Beer and Soap - Absolutely Small: How Quantum Theory Explains Our Everyday World - Michael D. Fayer

Absolutely Small: How Quantum Theory Explains Our Everyday World - Michael D. Fayer (2010)

Chapter 15. Beer and Soap

IN THIS CHAPTER, we will look at several types of molecules to see how differences in their nature influence chemical processes. First, alcohols will be introduced. An alcohol is an organic molecule that contains a particular type of chemical group. An alcohol can be a relatively small molecule, such as ethyl alcohol, which chemists usually call ethanol. Ethanol is the alcohol in beer, wine, and vodka. However, large important biological molecules, such as cholesterol, are also alcohols. We will get to such large molecules in Chapter 16. We will see why ethanol dissolves in water, how it can turn into vinegar, and outline the chemical reactions in your body that makes methanol (wood alcohol) poisonous, but ethanol safe, at least in moderation. Building on the mechanism that enables some molecules to dissolve in water, we will take a look at the structure of soap and oil molecules to see why you need soap to take grease off of your dishes and get it to wash down the drain.


Ethanol is ethane (Figure 14.10) with one of the hydrogens replaced by an OH group. The OH group is called a hydroxyl group. The chemical formula for ethanol is H3CH2COH. Figure 15.1 shows a diagram and a ball-and-stick model of ethanol. In ethanol as in ethane, the carbon atoms use four sp3 hybrid orbitals to form tetrahedral bonds. The oxygen also uses four sp3 hybrids. One of them is used to make the bond to the carbon, one is used to bond to the hydrogen, and the other two contain electron lone pairs. The lone pairs are not shown in the ball-and-stick model in Figure 15.1. (Figure 14.2 shows the oxygen lone pairs for the water molecule.)

Note that in the ball-and-stick model of ethanol, the hydrogen that is bonded to the oxygen is considerably smaller than the hydrogens bonded to the carbons. Going back to the Periodic Table (Chapter 11), we know that oxygen really wants to pick up electrons to obtain the neon closed shell configuration. It gains the electrons by making covalent electron pair sharing bonds. However, the sharing is not completely equal in a bond between oxygen and hydrogen. The oxygen has a very strong attraction for electrons, so it pulls some extra electron density away from the hydrogen. The extra electron density on the oxygen makes it a little bit negative and the loss of electron density makes the hydrogen a little bit positive. This loss of electron density reduces the size of the electron cloud of hydrogen, which is represented by the reduced size of sphere that represents the hydrogen bonded to the oxygen. Carbon and hydrogen have almost the identical attraction for electrons when they make a covalent bond. They share the electrons almost equally. Thus on average, there is more electron density on a hydrogen bound to a carbon than on a hydrogen bound to an oxygen. In general, an alcohol is a molecule that has a carbon with an OH group and that carbon is also bonded only to hydrogens or other carbons.


FIGURE 15.1. Ethanol (ethyl alcohol) diagram showing atom connectivity (top), ball-and-stick model (bottom). The hydrogens are light gray, the carbons are dark gray, and the oxygen is black.


At room temperature, ethane is a gas but ethanol is a liquid. You have to make ethane colder than—89° C (—128° F) before it becomes a liquid, and you have to make ethanol hotter than 78° C (172° F) before it boils and goes from a liquid to a gas. Ethane and larger hydrocarbons like oil do not dissolve in water, whereas ethanol and larger alcohols do dissolve in water. Ethane and ethanol are almost the same size, and they have similar shapes. So why, in contrast to ethane, does ethanol dissolve in water, and why is it a liquid at room temperature?

As we briefly discussed, ethanol’s hydroxyl group, the OH, has two partial changes. The O is a little bit negative, and the H is a little bit positive. A diagram showing the partial charges is: Oδ-—Hδ+. The δ (Greek letter delta) is used to mean “a little bit.” The δ is followed by the sign of the electrical charge on the atom. The amount of electron density transferred from the H to the O is very small, much less than one full electron that is transferred in a salt like NaCl, which is Na+ and Cl-. The bond between the oxygen and the hydrogen is mainly covalent, not ionic as in NaCl. However, the partial charges on the O and H are unbelievably important. They result from the details of the quantum mechanical molecular orbitals responsible for the oxygen-hydrogen covalent bond. These partial charges result in ethanol being a liquid. If you will permit only mild hyperbole, without the same type of partial charges on the oxygen and hydrogen atoms of water molecules, life would not exist.

Ethanol is a liquid because the partial changes give rise to a type of chemical interaction between molecules called hydrogen bonds. Hydrogen bonds are much weaker than a real covalent chemical bond. They are about a 10th the strength or less. While an accurate description of hydrogen bond formation requires quantum theory, a qualitative understanding can be obtained by considering electrostatic interactions between partial charges. A hydrogen bond is formed when the slightly positive H on one molecule is attracted to the slightly negative O on a different molecule. That attraction will tend to maintain the H of one ethanol in a pretty well-defined position relative to the O of another ethanol. These attractions hold the ethanols together to form a liquid at room temperature. Ethane molecules do not have these relatively strong intermolecular attractions.

Heat is a form of kinetic energy. Heat jiggles the molecules around. In ethane, the molecules do not have a substantial attraction of one for another. At room temperature, the heat-induced motions cause ethane molecules to fly apart, so ethane is a gas. Imagine you and another person hold hands and run in opposite directions. If you hold hands weakly, you will fly apart and take off in different directions, like ethane molecules. If you hold hands very tightly, the two of you will stay together, but move around somewhat as you tug each other to and fro, like ethanol molecules.

Figure 15.2 shows four ethanol molecules hydrogen bonded together into a chain. The dashed lines go from the hydrogen of the OH group of one ethanol to a lone pair on an oxygen of another ethanol. The lone pair has a good deal of electron density, so the slightly positive H is attracted to the electrons of a lone pair on oxygen. This continues from one ethanol to another to form chains. Ethanol liquid is composed of hydrogen-bonded chains that hold the molecules together. The hydrogen bonds make ethanol a liquid at room temperature, but they are relatively weak. These bonds are constantly being broken and reformed, but on average, each ethanol has hydrogen bonds (H-bonds) to one or more other ethanols. However, if you heat the ethanol up enough, the thermally (heat) induced jiggling causes the molecules to break the H-bonds and fly apart. The temperature at which the thermal energy is enough to cause the ethanols to fly apart is the boiling point (78° C). At this temperature or above, ethanol is a gas.


FIGURE 15.2. Four ethanol molecules that form a chain. The oxygens are almost black in the figure. An oxygen has two lone pairs in addition to the hydrogen and carbon bonded to it. The dashed lines show hydrogen bonds that go from the hydroxyl’s H on one ethanol to an oxygen lone pair on another ethanol.


Back to why hydrogen bonds are necessary for life. Water, H2O, is a very small molecule. It has a molecular weight that is similar to that of oxygen (O2), nitrogen (N2), or methane (CH4), all of which are gases at room temperature. Water has a single oxygen bonded to two hydrogens. Like the situation in ethanol, the oxygen makes covalent bonds to the hydrogens, but in an O-H covalent bond, the electrons are not shared perfectly equally. In a water molecule, the oxygen withdraws some electron density from the H’s. A diagram of water showing this is: Hδ+—Oδ-—Hδ+. The slightly positive hydrogens on one water molecule are attracted to the slightly negative oxygens on another water molecule. One water molecule can make up to four hydrogen bonds.

A schematic illustration of water hydrogen bonding is shown in Figure 15.3. The central water molecule is making four hydrogen bonds with the surrounding four waters. The central water’s two hydroxyls are hydrogen bonded to two oxygens of other water molecules. And two other water molecules’ hydroxyls are forming hydrogen bonds with the oxygen of the central water. In contrast to the depiction in Figure 15.3, the hydrogen bonding does not stop with the five water molecules. The outer four water molecules will each make approximately four hydrogen bonds with other water molecules. The result is a hydrogen bonding network.


FIGURE 15.3. A central water molecule hydrogen bonded to four surrounding water molecules. The central water’s two hydroxyl hydrogens bond to two oxygens, and the central water’s oxygen accepts two hydroxyl bonds from two other water molecules.

At room temperature, there is enough heat so that hydrogen bonds between water molecules are continually breaking and new hydrogen bonds are formed with different water molecules. So the hydrogen bond network is not static. It is continually reforming and rearranging. The time scale for these hydrogen bond rearrangements has been measured with ultrafast infrared spectroscopy, and it is approximately 3 ps (picosecond, a trillionth of a second, 10-12 s).

Life is based on chemistry that occurs in water. Recent spacecraft sent to Mars are not looking directly for signs of past life, but rather for signs of past liquid water. Liquid water is so fundamental to life that its presence is a necessary and perhaps a sufficient condition for life to exist. Water’s amazing properties, which are essential to the biochemistry of life, are a result of its hydrogen bonding network structure and the ability of that structure to rearrange. Water’s properties allow it to accommodate a vast array of chemical processes necessary for life. For example, proteins fold in water. Proteins are the very large, extremely complex molecules that are responsible for most of the chemical processes in our bodies. When proteins are chemically produced by other proteins, initially they are not in the correct configuration to function. This is called unfolded. Proteins have regions that will readily form hydrogen bonds to water and regions that are more like hydrocarbon and do not want to mix with water. The protein rearranges its structure, folds, such that the hydrophilic (likes water) regions are on the outside exposed to water and form hydrogen bonds with water and the hydrophobic (dislikes water) regions are on the inside away from water. Such selective interactions with water are an important driving force that helps proteins to obtain the correct shapes necessary to function. It is because water can readily reform its network structure through the making and breaking of hydrogen bonds that it can accommodate the proteins’ structural transformation and a vast number of other chemical processes that occur in living organisms.


One of water’s properties is its ability to dissolve a very wide variety of chemical compounds. We have discussed that NaCl will dissolve in water to give ions, Na+ and Cl-. The positive ions are surrounded by the slightly negative oxygens of water, and the negative ions are surrounded by the slightly positive hydrogens of water. Salt dissolves because of water’s ability to interact favorably with both cations and anions. Water can also dissolve a very wide variety of organic compounds. Water will not dissolve hydrocarbons such as ethane, but it will dissolve organic molecules like ethanol that have hydroxyl groups (—OH) or other groups that have slightly charged or fully charged portions. Water dissolves ethanol by forming hydrogen bonds to the hydroxyl group of ethanol. In pure ethanol, ethanol molecules hydrogen bond one to another to form chains, as shown in Figure 15.2. When ethanol is put in water, water can form hydrogen bonds to ethanol’s hydroxyl, incorporating ethanol molecules into water’s hydrogen bonding network. Vodka is essentially ethanol in water. Wine is water with less ethanol in it than vodka. Wine also has large organic molecules that give red wine its color and all wines their distinct aroma and taste.


If wine is exposed to air for too long, it will “go bad” and turn into vinegar. Vinegar is made by purposely making wine go bad. The chemical reactions that convert wine into vinegar are actually facilitated by a bacteria, acetobacter, which has the ability to convert ethanol into acetic acid when oxygen is present. The process requires two chemical reaction steps. The chemical reactions are written as



First ethanol (CH3CH2OH) is converted to acetaldehyde (CH3CHO) and hydrogen gas (top line), and then two molecules of acetaldehyde and one oxygen molecule (two oxygen atoms) are converted into two molecules of acetic acid (CH3COOH), which is vinegar. The structure of ethanol is shown in Figure 15.1. The structures of acetaldehyde (top) and acetic acid (bottom) are shown in Figure 15.4. In both acetaldehyde and acetic acid, the carbon labeled C1 forms a methyl group. C1 is bonded to three hydrogens and carbon C2. Acetaldehyde has C2 also bonded to a single hydrogen and double bonded to an oxygen. In general, an aldehyde has a carbon double bonded to an oxygen, bonded to a hydrogen, and bonded to another carbon. In formaldehyde (Figure 14.3), instead of bonding to another carbon, C2 is bonded to a second hydrogen. C2 uses three sp2 hybrid orbitals to form three σ bonds and an additional 2p orbital to combine with a 2p orbital on the oxygen to formaπbond to make the double bond. As shown in the top line of the chemical reaction equations and looking at the structure of ethanol in Figure 15.1, ethanol goes to acetaldehyde by eliminating two hydrogen atoms to give acetaldehyde and an H2 molecule. Two molecules of acetaldehyde each pick up one oxygen atom from O2 to yield two acetic acid molecules (bottom of Figure 15.4). Acetic acid has C2 bonded to two oxygen atoms. C2 is double bonded to one and single bonded to the oxygen of the hydroxyl group—OH.


FIGURE 15.4. Acetaldehyde (top) and acetic acid (bottom). Oxygens are the almost black spheres. Acetaldehyde C2carbon is bonded to C1, a hydrogen, and double bonded to an oxygen. Acetic acid C2is bonded to C1, double bonded to an oxygen and single bonded to another oxygen that is part of a hydroxyl group.

Acetic acid is an organic acid, also called a carboxylic acid. The -COOH is the acid group. The simplest definition of an acid is a chemical that when put in water results in the generation of hydrogen ions, H+. Hydrochloric acid is HCl. Like NaCl, when put in water HCl comes apart to give H+ and Cl- ions. HCl dissociates into ions completely. It is called a strong acid because for every molecule of HCl put in water, you get one H+ ion. The H+ will be associated with (solvated by) water molecules. The slightly negative oxygens of water molecules surround the H+. The acid dissociation reaction of acetic acid in water can be written in the following manner.


The -COOH group is the organic acid group. This group ionizes to give—COO- + H+ as shown. The diagram indicates that the negative charge is stuck on one of the oxygens. In fact, it is shared equally by both oxygens. This equal sharing is represented by the following diagram.


The dashed curve indicates that the molecular orbital that contains the electron that gives rise to the negative charge is spread out over both oxygens. Each oxygen can be thought of as having one-half of a negative charge. Organic acids are very soluble in water. Ethanol is soluble because the hydroxyl oxygen has a small negative charge that leads to hydrogen bond formation with water molecules. Nondissociated acetic acid has two oxygens with partial charges that can hydrogen bond to water. Dissociated acetic acid has two oxygens, each with a half negative charge that readily hydrogen bonds to water.

Organic acids, such as acidic acid, are weak acids. In water, only a small fraction of the acetic acid molecules ionize to produce H+ and acetate anion. If you have 60 grams of acidic acid in a liter of water, which is about four tablespoons (2 ounces) in a quart of water, only about 0.4% of the acidic acid molecules will be ionized to produce H+ and acetate anion. This concentration, four tablespoons of acetic acid in a quart of water, is approximately the concentration of acetic acid in common vinegar. It is the acetic acid that gives vinegar its snappy taste. Organic acids are very common in chemistry and biology. All proteins are composed of combinations of 20 amino acids. Prior to reactions to form a protein, each amino acid is a type of organic acid containing the organic acid group —COOH.


Methanol is the smallest alcohol molecule. Ethanol is ethane with one hydrogen replaced by a hydroxyl group,—OH. Methanol is methane with one hydrogen replaced by a hydroxyl. While ethanol can be ingested in reasonable quantities without dire consequences, methanol is highly toxic. Methanol, also called wood alcohol, is a common contaminant of moonshine liquor. As little as 20 ml (milliliters) has been known to cause death, and 15 ml can cause blindness. Fifteen ml is one tablespoon. Two tablespoons are about one ounce of liquid. Eight ounces of wine contain about one ounce of ethanol. So replacing the ethanol with methanol in a class of wine can cause blindness and death. This is rather remarkable since ethanol is just methanol with one additional methyl group (—CH3).

It is not methanol itself that is poisonous, but rather the metabolic products of methanol. In humans and other living organisms, alcohols are converted to other chemicals by enzymes (proteins that are responsible for chemical reactions) called alcohol dehydrogenases. In humans, these enzymes are found in the liver and the lining of the stomach. The evolutionary development of these enzymes probably occurred to break down alcohols that are produced by bacteria in the digestive tract or to break down alcohols that are natural components of some foods. Ethanol is converted first to acetaldehyde and then to acetic acid as discussed in connection with Figure 15.4 above, which shows the structures of these two molecules. Acetaldehyde and acetic acid are not harmful to the body and can be readily eliminated. Larger alcohol molecules are also converted first to aldehydes and then to organic acids, which can be eliminated from the body without harm. However, alcohol dehydrogenases convert methanol to formaldehyde and then to formic acid. The structure of formaldehyde is shown in Figure 14.3. Formaldehyde is like acetaldehyde except carbon C2 of acetaldehyde (Figure 15.4, top) is bonded to a single hydrogen rather than to a methyl group (C1 in Figure 15.4). Formic acid is like acetic acid (Figure 15.4, bottom) but again with C2 bonded to a hydrogen rather than a methyl group. Formaldehyde in particular, but also formic acid, are highly toxic. They damage the retina and the optic nerve, leading to vision impairment and blindness. However, formic acid also is responsible for serious acidosis, which involves interfering with enzymes that break down carbohydrates. Relatively low concentrations of formaldehyde and formic acid, as well as other metabolites derived from them, can cause death.


As we discussed, alcohols like ethanol and organic acids, such as acetic acid, are very soluble in water because the organic groups containing oxygens can form hydrogen bonds with water. In contrast, ethane, which is very similar to ethanol, does not dissolve in water, because it does not have oxygen containing groups that can form hydrogen bonds. Ethane is a hydrocarbon, that is, it is composed only of hydrogens and carbons. Methane and ethane are gases. Larger hydrocarbons, beginning with pentane (five carbons), are liquids at room temperature. The smaller of these liquid hydrocarbons like pentane and octane (a component of gasoline) are very thin liquids, that is, they are not viscous.


As the number of carbons increases, the liquid hydrocarbons become increasingly viscous. The heating oil, used in homes and businesses in many areas of the United States, is composed of a mixture of hydrocarbons typically spanning the range of 14 to 20 carbons. At room temperature, oil flows readily but it is much more viscous than gasoline. Grease is composed of really large hydrocarbons. It is very viscous and will not pour at room temperature.

The hydrocarbons that comprise heating oil are liquids at room temperature, but as discussed, they do not dissolve in water. Molecules with 14 carbons are the lightest component of heating oil. Figure 15.5 shows n-tetradecane. Decane has 10 carbons. Tetradecane has four (tetra) additional carbons. The n (normal) means that all of the carbons are connected one to another with no branches. That is, each carbon is connected to at most two other carbons. The upper portion of the figure shows a ball-and-stick model of n-tetradecane. It is important to remember that the electron density surrounding the atoms in a molecule is space filling. The bottom portion of the figure shows a space-filling model of n-tetradecane.


FIGURE 15.5. n-tetradecane, C14H30, ball-and-stick model (top) and space-filling model (bottom). The molecule has 14 carbons connected one to the next without branching.


Many other hydrocarbons have 14 carbons. These are branched. Figure 15.6 shows a ball-and-stick model (top) and a space-filling model (bottom) of one of them, 2,8-dimethyldodecane. Dodecane has 12 carbons. Two additional methyl groups branch off of the main chain at the second and eighth carbon from the left. n-tetradecane and 2,8-dimethyldodecane are structural isomers. They have the same number of hydrogens and carbons, but no amount of rotation about the bonds can convert one into the other. Both n-tetradecane and 2,8-dimethyldodecane have many conformers, that is, it is possible to rotate about various carbon-carbon single bonds to produce different shapes without changing how the carbons are connected. Structural isomers and conformers were discussed in connection with butane (see Figures 14.12 and 14.13).


FIGURE 15.6. 2,8-dimethlydodecane, C14H30, ball-and-stick model (top) and space-filling model (bottom). The molecule has 14 carbons. There is a chain of 12 carbons, with two methyl groups branching off at the second and eighth carbons from the left.


Heating oil is a relatively viscous liquid, although the hydrocarbon molecules have relatively weak attractive interactions one for another. The large number of sizes, structural isomers, and conformers cause the molecules to become entangled, which contributes to the viscosity. If oil is in water, it will float on top. If you shake it up, it will appear to mix for a while. However, if you let it stand, the oil will separate and again float on top. Anyone who has made their own oil-and-vinegar salad dressing knows this. You mix olive oil, vinegar, and possibly some water, and then you shake it up. If you let it stand, the olive oil floats right back to the top. In commercial oil and vinegar salad dressings, emulsifiers are added to keep the oil and vinegar from separating. Emulsifiers are very similar to the soap that we are about to describe. We know that an oxygen atom in a water molecule is slightly negative and is attracted to atoms that are positively charged or at least have a partial positive charge. The hydrogens of water molecules are slightly positive and are attracted to negatively charged or partially negatively charged atoms. Hydrocarbons have carbons and hydrogens that are essentially neutral in charge. Therefore, water molecules are attracted to each other much more strongly than they are attracted to oil. The result is oil does not dissolve in water.


Soap makes oil dissolve in water. Many different molecules are used as soaps or detergents. The more formal name for a soap molecule is a surfactant. While the chemical nature and structure of surfactants vary widely, all surfactants have a common feature. A section of a surfactant molecule, if taken by itself, would be very soluble in water. The other part of the surfactant molecule by itself would be very soluble in oil and grease.

One such molecule is sodium n-heptadecaneacetate, which is shown as both a ball-and-stick model and a space-filling model in Figure 15.7. n-heptadecane is an unbranched 17-carbon chain. This hydrocarbon portion of the molecule is shown in the figure as a particular conformer with a couple of rotations about carbon-carbon bonds that produces the bent shape. Tetradecane, shown in Figure 15.5, is all trans. It does not have any rotations to give some gauche conformations. Large hydrocarbons have many different conformers that can interconvert. By itself, heptadecane would be one component of heating oil.

The n-heptadecane hydrocarbon is attached to an acetate group or acetate anion. The acetate group comprises the last two carbons and two oxygens on the right side of the molecule shown in Figure 15.7. The acetate anion is shown on page 260 in the chemical diagram representing the dissociation of acetic acid. For dissociated acetic acid, the cation is H+. Here the cation is sodium, Na+, which is not shown in Figure 15.7. Sodium acetate is represented in the following diagram.


Sodium acetate is a sodium salt like sodium chloride, NaCl. However, here the anion is an organic anion rather than the elemental anion, Cl-. Sodium acetate dissolves completely in water just like NaCl does.


FIGURE 15.7. Sodium heptadecaneacetate, C18H37COO-Na+, ball-and-stick model (top) and space-filling model (bottom). The dissociated sodium ion is not shown. The molecule has 19 carbons. There is a chain of 17 carbons and then an acetate group. The δ-indicates that each oxygen (darkest spheres) carries an approximately negative one-half charge.

Putting Soap in Water Forms Micelles

So the molecule sodium n-heptadecaneacetate is composed of a long hydrocarbon chain that will not dissolve in water and sodium acetate that readily dissolves in water. What happens if you put a substantial amount of soap, in this case sodium n-heptadecaneacetate, in water with no oil or grease around? The hydrocarbon portions of the molecules hate water, so they want to avoid it. As a pure hydrocarbon, n-heptadecane would completely separate from the water and float on top of it. However, the sodium acetate portion loves water. It will dissociate into an acetate anion and a sodium cation, both of which will have strong favorable interactions with water molecules. Both portions of the surfactant molecule get what they want by forming micelles. Micelles are nanoscopic structures, that is, they are structures with a size on the order of nanometers. A common shape for a micelle is spherical, or near spherical, although a variety of shapes occur, depending on the surfactant and the concentration of surfactant in the water. Their sizes are typically about 10 nanometers, that is, 10 billionths of a meter in diameter. The size and structure of the surfactant molecules determine the size of the micelles.

Figure 15.8 shows a schematic illustration of a spherical micelle. The balls represent the acetate or other charged or hydrophilic portions of the surfactant molecules. The hydrophilic portion of a surfactant molecule is often referred to as the head group. The squiggly lines represent the hydrophobic hydrocarbon tails of the surfactants. The head groups are very soluble in water and form an outer shell. The hydrocarbon tails avoid contact with water by clustering together to form a nanodroplet of oil called the core of the micelle. The formation of micelles enables soap to readily dissolve in water.


FIGURE 15.8. A schematic of a spherical micelle. The balls represent the acetate anion head groups. Wiggly lines represent the hydrocarbon tails. The micelle is surrounded by water that hydrogen bonds to the head groups. The hydrocarbon chains clump together to form a nanodroplet of oil, which is protected from water by the head groups.

Soap Dissolves Grease

Now consider what happens when plates or hands, with grease or oil on the surface, are put into soapy water. In pure water, the hydrocarbons on a surface are repealed by the water. However, with soap micelles in the water, the situation is very different. The charged head groups of the micelles come in contact with the oily surface. The head groups want to avoid the oil, which causes the micelles to open up, exposing the surfactants’ hydrocarbon tails to the grease. The tails of surfactants are perfectly happy to be embedded in the oil and grease. The oily hydrocarbons become entangled with the surfactant tails. Helped by agitation, some of the oil hydrocarbons lift off from the rest of the oily surface. The surfactant head groups close up around the core, reforming a micelle. However, some of the oil and grease hydrocarbons have been incorporated into the micelle’s core. The containment of hydrocarbons in the interior of a micelle is shown schematically in Figure 15.9. The hydrocarbon tails of the surfactants are the double lines, while the oil hydrocarbons are single spotted lines. The oil and grease molecules remain in the micelle core as part of the oil nanodroplet. The additional hydrocarbons in the core make the micelles bigger. More surfactant molecules, which are in the water, can join a micelle to fully surround the enlarged oil nanodroplet. The charged head groups of one micelle repel those of other micelles, which prevents the grease from coagulating and forming grease globs that are not soluble in water.

Soap-like materials are reported to have been produced as early as 2800 BCE. True soaps, basically the same as those used today, were made by chemists in the Islamic world in the seventh century. Today, we hear a lot about the coming of nanotechnology, in which nanometer scale assemblies of molecules or atoms can perform very specialized functions. It is remarkable that soap in water is a nanoscopic material. Surfactants form nanometer-size micelles that can encapsulate grease and oil. The micelles with encapsulated hydrocarbons are soluble in water, which makes it possible for us to wash away otherwise water-insoluble molecules.


FIGURE 15.9. A schematic of hydrocarbons from oil or grease (single spotted lines) contained in the interior of a soap micelle.