ring has a nitrogen on the left side, as well as an amine group, that have two dotted lines leading from them to an oxygen and amine group on a six membered ring. These dotted lines are labeled “Hydrogen bonds.” The six membered ring also has a double bonded oxygen on its lower side and a nitrogen atom on its left side that is single bonded to a five-membered ring. This ring connects to the two phosphate groups mentioned at the start of this to form a large circle. The name “guanine” is written below the lower left side of this image while the name “cytosine” is written on the lower right. The name “thymine” is written above the right side of the image and “adenine” is written on the top right. Three sections are indicated below the images where the left is labeled “Sugar, dash, phosphate backbone,” the middle is labeled “Bases” and the right is labeled “Sugar, dash, phosphate backbone.”
Figure 10. Two separate DNA chains form a double-stranded helix in which the strands are held together via hydrogen bonding. (credit: modification of work by Jerome Walker, Dennis Myts)

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 11).

A large Lewis structure is shown. The top left corner of this structure, labeled “5, prime,” shows a phosphorus atom single bonded to three oxygen atoms, one of which has a superscripted negative charge, and double bonded to a fourth oxygen atom. One of the single bonded oxygen atoms is single bonded to the left corner of a five-membered ring with an oxygen atom at its top point and which is single bonded to an oxygen atom on the bottom left. This oxygen atom is single bonded to a phosphorus atom that is single bonded to two other hydrogen atoms and double bonded to a fourth oxygen atom. The lower left of these oxygen atoms is single bonded to another oxygen atom that is single bonded to a five-membered ring with an oxygen in the upper bonding site. The bottom left of this ring has a hydroxyl group attached to it while the upper right carbon is single bonded to a nitrogen atom that is part of a five-membered ring bonded to a six-membered ring. Both of these rings have points of unsaturation and nitrogen atoms bonded into their structures. On the right side of the six-membered ring are two single bonded amine groups and a double bonded oxygen. Three separate dotted lines extend from these sites to corresponding sites on a second six-membered ring. This ring has points of unsaturation and a nitrogen atom in the bottom right bonding position that is single bonded to a five-membered ring on the right side of the image. This ring is single bonded to a carbon that is single bonded to an oxygen that is single bonded to a phosphorus. The phosphorus is single bonded to two other oxygen atoms and double bonded to a fourth oxygen atom. This group is labeled “5, prime.” The five-membered ring is also bonded on the top side to an oxygen that is bonded to a phosphorus single bonded to two other oxygen atoms and double bonded to a fourth oxygen atom. The upper left oxygen of this group is single bonded to a carbon that is single bonded to a five-membered ring with an oxygen in the bottom bonding position. This ring has a hydroxyl group on its upper right side that is labeled “3, prime” and is bonded on the left side to a nitrogen that is a member of a five-membered ring. This ring is bonded to a six-membered ring and both have points of unsaturation. This ring has a nitrogen on the left side, as well as an amine group, that have two dotted lines leading from them to an oxygen and amine group on a six membered ring. These dotted lines are labeled “Hydrogen bonds.” The six membered ring also has a double bonded oxygen on its lower side and a nitrogen atom on its left side that is single bonded to a five-membered ring. This ring connects to the two phosphate groups mentioned at the start of this to form a large circle. The name “guanine” is written below the lower left side of this image while the name “cytosine” is written on the lower right. The name “thymine” is written above the right side of the image and “adenine” is written on the top right. Three sections are indicated below the images where the left is labeled “Sugar, dash, phosphate backbone,” the middle is labeled “Bases” and the right is labeled “Sugar, dash, phosphate backbone.”
Figure 11. The geometries of the base molecules result in maximum hydrogen bonding between adenine and thymine (AT) and between guanine and cytosine (GC), so-called “complementary base pairs.”

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 function as a template for replication.

Dipole-Induced Dipole Interactions

In the Dispersion Forces section, we learned that at any moment, an atom or nonpolar molecule can experience an instantaneous dipole if its electrons are asymmetrically distributed, and that this instantaneous dipole can induce a dipole in a neighboring atom or molecule. If a molecule with an established dipole were in proximity to a nonpolar molecule, however, it follows that the established dipole will induce a dipole in the nonpolar molecule. This type of interaction between a polar molecule and a nonpolar molecule is known as a dipole-induced dipole interaction (Figure 12).

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This gif shows two molecules, represented as ovals, interacting. One molecule is polar and is labeled with a negative charge on one side and the other side is labeled as positive. The other molecule is a non-polar molecule. As the two molecules come into contact, the oval around the non-polar molecule shifts in the presence of the polar molecule and becomes polar, with one side labeled as positive and the other labeled as negative.
Figure 12. A molecule with a dipole induces a dipole in a non-polar molecule.

Ion-Dipole Interactions

Ion-dipole interactions are the result of the Coulombic electrostatic interactions between an ion and the charged ends of a dipole. Since a dipole consists of an end that is slightly positively charged and another end that is slightly negatively charged, there will be Coulombic attraction and repulsion in this interaction. Consider a cation and a polar molecule. The positive-positive repulsion means the positive end of the dipole is going to be farther away from the cation than the negative end that is attracted (Figure 13).

The image shows a cation, represented as a circle with a plus sign and a polar molecule, represented as an oval with a plug sign on one end and a negative sign on the other. There is an arrow between the cation and the positive end of the dipole that is labeled "repulsive". There is another arrow that points to the cation and the negative end of the dipole labeled "attractive".
Figure 13. There is an attractive and repulsive force between an ion and a dipole. The cation repels the positively charged end of the dipole and attracts the negatively charged end. (credit: “Cnrowley”/Wikimedia Commons)

Ionic Interactions

In Module 2, you learned about ionic compounds. Recall that when an element composed of atoms that readily lose electrons (a metal) reacts with an element composed of atoms that readily gain electrons (a nonmetal), a transfer of electrons usually occurs, producing a positively charged cation and a negatively charged anion. These ions then experience electrostatic attractions (ionic bonds) due to the ions of opposite charge present in the compound.

The Special Role of Hydrogen Bonding in Water

All substances become less dense when they are heated (as they expand slightly) and more dense when they are cooled. Above 4°C water is no exception, but below 4°C as it is cooled further, liquid water actually expands a little and the density decreases slightly. Upon actually freezing at 0 oC, there is a much larger expansion as the volume increases by about 9% with a corresponding drop in density as the liquid water becomes solid ice. In H2O(s) each water molecule makes four hydrogen bonds with its neighbors resulting in a three dimensional array with an open hexagonal structure, more open and therefore with a greater volume compared to the same amount of H2O(l). Between 4 oC and 0 oC, the water molecules are beginning to orient themselves and so the liquid expands as it gets close to actually freezing.  (Figure 9).  You have likely found that if you place a full water bottle in the freezer, the plastic bottle will begin to bulge as the solid water occupies more space than the liquid water.

The expansion of ice as it freezes means that ice floats on top of water. This characteristic is vital to the survival of organisms in cold climates like Wisconsin. During the winter, as Lake Mendota begins to freeze, the surface of the water freezes first; people often enjoy skating or skiing across frozen lakes when the ice is thick enough. The ecosystem is able to survive in the unfrozen water beneath the surface because the ice insulates it, providing an adequate living temperature. If water froze and sank to the bottom, lakes would freeze from the bottom up and become completely frozen during winter. The summer is not long enough in Wisconsin to melt all of the water in the frozen lakes, and fish and other organisms would be unable to survive.

Key Concepts and Summary

The physical properties of condensed matter (liquids and solids) can be explained in terms of the kinetic molecular theory. In a liquid, intermolecular attractive forces hold the molecules in contact, although they still have sufficient KE to move past each other.

Intermolecular attractive forces, collectively referred to as van der Waals forces, are responsible for the behavior of liquids and solids and are electrostatic in nature. Dipole-dipole attractions result from the electrostatic attraction of the partial negative end of one polar molecule for the partial positive end of another. The temporary dipole that results from the motion of the electrons in an atom can induce a dipole in an adjacent atom and give rise to the London dispersion force. London forces increase with increasing molecular size. Hydrogen bonds are a special type of dipole-dipole attraction that results when hydrogen is bonded to one of the three most electronegative elements: F, O, or N.

Glossary

dipole-dipole attraction
intermolecular attraction between two permanent dipoles
dispersion force
(also, London dispersion force or London force) attraction between two rapidly fluctuating, temporary dipoles; significant only when particles are very close together
hydrogen bonding
occurs when exceptionally strong dipoles attract; bonding that exists when hydrogen is bonded to one of the three most electronegative elements: F, O, or N
induced dipole
temporary dipole formed when the electrons of an atom or molecule are distorted by the instantaneous dipole of a neighboring atom or molecule
instantaneous dipole
temporary dipole that occurs for a brief moment in time when the electrons of an atom or molecule are distributed asymmetrically
intermolecular force
noncovalent attractive force between atoms, molecules, and/or ions
polarizability
measure of the ability of a charge to distort a molecule’s charge distribution (electron cloud)
van der Waals force
attractive or repulsive force between molecules, including dipole-dipole, dipole-induced dipole, and London dispersion forces; does not include forces due to covalent or ionic bonding, or the attraction between ions and molecules

Chemistry End of Section Exercises

  1. Open the PhET States of Matter Simulation to answer the following questions:
    1. Select the Solid, Liquid, Gas tab. Explore by selecting different substances, heating and cooling the systems, and changing the state. What similarities do you notice between the four substances for each phase (solid, liquid, gas)? What differences do you notice?
    2. For each substance, select each of the states and record the given temperatures. How do the given temperatures for each state correlate with the strengths of their intermolecular attractions? Explain.
    3. Select the Interaction Potential tab, and use the default neon atoms. Move the Ne atom on the right and observe how the potential energy changes. Select the Total Force button, and move the Ne atom as before. When is the total force on each atom attractive and large enough to matter? Then select the Component Forces button, and move the Ne atom. When do the attractive (van der Waals) and repulsive (electron overlap) forces balance? How does this relate to the potential energy versus the distance between atoms graph? Explain.
  2. Define the following and give an example of each:
    1. dispersion force
    2. dipole-dipole attraction
    3. hydrogen bond
  3. Explain why a hydrogen bond between two water molecules is weaker than a hydrogen bond between two hydrogen fluoride molecules.
  4. Identify the intermolecular forces present in the following solids:
    1. CH3CH2OH
    2. CH3CH2CH3
    3. CH3CH2Cl

Answers to Chemistry End of Section Exercises

  1. (a) Similarities: Solid and liquid are more dense compared to gas. The higher temperature the system has, the faster the molecules move. Differences: For water, the liquid state is denser than solid state; while other solids are denser than their liquid state. The various substances reach each state at different temperatures.
    (b) Neon: -259 °C (solid), -246 °C (liquid), -217 °C (gas);  Argon: -230 °C (solid), -187 °C (liquid), -84 °C (gas);  Oxygen: -246 °C (solid), -216 °C (liquid), -160 °C (gas);  Water: -127 °C (solid);, 13 °C (liquid), 156 °C (gas).  Solids have the strongest IMFs and gases have the weakest. For each state, IMF decreases with increasing temperature.
    (c) When the distance between two atoms is greater than ε , the total force on each atom attractive and large enough to matter. When the distance between two atoms is ε , the attractive (van der Waals) and repulsive (electron overlap) forces balance. At this point, the whole system has the lowest potential energy.
  2. (a) Dispersion forces occur as an atom develops a temporary dipole moment when its electrons are distributed asymmetrically about the nucleus. This structure is more prevalent in large atoms such as argon or radon. A second atom can then be distorted by the appearance of the dipole in the first atom. The electrons of the second atom are attracted toward the positive end of the first atom, which sets up a dipole in the second atom. The net result is rapidly fluctuating, temporary dipoles that attract one another (example: Ar).
    (b) A dipole-dipole attraction is a force that results from an electrostatic attraction of the positive end of one polar molecule for the negative end of another polar molecule (example: ICI molecules attract one another by dipole-dipole interaction).
    (c) Hydrogen bonds form whenever a hydrogen atom is bonded to one of the more electronegative atoms, such as a fluorine, oxygen, or nitrogen atom. The electrostatic attraction between the partially positive hydrogen atom in one molecule and the partially negative atom in another molecule gives rise to a strong dipole-dipole interaction called a hydrogen bond (example: HF • • • HF).
  3. The hydrogen bond between two hydrogen fluoride molecules is stronger than that between two water molecules because the electronegativity of F is greater than that of O. Consequently, the partial negative charge on F in HF is greater than that on O in H2O.
  4. (a) hydrogen bonding, dipole-dipole attraction, and dispersion forces; (b) dispersion forces; (c) dipole-dipole attraction and dispersion forces
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Chem 103/104 Resource Book by crlandis and Chem 104 Textbook Team is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.