2.6 Melting Points, Boiling Points, and Intermolecular Interactions

Melting points and boiling points can provide a wealth of information about intermolecular interactions. To help you see why, study Figure 2-11 to review what the different phases of a compound look like on the molecular level.

A three-part illustration shows the molecular arrangement in the three different states of water. The first part shows a block of ice with a microscopic view depicting closely spaced molecules of water. A note below it reads, �Solid: Intermolecular interactions are maximized.� An arrow labeled �Melting� leads to the next part of the illustration. The second part shows a droplet of water with a microscopic view depicting loosely spaced molecules of water. A note below it reads, �Liquid: Intermolecular interactions are less substantial.� An arrow labeled �Boiling� leads to the final part of the illustration. The third part shows water boiling in a pan with a microscopic view depicting the molecules of water dispersed widely. A note below it reads, �Gas: Intermolecular interactions are effectively absent.� The caption reads, Microscopic structure of the three phases of matter: a. In a crystalline solid, molecules or ions form a well-ordered structure called a crystal lattice, in which movement is limited to vibration and intermolecular forces are maximized. b. In a liquid, molecules are free to move around, because intermolecular forces are somewhat less substantial. c. In the gas phase, molecules are so far apart that intermolecular forces are effectively absent.
FIGURE 2-11 Microscopic structure of the three phases of matter (a) In a crystalline solid, molecules or ions form a well-ordered structure called a crystal lattice, in which movement is limited to vibration and intermolecular forces are maximized. (b) In a liquid, molecules are free to move around, because intermolecular forces are somewhat less substantial. (c) In the gas phase, molecules are so far apart that intermolecular forces are effectively absent.

●  Solids consist of atoms, ions, or molecules that are in contact with one another and are essentially immobile. This allows intermolecular interactions to be maximized.

●  In a liquid, the species are also in close contact, but they can rotate and slide past one another, so intermolecular interactions are less substantial than in a solid.

●  In a gas, the species are far apart and are effectively isolated from one another, so they can move freely. Intermolecular interactions are essentially nonexistent.

Melting, therefore, decreases the intermolecular interactions that exist in the solid phase, and boiling effectively overcomes the remaining intermolecular interactions that exist in the liquid phase. Consequently, as the strength of the intermolecular interactions that exist in a particular substance increases, more energy (in the form of heat) is required for the substance to melt or boil.

●  Melting points increase as the intermolecular interactions in a solid increase.

●  Boiling points increase as the intermolecular interactions in a liquid increase.

Let’s now apply these ideas to interpret the relative strengths of various types of intermolecular interactions.

2.6a Ion–Ion Interactions

Of all the compounds in Table 2-4 on page 81, sodium methanoate (sodium formate, Na+–OCH═O) has the highest melting point and boiling point, suggesting that it has particularly strong intermolecular attractions in both its solid and liquid phases. Sodium methanoate is an ionic compound, composed of Na+ and ions held together (as we saw in Chapter 1) by the electrostatic attraction of oppositely charged ions, called ionic bonds or, more generally, ion–ion interactions.

Ion–ion interactions are the strongest intermolecular interactions because ions have very high concentrations of positive and negative charge.

2.6b Dipole–Dipole Interactions

Referring again to Table 2-4, notice that the compounds with significant dipole moments—methanoic acid (formic acid), ethanol, ethanal (acetaldehyde), and methoxymethane (dimethyl ether)—have boiling points and melting points that are significantly higher than those of the remaining compounds, which are essentially nonpolar.

Polar molecules are attracted to each other more strongly than similar nonpolar molecules.

The basis for this behavior is dipole–dipole interactions. Dipole–dipole interactions arise because the positive end of one molecule’s net dipole is attracted to the negative end of another’s, as shown in Figure 2-12. Therefore:

Electrostatic potential map shows the attraction between two dimethyl ether molecules. The two molecules are placed one above the other, and each molecule is represented by a ball-and-stick model in an electrostatic cloud. The structure of each molecule shows an oxygen atom bonded to two methyl groups, one on each side in a bent fashion. The electrostatic map is shaded red at the central upper portion representing the oxygen atom, and blue at the lower ends representing the hydrogen atoms. The central portion, from top to bottom, is shaded in yellow, green, and turquoise. An arrow labeled �partial positive� points to the blue regions of the first molecule, and an arrow labeled �partial negative� points to the red region of the second molecule. A two-way arrow labeled �attraction� connects these two regions. The caption reads, �Dipole�dipole interaction: The dominant intermolecular force between two molecules of dimethyl ether is a dipole�dipole interaction. The positive end of one ether molecule attracts the negative end of the other.�
FIGURE 2-12 Dipole–dipole interaction The dominant intermolecular force between two molecules of dimethyl ether is a dipole–dipole interaction. The positive end of one ether molecule attracts the negative end of the other.

All other factors being equal, the strength of dipole–dipole interactions increases as the dipole moment increases.

Connections Ethanol is probably most commonly known as the alcohol found in alcoholic beverages, but is also an important organic solvent, fuel, and antiseptic.

Notice in Table 2-4 (p. 81), for example, that the highly polar ethanal (CH3CH═O) has a higher melting point and a higher boiling point than the less polar dimethyl ether (CH3OCH3).

YOUR TURN 2.10
SHOW ANSWERS

The boiling point of CH3CH2F is –37.1 °C. How does this compare to the boiling point of CH3CH═O? Which compound, therefore, is more polar?

Condensed structural formulae of CH3CH2F and CH3CH double bond O. The structure of CH3CH2F shows a central carbon atom single bonded to two hydrogen atoms. It is also single bonded to a methyl group by a dash bond and single bonded to a fluorine atom by a dash bond. A thin short red arrow points from the carbon atom toward the fluorine atom. The label below this structure reads, �Boiling point equals minus 37.1 degree Celsius�. The structure of CH3CH double bond O shows a central carbon atom single bonded to a hydrogen atom and a methyl group. It is also double bonded to an oxygen atom. A thin long red arrow points from the carbon atom toward the oxygen atom. The label below this structure reads, �Boiling point equals plus 20 degree Celsius. Higher boiling point, large dipole moment�.

Although dipole–dipole interactions can be quite strong, they involve the attraction of partial charges, which are smaller in magnitude than the full charges in ion–ion interactions. Therefore:

Dipole–dipole interactions are generally much weaker than ion–ion interactions.

As a result, compounds such as ethanal (acetaldehyde) melt and boil at lower temperatures than ionic compounds.

Solved Problem 2.6

Which compound has the higher melting point, NaOCH2CH3 or CH3CH═O?

Think
SHOW SECTION

What is the strongest intermolecular interaction in NaOCH2CH3? In CH3CH═O? Which type of interaction is stronger? How does that affect the melting point?

Solve
SHOW SECTION

NaOCH2CH3 is an ionic compound. It consists of CH3CH2O– and Na+, which are held together by ion–ion interactions. CH3CH═O, on the other hand, is a polar covalent compound, so it experiences dipole–dipole interactions. Because ion–ion interactions are stronger than dipole–dipole interactions, the melting point of NaOCH2CH3 is higher than the melting point of CH3CH═O.

problem 2.7 Which compound has the higher boiling point, NaOCH2CH3 or CH3CH═O?

problem 2.8 Which compound has the higher boiling point, CH4 or CH3F?

2.6c Hydrogen Bonding

An illustration shows a hydrogen bond donor and a hydrogen bond acceptor. The illustration shows an atom �D� carrying a delta negative charge bonding with a hydrogen atom carrying a delta positive charge. A dashed line connects the hydrogen atom to another atom �A� carrying a delta negative charge. The atom �A� has a lone pair of electrons. An arrow pointing to atom �D� has a note that reads: �Hydrogen-bond donor (D equals nitrogen, oxygen, or fluorine)� An arrow pointing to atom �A� has a note that reads: �Hydrogen-bond acceptor (A equals nitrogen, oxygen, or fluorine)� The caption reads, Hydrogen bonds: A hydrogen bond consists of a hydrogen-bond donor D and a hydrogen-bond acceptor A with a lone pair. If the D and A atoms are uncharged, they must be nitrogen, oxygen, or fluorine for the hydrogen bond to be substantial.
FIGURE 2-13 Hydrogen bonds A hydrogen bond consists of a hydrogen-bond donor (D—H) and a hydrogen-bond acceptor (: A). If the D and A atoms are uncharged, they must be nitrogen, oxygen, or fluorine for the hydrogen bond to be substantial.

Dipole–dipole interactions alone cannot account for some of the data in Table 2-4 (p. 81). Methanoic acid (formic acid, HCO2H) and ethanol (CH3CH2OH), for example, have substantially higher melting points and boiling points than ethanal (acetaldehyde, CH3CH═O), despite the fact that ethanal’s net dipole moment is the highest of the three. These apparent anomalies arise because methanoic acid and ethanol can form hydrogen bonds, whereas acetaldehyde cannot.

Each hydrogen bond (H bond) requires a hydrogen-bond donor and a hydrogen-bond acceptor, as shown in Figure 2-13.

●  A hydrogen-bond donor is a covalent bond, D—H, where D is a highly electronegative atom, such as N, O, or F.

●  A hydrogen-bond acceptor, A, can be any atom with a large concentration of negative charge and a lone pair of electrons. For the hydrogen bond to be substantial, an uncharged H-bond acceptor must be F, O, or N.

The high electronegativity of the F, O, or N atom ensures that the D—H bond is highly polar, giving the H atom a large partial positive charge. This sets up a strong attraction between the H atom and the oppositely charged acceptor.

The actual H bond is created when a lone pair of electrons on the H-bond acceptor is shared with the hydrogen atom of the H-bond donor. The H bond is often depicted by a dashed line (different from the dash bond that is part of dash–wedge notation). Figure 2-14 shows how hydrogen bonding might be depicted between two molecules of ethanol (Fig. 2-14a) and between two molecules of methanoic acid (Fig. 2-14b). Notice that either O atom in methanoic acid can serve as a H-bond acceptor.

Lewis structures show the hydrogen bonding between two ethanol molecules and two methanoic acid molecules. The first illustration shows two ethanol molecules. Each consists of two carbon atoms bonded by a single bond. Carbon 1 is bonded to two hydrogen atoms and a hydroxyl group, and carbon 2 is bonded to three hydrogen atoms. In the first molecule, the oxygen atom in the hydroxyl group carries a partial negative charge, and the hydrogen atom in the group carries a partial positive charge. In the second molecule, the oxygen atom in the hydroxyl group carries a partial negative charge and has two lone pairs of electrons. A dashed line labeled �a hydrogen bond� connects this oxygen atom with the hydrogen in the hydroxyl group of the first molecule. The oxygen atom in the first molecule is the hydrogen-bond donor, and the oxygen with the lone pairs in the second molecule is the hydrogen-bond acceptor. The second illustration shows two sets of two methanoic acid molecules. Each molecule consists of a carbon atom bonded to a hydrogen atom and a hydroxyl group by single bonds and to an oxygen atom by a double bond. In the first molecule of the first set, the oxygen atom in the hydroxyl group carries a partial negative charge, and the hydrogen atom in the group carries a partial positive charge. In the second molecule of the first set, the oxygen atom in the hydroxyl group carries a partial negative charge and has two lone pairs of electrons. A dashed line labeled �a hydrogen bond� connects this oxygen atom with the hydrogen in the hydroxyl group of the first molecule. In the first molecule of the second set, the oxygen atom in the hydroxyl group carries a partial negative charge, and the hydrogen atom in the group carries a partial positive charge. In the second molecule of the second set, the oxygen atom double-bonded to the carbon carries a partial negative charge and has two lone pairs of electrons. A dashed line labeled �a hydrogen bond� connects this oxygen atom with the hydrogen in the hydroxyl group of the first molecule. The caption reads, Hydrogen bonding between ethanol molecules and between methanoic acid molecules: a. A hydrogen bond can form between two molecules of ethanol because the H atom is covalently bonded to an O atom in one molecule and is attracted to the second molecule�s O atom, which has a lone pair of electrons and a significant partial negative charge. The H bond is indicated by a dashed line. b. Two different H bonds between two molecules of methanoic acid (formic acid).
FIGURE 2-14 Hydrogen bonding between ethanol molecules and between methanoic acid molecules (a) A hydrogen bond can form between two molecules of ethanol because the H atom is covalently bonded to an O atom in one molecule and is attracted to the second molecule’s O atom, which has a lone pair of electrons and a significant partial negative charge. The H bond is indicated by a dashed line. (b) Two different H bonds between two molecules of methanoic acid (formic acid).

YOUR TURN 2.11
SHOW ANSWERS

Circle and label the H-bond donor and H-bond acceptor in each H bond shown in Figure 2-14b.

An illustration shows two sets of Lewis structures with H-bond donor and H-bond acceptor between two methanoic acid molecules. Each molecule consists of a carbon atom bonded to a hydrogen atom and a hydroxyl group by single bonds and to an oxygen atom by a double bond. In the first molecule of the first set, the oxygen atom in the hydroxyl group carries a partial negative charge, and the hydrogen atom in the group carries a partial positive charge. This bonding is circled and labeled as �H-bond donor�. In the second molecule of the first set, the oxygen atom in the hydroxyl group carries a partial negative charge and has two lone pairs of electrons. This bonding is circled and labeled as �H-bond acceptor�. A dashed line labeled �a hydrogen bond� connects this oxygen atom with the hydrogen in the hydroxyl group of the first molecule. In the first molecule of the second set, the oxygen atom in the hydroxyl group carries a partial negative charge, and the hydrogen atom in the group carries a partial positive charge. This bonding is circled and labeled as �H-bond donor�. In the second molecule of the second set, the oxygen atom double-bonded to the carbon carries a partial negative charge and has two lone pairs of electrons. This bonding is circled and labeled as �H-bond acceptor�. A dashed line labeled �a hydrogen bond� connects this oxygen atom with the hydrogen in the hydroxyl group of the first molecule.

Hydrogen bonding in uncharged species involves only partial charges, so like dipole–dipole interactions:

Hydrogen bonding is weaker than ion–ion interactions.

Hydrogen bonds, however, are distinct from dipole–dipole interactions for two main reasons: (1) fluorine, oxygen, and nitrogen are highly electronegative, so the partial charges involved are large, and (2) the hydrogen atom is very small, which allows the partial positive charge on hydrogen to be very close to the partial negative charge on the H-bond acceptor. For these reasons:

Hydrogen bonding is often (but not always) stronger than dipole–dipole interactions.

problem 2.9 Which pair of species will give rise to the strongest intermolecular interactions? The weakest?

Condensed skeletal formulas show three pairs of species. The first pair shows two structures, the first of which is a central carbon atom bonded to two other carbon atoms and an NH ion with a negative charge. The second structure shows a lithium ion with a positive charge. The second pair shows two structures, each of which is a zigzag line with two crests and two troughs, with a nitrogen atom at the third position. A double bond exists between carbon 2 and the nitrogen atom. The third pair shows two structures, each of which has a central carbon bonded to two other carbon atoms and an amine group.

It is important to be able to gauge the extent of hydrogen bonding among molecules of a particular compound—that is, the collective strength of all the hydrogen bonds present. As a general rule:

The extent of hydrogen bonding increases as the total number of potential H-bond donors and acceptors increases between the species involved.

With more potential donor–acceptor pairs involving two molecules, there are more ways in which hydrogen bonding can take place.

This explains why methanoic acid (formic acid) has a higher boiling point and melting point than ethanol. As shown in Figure 2-15a, a single molecule of ethanol has one potential H-bond donor and one potential H-bond acceptor. In a pair of ethanol molecules undergoing the intermolecular interaction, therefore, there are two potential donors and two potential acceptors. In a single molecule of methanoic acid (formic acid, Fig. 2-15b), on the other hand, there is one potential donor and two potential acceptors. In a pair of methanoic acid molecules, therefore, there are two potential donors and four potential acceptors. As a result, there are more ways in which hydrogen bonding can take place in methanoic acid than in ethanol.

Lewis structures show an ethanol molecule and a methanoic acid molecule. The structure of ethanol consists of two carbon atoms bonded by a single bond. Carbon 1 is bonded to two hydrogen atoms and a hydroxyl group, and carbon 2 is bonded to three hydrogen atoms. The oxygen atom in the hydroxyl group has two lone pairs of electrons. An arrow labeled �potential H-bond acceptor� points to the oxygen atom, and the O-H bond is labeled �potential H-bond donor.� The structure of methanoic acid, also known as formic acid, shows a carbon atom bonded to a hydrogen atom and a hydroxyl group by single bonds and to an oxygen atom by a double bond. Each oxygen atom has two lone pairs of electrons and is labeled �potential H-bond acceptor.� The O-H bond is labeled �potential H-bond donor.� The caption reads, Potential hydrogen bond donors and acceptors: a. In a molecule of ethanol, the OH bond is a potential H-bond donor and the O atom is a potential H-bond acceptor. b. In a molecule of formic acid, the OH bond is a potential H-bond donor and each O atom is a potential H-bond acceptor.
FIGURE 2-15 Potential hydrogen bond donors and acceptors (a) In a molecule of ethanol, the OH bond is a potential H-bond donor and the O atom is a potential H-bond acceptor. (b) In a molecule of formic acid, the OH bond is a potential H-bond donor and each O atom is a potential H-bond acceptor.

YOUR TURN 2.12
SHOW ANSWERS

High levels of cholesterol, a naturally occurring steroid, have been linked to cardiovascular disease. Octanoic acid (caprylic acid) is a fatty acid found in milk and is used commercially to manufacture perfumes and dyes. Identify the number of potential H-bond donors and acceptors in cholesterol and in octanoic acid.

Two condensed skeletal formulas show a molecule of cholesterol and octanoic acid. The dash-wedge structure of cholesterol shows two hexagonal six-carbon rings fused together. The second ring is fused with a third hexagonal ring in a bent fashion. The third ring is fused with a pentagonal five-carbon ring beside it. Double bonds exist between carbons 5 and 6. Carbon 3 is bonded to a hydroxyl group, and carbons 10, 13, and 25 are each bonded to a methyl group. Carbon 17 is bonded to carbon 20, which is in the second position in a zigzag seven-carbon chain. Solid wedges connect carbon 3 and the hydroxyl group, carbon atoms 10 and 19, carbon atom 8 and a hydrogen atom, and carbons 13 and 18. Dashed wedges connect carbon atom 9 and a hydrogen atom, carbon atom 14 and a hydrogen atom, carbon atom 17 and a hydrogen atom, and carbons 20 and 21. The structure of octanoic acid, also known as caprylic acid, shows a zigzag eight-carbon chain where carbon 1 is part of a carboxyl group.

Cholesterol has one potential H-bond acceptor (the O atom) and one potential donor (the OH bond). Octanoic acid has two potential H-bond acceptors (each O atom) and one potential donor (the OH bond).

problem 2.10 How many potential H-bond donors and H-bond acceptors are there in each of the following molecules?

Skeletal structural formulas of five different molecules. The first structure shows a benzene ring with two hydroxyl groups bonded at para positions. The second structure shows a zigzag four-carbon chain where carbon 1 is part of a carboxyl group and carbon 2 is bonded to an amine group. The third structure shows a five-carbon chain where carbon 1 is bonded to three fluorine atoms and carbon 2 is bonded to an oxygen atom by a double bond. The fourth structure shows a nitrogen atom bonded to three ethyl chains. The fifth structure shows a hexagonal ring with one carbon atom bonded to an aldehyde group.

problem 2.11 Which functional groups in Table 1-6 (p. 35) possess at least one H-bond acceptor but no H-bond donors? Which functional groups possess at least one H-bond donor and one H-bond acceptor? Which functional groups possess no H-bond donors and no H-bond acceptors?

Solved Problem 2.12

1,2-Ethanediol (ethylene glycol, A) is used as an automotive antifreeze. Hydroxyacetaldehyde (B) is believed to be an intermediate in the metabolism of proteins and carbohydrates. Which of these compounds would you expect to have a higher boiling point? Why?

Skeletal structural formula of two molecules labeled A and B. The first structure shows a two-carbon chain where each carbon is bonded to a hydroxyl group. The second structure a two-carbon chain where one carbon is bonded to a hydroxyl group by a single bond and the other carbon is bonded to an oxygen atom by a double bond.

Think
SHOW SECTION

What is the most important intermolecular interaction that will occur between two molecules of A? Between two molecules of B? Which interaction is stronger, and how would it affect the boiling points of A and B?

Solve
SHOW SECTION

Both A and B are polar molecules, so dipole–dipole interactions should be present in a pair of each type of molecule. However, hydrogen bonding, which is often stronger than dipole–dipole interactions, is also possible: Each molecule contains at least one potential H-bond donor (an OH bond) and at least one potential H-bond acceptor (an O atom). To estimate which pair of molecules has greater hydrogen bonding, we count the total number of potential H-bond donors and acceptors. In two molecules of A, there are four potential donors and four potential acceptors, because there are two donors and two acceptors from each molecule. In two molecules of B, there are two potential donors and four potential acceptors, because there are one donor and two acceptors from each molecule. As a result, we would expect compound A to have more hydrogen bonding, and thus a higher boiling point.

problem 2.13 Which compound, C or D, would you expect to have a higher boiling point? Why?

Skeletal structural formula of two molecules labeled C and D. The first structure shows a three-carbon chain where carbon atoms 1 and 3 are each bonded to an amine group. The second structure a two-carbon chain where one carbon is bonded to an amine group by a single bond and the other carbon is bonded to a nitrogen atom by a triple bond.

The strength of hydrogen bonding also depends on the concentrations of charge in the H-bond donors and H-bond acceptors. Ethanamine (CH3CH2NH2), for example, has a lower boiling point than ethanol (CH3CH2OH) because N is less electronegative than O, resulting in smaller concentrations of charge in ethanamine than in ethanol (Fig. 2-16).

A two-part illustration shows hydrogen bonding in ethanol and ethanamine. The first part shows the hydroxyl group segments of two ethanol molecules, bonded to the two-carbon chain. The oxygen atom in the hydroxyl group of the first molecule carries a partial negative charge, while the hydrogen atom in the group carries a partial positive charge. The oxygen atom in the hydroxyl group of the second molecule carries a partial negative charge and has two lone pairs of electrons. A dashed line representing a hydrogen bond connects this oxygen atom with the hydrogen in the hydroxyl group of the first molecule. An arrow with a note, �O is more electronegative than N, so this hydrogen bond is stronger,� points to the dashed line. A note below this part of the structure reads, �Ethanol: Boiling point equals 78 degrees Celsius.� The second part shows the amine group segments of two ethanamine molecules, bonded to the two-carbon chain. The nitrogen atom in the amine group of the first molecule carries a partial negative charge, while one of the hydrogen atoms in the group carries a partial positive charge. The nitrogen atom in the amine group of the second molecule carries a partial negative charge and has one lone pair of electrons. A dashed line representing a hydrogen bond connects this nitrogen atom with the hydrogen in the amine group of the first molecule. A note below this part of the structure reads, �Ethanamine: Boiling point equals 17 degrees Celsius.� The caption reads, �Hydrogen bonding and electronegativity: Hydrogen bonding is stronger in ethanol than in ethanamine because O is more electronegative than N, thus giving rise to larger concentrations of positive and negative charges. As a result, the boiling point of ethanol is higher than the boiling point of ethanamine.�
FIGURE 2-16 Hydrogen bonding and electronegativity Hydrogen bonding is stronger in ethanol than in ethanamine because O is more electronegative than N, thus giving rise to larger concentrations of positive and negative charges. As a result, the boiling point of ethanol is higher than the boiling point of ethanamine.

problem 2.14 Which H bond would you expect to be stronger, E or F? Why?

Lewis structures of two pairs of molecules labeled E and F. The first pair shows two molecules, each with a nitrogen atom carrying a lone pair of electrons bonded to a hydrogen atom and two methyl groups. A dashed line representing a hydrogen bond connects the nitrogen atom in the second molecule with the hydrogen atom in the first. The second pair shows two molecules, each with a fluorine atom carrying three lone pairs of electrons bonded to a hydrogen atom. A dashed line representing a hydrogen bond connects the fluorine atom in the second molecule with the hydrogen atom in the first.

2.6d Induced Dipole–Induced Dipole Interactions (London Dispersion Forces)

Nonpolar molecules must attract each other through intermolecular interactions; otherwise, nonpolar compounds could never condense from gas to liquid. The dominant intermolecular interaction between nonpolar molecules is called induced dipole–induced dipole interactions, or London dispersion forces.

How do induced dipole–induced dipole interactions arise? Although the average electron distribution in a molecule such as propane (CH3CH2CH3) does not give rise to a significant permanent dipole (Fig. 2-17a), any electron cloud can be distorted—that is, any electron cloud is polarizable. Electrons are constantly moving around and at some instant in time, there may be more electrons on one side of the molecule than there are on the other. The extra electrons on that one side give rise to an instantaneous dipole (Fig. 2-17b), which can alter the electron distribution on a second molecule by repelling or attracting nearby electrons. The second molecule then develops an induced dipole that is attracted to the first molecule (Fig. 2-17c).

Electrostatic potential maps show the three stages in formation of an induced dipole-induced dipole interaction in two propane molecules. The illustration is divided into three parts, each shown in a box. Each part shows the ball-and-stick models of two propane atoms placed one over the other, enclosed in an electrostatic cloud each. In the first part, the cloud for both the molecules is shaded in blue around the vertices representing the carbon atoms, and green in the central portion. A note below the box reads, �On average, propane is nonpolar.� In the second part, the electrostatic cloud for the first molecule is shaded in blue around the vertices representing the carbon atoms, and green in the central portion. The map for the second molecule is shaded in red at the upper portion, representing the central carbon atom. The areas around the other carbon atoms are shaded in blue, and the central portion is green. A note below the box reads, �An instantaneous dipole develops on one propane molecule.� In the third part, the cloud for both the molecules is shaded in red at the upper portion, representing the central carbon atom. The areas around the other carbon atoms are shaded in blue, and the central portion is green. Two arrows, both labeled �opposite partial charges attract,� point to the red upper portion of the second molecule and the blue lower portion of the first. A note below the box reads, �An induced dipole develops on the other molecule.� The caption reads, Induced dipole�induced dipole interaction: a. On average, two isolated molecules of propane are nonpolar. b. Electrons are not static, however, so electron density can build up on one side of a molecule at some instant in time, resulting in a temporary dipole. c. That temporary dipole, in turn, can alter the electron distribution of a second, adjacent molecule, giving the second molecule an induced dipole. The oppositely charged ends of these induced dipoles attract one another.
FIGURE 2-17 Induced dipole–induced dipole interaction (a) On average, two isolated molecules of propane are nonpolar. (b) Electrons are not static, however, so electron density can build up on one side of a molecule at some instant in time, resulting in a temporary dipole. (c) That temporary dipole, in turn, can alter the electron distribution of a second, adjacent molecule, giving the second molecule an induced dipole. The oppositely charged ends of these induced dipoles attract one another.

Although they are most pronounced in nonpolar molecules, induced dipole–induced dipole interactions are present when any two species interact.

To gain a sense of the relative strength of induced dipole–induced dipole interactions, notice in Table 2-4 (p. 81) that the nonpolar compounds tend to have the lowest boiling points and melting points. This suggests that:

Induced dipole–induced dipole interactions are generally the weakest of all intermolecular forces.

Their strength is highly variable, however, and can sometimes be quite significant, as shown in Table 2-5. A small molecule like CH4 is a gas at room temperature, indicating that its induced dipole–induced dipole interactions are extremely weak. A much heavier molecule like I2, on the other hand, is a solid at room temperature, indicating that its induced dipole–induced dipole interactions are quite strong—even stronger than the hydrogen bonding in water, which is a liquid at room temperature!

Table 2-5 is titled, melting points and boiling points of nonpolar compounds. The table has eight columns and six rows. The rows represent different molecules. The columns represent the total number of electrons, and the melting and boiling points for the molecules. Data are included in the accompanying table. | Molecule Total number of electrons Melting point (degree Celsius) Boiling point (degree Celsius) Molecule Total number of electrons Melting point (degree Celsius) Boiling point (degree Celsius) | Methane, with carbon bonded to four hydrogen atoms 10 Negative 182 Negative 161 Dimethylpropane represented by an inverted V with two single bonds arising from the tip. 42 Negative 17 10 | Ethane, with a two-carbon chain where each carbon is bonded to three hydrogen atoms. 18 Negative 183 Negative 89 Butane, represented by a zigzag line with two crests and three troughs. 42 Negative 130 36 | Propane, represented by an inverted V. 26 Negative 188 Negative 42 Bromine molecule 70 Negative 7 59 | Chlorine molecule 34 Negative 101 Negative 34 Iodine molecule 106 114 184 | Butane, represented by a zigzag line with two crests and two troughs. 34 Negative 138 Negative 1 Empty cell Empty cell Empty cell Empty cell

Connections Elemental iodine has use as a disinfectant. Because of its deep violet color and reactivity with compounds such as alkenes, I2 is used in analytical chemistry to determine end points of titrations.

The precise strength of induced dipole–induced dipole interactions in a particular species depends on its polarizability, which can be defined as the ease with which its electron cloud is distorted.

Climbing Like Geckos

In the opening to this chapter, we highlighted the secret behind the ability of geckos to climb just about any surface effortlessly: dispersion forces. Geckos have a specialized hierarchical structure in their toes, culminating in large numbers of very small hairs less than 200 nm in diameter, called setae. This creates a very large contact surface area for each toe, giving a gecko’s foot an adhesive pressure up to about 30 pounds per square inch on glass. This means that the gecko can, in principle, cling to glass by just one toe.

Scientists have known the gecko’s secret for many years but were unable to harness it until recently, when Professors Duncan Irschick and Alfred Crosby of the University of Massachusetts Amherst unveiled Geckskin in 2014. Geckskin is made from polydimethylsiloxane, a common polymer, and the adhesive pad mimics the structure of the gecko’s toe to maximize its contact surface area. One of the keys to the success of Geckskin is that it is woven into a synthetic tendon to maintain both stiffness and rotational freedom. And successful it is. A piece about the size of an index card has been shown to hold 700 pounds on a smooth wall and is easily removed without leaving any residue. Inspired by this, the Defense Advanced Research Projects Agency (DARPA) developed a system that enables an adult human weighing more than 200 pounds to scale a 25-foot vertical glass wall.

A photo shows a climber in a helmet and climbing harness, using adhesive pads to scale a surface.

Polarizability tends to increase as the number of total electrons increases.

A species with relatively few electrons like CH4 (melting point = –182 °C, boiling point = –161 °C), which has 10 total electrons, is not very polarizable. This is why its induced dipole–induced dipole interactions are weak. On the other hand, pentane (CH3CH2CH2CH2CH3), with 42 electrons, is significantly more polarizable, giving it a higher melting point (–130 °C) and boiling point (36 °C). Iodine (I2) has 106 electrons, so its melting point (114 °C) and boiling point (184 °C) are even higher.

YOUR TURN 2.13
SHOW ANSWERS

On the graph provided, plot boiling point as a function of total number of electrons for the straight-chain alkanes in Table 2-5 (i.e., methane, ethane, propane, butane, and pentane). What do you notice?

A line graph shows total number of electrons along the x axis and boiling point in degree Celsius along the y axis. The x axis is numbered from 0 to 50 in intervals of 10 and the y axis is numbered from negative 250 to 150 in intervals of 50. The approximate plot from the graph reads: 10, negative 140; 18, negative 90; 26, negative 45; 34, 0; 42, 40.

Boiling point increases as the number of electrons increases.

A graph showing boiling point in degrees Celsius along the vertical axis and total number of electrons along the horizontal axis. The vertical axis is numbered from negative 250 to 150, in intervals of 50, and the horizontal axis is numbered from 0 to 50 in intervals of 10.

Another important factor governing the strength of induced dipole–induced dipole interactions is contact surface area.

Induced dipole–induced dipole interactions tend to increase in strength as the contact surface area increases.

For instance, pentane [CH3(CH2)3CH3] and dimethylpropane [neopentane, C(CH3)4] have the same formula (C5H12), and thus the same number of electrons, but pentane has a more extended shape than the relatively compact dimethylpropane. As a result, two molecules of pentane have a greater surface area available for interaction than two molecules of dimethylpropane do, as shown in Figure 2-18. Greater contact surface area means that pentane has more effective induced dipole–induced dipole interactions and a higher boiling point than dimethylpropane.

Two illustrations show the condensed structural formulas and space-filling models of a pair of pentane molecules and dimethylpropane molecules. The first illustration shows the structure of two molecules of pentane, one above the other, with the space-filling models in between. The structure shows a five-carbon chain with carbons 1 and 5 each bonded to three hydrogen atoms, and carbons 2, 3, and 4 each bonded to two hydrogen atoms. The second illustration shows the structure of two molecules of dimethylpropane, one above the other, with the space-filling models in between. The structure shows a three-carbon chain with carbons 1 and 3 each bonded to three hydrogen atoms, and carbon 2 bonded to two methyl groups, one pointing away from the reader and one pointing toward. The caption reads, �Contact surfacearea and induced dipole�induceddipole interactions: The contactsurface area between two moleculesof pentane a is greater thanthat between two molecules ofdimethylpropane b. The result isstronger induced dipole�induceddipole interactions in pentane and acorrespondingly higher boiling point.�
FIGURE 2-18 Contact surface area and induced dipole–induced dipole interactions The contact surface area between two molecules of pentane (a) is greater than that between two molecules of dimethylpropane (b). The result is stronger induced dipole–induced dipole interactions in pentane and a correspondingly higher boiling point.

YOUR TURN 2.14
SHOW ANSWERS

In Figure 2-18, shade in the contact surface area for each pair of molecules. Indicate which pair of molecules has greater contact surface area.

As shown below, pentane has a greater intermolecular contact surface area than dimethylpropane.

The condensed structural formulae and space-filling models of a pair of pentane molecules and dimethylpropane molecules with their contact surface area highlighted. The first illustration shows the structure of two molecules of pentane, one above the other, with the space-filling models in between. The structure shows a five-carbon chain with carbons 1 and 5 each bonded to three hydrogen atoms, and carbons 2, 3, and 4 each bonded to two hydrogen atoms. The space between the two pentane structures is highlighted as a wide rectangle. The caption reads, �(a) Pentane: Boiling point equals 36 degree Celsius�. The second illustration shows the structure of two molecules of dimethylpropane, one above the other, with the space-filling models in between. The structure shows a three-carbon chain with carbons 1 and 3 each bonded to three hydrogen atoms, and carbon 2 bonded to two methyl groups, one pointing away from the reader and one pointing toward. The space between the two dimethylpropane structures is highlighted as a small rectangle. The caption reads, �(b) Dimethylpropane: Boiling point equals 10 degree Celsius�.

problem 2.15 Which molecule, in each pair, would you expect to have a higher boiling point? Why?

A two-part illustration shows skeletal structures of two pairs of molecules. The first part shows the structure of 2,2-Dimethylbutane and 2-Methylpentane. The first structure shows a zigzag four-carbon chain with two methyl groups bonded to the second carbon atom. The second structure shows a zigzag five-carbon chain with a methyl group bonded to the second carbon atom. The second part shows the structure of 1,1-Dimethylcyclopropane and 1,2-Dimethylcyclopropane. The first structure shows a triangular three-carbon ring with two methyl groups bonded to the first carbon atom. The second structure shows a triangular three-carbon ring with the first and second carbon atoms each bonded to a methyl group.