Organic Acids and Bases

Organic acids and organic bases are acids and bases that contain a carbon skeleton. Within these categories are a number of classes of neutral proton acids and bases (that is, uncharged acids and bases.) The first part of this section examines the three main types of

neutral organic proton acids to see why they are acids and why they have widely different acid strengths. The second part looks at the two main types of neutral organic bases. The last part looks at positively charged carbon acids and negatively charged carbon bases.

Three main types of neutral organic Brønsted-Lowry acids are carboxylic acids, phenols, and alcohols. Each of these three functional groups has an —OH group. Each is acidic because of the electronegativity difference between the oxygen and the hydrogen involved in the O—H bond. The differences in acid strength of the three functional groups are due to the differences in stability of the conjugate base. The most acidic of the three groups are the carboxylic acids. Carboxylic acids are characterized by the presence of the carboxyl group:

 

Carboxylic acids are among the most acidic of the neutral organic acids, but they are rather weak acids. For example, the pKa of acetic acid, a common carboxylic acid, is 4.8, indicating that only a small portion of the molecules of acetic acid ionize in an aqueous solution. In contrast, mineral acids, such as HCl, with a pKa of –7.0, and HNO3, with a pKa of –5.2, completely ionize in aqueous solutions. Although carboxylic acids are weaker than mineral acids, they are the strongest of the neutral organic acids that you will study.

The reason for the relative strength of the carboxylic acids is the conjugate base is resonance-stabilized, which makes it a weak base.

 

 In the carboxylate ion the negative charge spreads over the two oxygen atoms as a resonance hybrid. This reduces the energy of the anion and makes the carboxylic acid more acidic.
Another way of visualizing the reason for the acid strength of carboxylic acids is to look at the molecular orbital system of the carboxylate ion. The carboxylate ion includes three p orbitals that contain a total of four electrons. The overlap of these three p orbitals results in a three-centered π molecular orbital system.

The carbon is joined to each oxygen atom by the equivalent of ½ of a π bond. Each oxygen atom bears ½ of the negative charge.

 

The second main type of neutral organic acids are the phenols. Phenols are much less acidic than carboxylic acids. An —OH group attached to an aromatic ring is characteristic of phenols:

 

Phenol has a pKa of 10.0 in aqueous media, indicating that in water only a very small portion of it ionizes.

 

Phenols are moderately strong organic acids because their conjugate bases are resonance-stabilized. The aromatic ring is involved in resonance, which stabilizes the negative charge.

 

However, this stabilization is less significant than it is for carboxylic acids for two reasons: the resonance stabilization of the phenolate ion disrupts the aromaticity of the aromatic ring, and the resonance stabilization places a negative charge on the carbon atoms, which, when compared to oxygen, are not very electronegative.
The third type of neutral organic acids are alcohols. An —OH group attached to an alkyl group characterizes an alcohol:

 R---OH

Alcohols are much less acidic than phenols. In fact, most alcohols have an acid strength slightly lower than that of water.

 

 

A typical alcohol has a pKa of 15 to 18 in aqueous media, indicating only a very small amount of ionization. Alcohols have such a low acidity because there is no resonance stabilization of the conjugate base.
This section discusses only two of the many types of neutral organic bases: amines and ethers. The primary characteristic of neutral organic bases is they contain one or more pairs of nonbonding electrons. These pairs of electrons are available to donate to a Lewis acid or to accept a proton when the base is acting as a Brønsted-Lowry base. The more available the pair of electrons, often called a lone pair, the stronger the base. Any molecule with a lone pair of electrons can act as a base.
The most common of the organic bases are the amines. Amines are derivatives of ammonia (NH3) and most are weak bases in aqueous media.

The pKa of methyl ammonium ion is 10.6 meaning that the methyl ammonium ion is a relatively weak acid. Thus, methylamine is a moderately strong base.
Amines are stronger bases than other neutral organic bases because the nonbonding pair of electrons on the nitrogen is more available than nonbonding pairs of electrons on other neutral organic bases. The atoms that are found in these other neutral organic bases are oxygen, sulfur, or the halogens. Nitrogen holds its electrons less tightly than these other atoms, so its compounds are the stronger bases. Figure 5.3 illustrates the structure of an amine.

Figure 5.3. Structure of the amine nitrogen.
Ethers, the second type of neutral organic bases, have the general structure ROR′. Ethers are weak bases in aqueous media. In fact, they are so weak that they do not appreciably protonate, or accept a proton, even in 1 M HCl. The pKa of the conjugate acid of ethyl ether is –3.8. A pKa of this magnitude indicates that water is a better base than is an ether.
In nonaqueous media, ethers are good Lewis bases, forming stable complexes with Lewis acids. The ability to form stable complexes is extremely important in organic reactions. For example, in organic synthesis, chemists widely use the complex of BH3 with the cyclic ether tetrahydrofuran:

The third category of organic acids and bases discussed in this section are the positively charged acids and the negatively charged bases. Positively charged acids are electron-deficient. That is, they are organic acids that contain a carbon without an octet of electrons. The most significant electron-deficient organic acid is the carbocation (formerly called a carbonium ion). Carbocations are very reactive reaction intermediates, so chemists seldom observe them directly. A carbocation is a Lewis acid because, without a full octet of electrons, it is electron-deficient and "needs" electrons. As a result of this need for electrons, it reacts with the first available Lewis base—although it prefers a hard one because it is a hard acid. As Figure 5.4 shows, the positively charged carbon forms three sp2 hybridized bonds in a plane with an empty p orbital perpendicular to that plane. Chapter 12 examines nucleophilic substitution reactions that involve carbocations.

The negatively charged organic base discussed in this section is the carbanion. A carbanion has bonds to three other atoms and one pair of nonbonding electrons. The structure of a carbanion is much like the structure of an amine (See Figure 5.5). Because carbon is not very electronegative, it holds these nonbonding electrons loosely. Thus, a carbanion is a strong base. (Chapters 19 and 20 cover carbanion reactions extensively.)

Now that you have seen the various types of organic acids and bases, Section 5.5 examines the factors that modify the strength of the specific acids and bases.

Polarizability

Polarizability means the ability of an atom to have a distorted distribution of electrons.

Charge density

Charge density is the volume of space occupied by a charge. A large ion has a lower charge density than a small ion does.

Hardness or softness

Hardness or softness is a qualitative measure of the reactivity of acids and bases. Hard or soft is independent of strength or weakness of acids and bases.

Hard and Soft Acids and Bases

The ease with which an acid-base reaction occurs depends on the strength of both the acid and the base. Strong acids and bases are generally more reactive than weak acids and bases. However, the direction of the reaction and the stability of the products often depend on another quality—the hardness or softness of the acid and base. Although chemists have not created a quantitative measure to describe the qualities that makes an acid or base hard or soft, they do describe them qualitatively. As you look at the following list of characteristics that describe hard and soft acids and bases, remember that an acid has an empty orbital and an unfilled valence shell, and a base has in its valence shell a pair of nonbonding electrons that is available for donation.

Soft Acids. For soft acids, the electron-pair acceptor atoms are large, have a low positive charge density, and contain unshared pairs of electrons in their valence shells. The unshared pairs of electrons are in the p or d orbitals. Also, soft acids have a high polarizability and a low electronegativity. In organic chemistry, the soft acids usually include only the halogens, phosphorus, and sulfur compounds.

Hard Acids. For hard acids, the acceptor atoms are small, have a high positive charge density, and contain no unshared pairs of electrons in their valence shells. They have a low polarizability and a high electronegativity. The hydrogen ion is a good example of a hard acid.

Soft Bases. For soft bases, the donor atoms hold their valence electrons loosely. They have high polarizability, low negative charge density, and low electronegativity. Common soft bases are the cyanide (-CN) and iodide (I-) ions.
Hard Bases. For hard bases, the donor atoms are small, have a high negative charge density, and hold their valence electrons tightly. They have a low polarizability and a high electronegativity. The hydroxide ion is a good example of a hard base.
To visualize a polarizable atom, imagine that an atom is a large floppy ball and you are holding it cupped in both hands. The ball tends to be spherical, but, as you shift one hand higher than the other, it easily deforms. If you raise your left hand a little, the portion of the ball in your right hand becomes larger. Then, if you raise your right hand a little, the portion of the ball in your left hand becomes larger. A polarizable atom shifts its electron density from one part of the atom to another: at one instant, one portion of the atom has the higher electron density; then the next instant, another portion has the higher electron density.

For the concepts of hardness or softness of acids and bases to be of value to you, you must be able to differentiate between them. To do this, your most useful tool is the periodic table. A general rule is that hardness goes to softness moving from the top to the bottom on the periodic table because the size of the atoms increases with increasing numbers of electrons. A larger acid or base has a lower charge density and is more polarizable. For example, base softness in Group VII A on the periodic table decreases in this order: I- > Br- > Cl- > F-. Also, the elements on the left side tend to be acids, and elements on the right side tend to be bases. In this way, chemists approximately rank acids and bases in order of hardness or softness. Base softness within a period on the periodic table decreases in order of increasing electronegativity; for example, -CH3 > -NH2 > -OH > F.
Hardness and softness are difficult to quantify. Rather than relying specifically on these types of sequences, chemists divide acids and bases into three groups: (1) hard acids or bases, (2) soft acids or

bases, and (3) borderline acids or bases. Table 5.2 lists a few examples of each category.

H⊕ is a hard acid because it has no electrons and has a high positive charge density. The HO- ion is a soft base because it has a pair of electrons and only one proton, so it holds the electrons rather loosely. Thus, it is quite polarizable and soft.

Acid and Base Strength

The strength of a Brønsted-Lowry acid or base depends on the extent to which it ionizes in water. Although there are numerous solvents besides water, chemists discuss acid and base strength in relation to water because they use it so widely as a solvent. Chemists use the autoionization of pure water to determine the values for the concentrations of acidic and basic solutions. Autoionization is the reaction of two molecules of water with each other to give a hydronium ion, H3O⊕, and a hydroxide ion, -OH.

For this reaction, the amount of autoionization is extremely slight—at 25oC, it is 10–7 M (moles/liter). The concentrations of H3O⊕ and -OH are equal; that is, both measure 10–7 M. Chemists call this a neutral solution. If you add a compound that is more acidic than water, you increase the concentration of H3O⊕ ions and make the solution acidic. If you add a compound that is more basic than water, you increase the concentration of -OH ions and make the solution basic.
The product of the H3O⊕ and -OH concentrations in water is equal to 10–14 and is a constant, Kw. Chemists define Kw with the following equation.

Because the concentrations of H3O⊕ and -OH are equal in a neutral solution, you can easily calculate the concentration of both:

Because the product of the two concentrations is a constant, Kw, when one concentration increases, the other must decrease. For example, if you add -OH ions to water the concentration of the H3O⊕ decreases by whatever amount is necessary for the product of the two concentrations to still equal 10–14.
Because the hydronium ion concentrations can span a very wide range of values, from greater than 1 M down to less than 10–14 M, chemists measure the concentration of H3O⊕ on a logarithmic scale called pH. The pH values give the hydronium ion concentration of a solution. Therefore, measuring the pH of a solution is a means of quantifying the acidity of that solution. Chemists define this

measurement as the negative logarithm (base 10) of the H3O⊕ concentration, represented by the following equation:

pH = –log10[H3O⊕]

For simplicity, this book will normally refer to the H3O⊕ ion as the H⊕ ion from now on. If an equation shows the H⊕ ion present in aqueous solution, remember that it is actually the H3O⊕ ion.
This equation shows the general reaction of an acid in water:

Note that this reaction is an equilibrium. Most acid-base reactions are equilibrium reactions because the reactants only partly ionize. Strong acids and bases ionize completely in water. Weak acids and bases ionize only partly in water. An acidic, aqueous solution is any solution with a concentration of hydrogen ions greater than 10–7 M. Similarly, a basic solution is any solution with a concentration of hydroxide ions greater than 10–7 M.
To determine the relative strength of an acid or a base, you need to find out how much the acid or the base ionizes, or dissociates, in water at equilibrium. The equilibrium constant, Ke, gives this information and is defined as follows:

However, because water is the solvent and its concentration is essentially constant, a more meaningful value for acid ionization comes from multiplying the equilibrium constant by the water concentration:

Chemists call Ka the acid dissociation constant. The value of Ka specifies the strength of the acid. The stronger the acid, the larger the amount of dissociation and the larger the concentration of H3O⊕ ions. Thus, the stronger the acid, the larger the value of Ka. Strong acids completely dissociate in water and have large dissociation constants. Most organic compounds are weak acids and have dissociation constants in the range from 10–2 to 10–60.

Because acids have such a large range of values for their dissociation constants, chemists often convert those values to a logarithmic scale, similar to pH. The following equation defines this scale:

pKa = –log10 Ka

Autoionization

Autoionization is a process by which one molecule of a compound reacts with another molecule of the same compound in an acid-base reaction

Acids and Bases

Three major definitions of acids and bases have influenced the thinking of chemists. In 1884, Svante Arrhenius formulated the first of these definitions. Then, in 1923, independently of each other, Johannes N. Brønsted and Thomas M. Lowry developed the second. The third definition grew from Gilbert Newton Lewis's theory of covalent bonding, which he proposed in 1916.
The first definition, proposed by Svante Arrhenius in his doctoral dissertation, was so revolutionary that he was almost denied his Ph.D. However, in 1903, he received the Nobel Prize in chemistry for his theory. His theory states that a stable ionic compound that is soluble in water will break down, or dissociate, into its component ions. This dissociation, or ionization, of a compound in water, leads to Arrhenius' definition of an acid and a base. An acid is a substance that, when added to water, increases the concentration of hydronium ions, H3O⊕. Because Arrhenius regarded acid-base reactions as occurring only in water, he frequently called the hydronium ion a hydrogen ion, H⊕. An H⊕ ion is a proton, or a hydrogen that is electron-deficient. Thus, a base is a substance that, when added to water, increases the concentration of hydroxide ions, -OH. The following statements summarize his definition.

An Arrhenius acid is a source of H⊕ ion.
An Arrhenius base is a source of -OH ion.

The Arrhenius acid-base theory provided a good start toward understanding acid-base chemistry, but it proved much too limited in its scope.
Brønsted and Lowry developed a more general acid-base definition than that of Arrhenius. Although they considered reactions other than those that take place in aqueous solutions, they still said acids were molecules that donate a hydrogen ion—such as HCl and H2SO4. However, they broadened the definition of bases to include any compound that accepts a proton. The basis of their acid-base definition is that in a reaction a proton transfers between reactants. Thus, acids involving a transfer of H⊕ ions are sometimes called proton acids. According to the Brønsted-Lowry definition, an acid is any molecule or ion that donates a proton to another molecule or ion, and a base is any molecule or ion that receives that proton. The following statements briefly summarize the Brønsted-Lowry definition.
A Brønsted-Lowry acid is a proton donor.
A Brønsted-Lowry base is a proton acceptor.
An example of the Brønsted-Lowry definition is the reaction between hydrogen chloride and sodium hydroxide:

 

In this reaction, HCl is the acid because it is the source of protons, or hydrogen ions; NaOH is a base because the hydroxide ion is the proton acceptor. The following reactions further illustrate the Brønsted-Lowry acid-base definition.

When an acid and a base react with each other, the reactants and products are in equilibrium with each other. Note the two-way arrows. They indicate that this is an equilibrium reaction. That is, the reactants on the left side of the equation are reacting and forming product, and the products on the right side are also reacting and forming the starting reactants. Chemists call the acid and base on the right side of the equation the conjugate acid and conjugate base. The reaction below is labeled to show the conjugate acid and conjugate base.

 

A hydrogen of sulfuric acid (H2SO4) is the acid, and the nitrogen of ammonia (NH3) is the base. They react to form the hydrogen sulfate anion (HSO4-) and the ammonium ion (NH4⊕). The ammonium ion is the conjugate acid of ammonia. The bisulfate ion is the conjugate base of the sulfuric acid.

Like Brønsted and Lowry, G. N. Lewis defined acids and bases in a broader scheme than Arrhenius did. Lewis noted that there are a number of reactions that look like acid-base reactions but do not involve the transfer of a proton. Instead, they involve the interaction of a pair of nonbonding electrons. From that observation, he defined an acid as a molecule that forms a covalent bond by accepting a pair of electrons and a base as a molecule that forms a covalent bond by donating a pair of electrons. Below is a simplified statement of the Lewis definition of acids and bases.
A Lewis acid is an electron-pair acceptor.
A Lewis base is an electron-pair donor.

 

Because a Lewis acid accepts a pair of electrons, chemists call it an electrophile, from the Greek meaning "lover of electrons." They call the base a nucleophile, or "lover of nuclei," because it donates the electrons to a nucleus with an empty orbital. In a chemical reaction, a nucleophile seeks a nucleus, or a positive charge, and an electrophile seeks electrons, or a negative charge. Fundamental to organic chemistry is the fact that nearly all the reactions that you will study are reactions of an acid with a base or, more commonly, of an electrophile with a nucleophile.

 

Chemists use a curved arrow () to show electron movement. A curved arrow points from the electron-rich reactant, the base or nucleophile, toward the electron-poor reactant, the acid or electrophile. Rewriting the previous two reactions using a curved arrow shows the movement of electrons. In each reaction, a pair of nonbonding electrons from a nucleophile reacts with an electrophile to form a bond.

 

All chemical reactions involve orbital interactions. The orbital description of a reaction can help you understand how chemical reactions occur. As you study the various reactions presented in this book, think about the orbitals involved in the reactions. Figure 5.1 is a molecular orbital picture of ammonia reacting with boron trifluoride to form a new bond. Ammonia is a base with a pair of nonbonding electrons. The nitrogen of ammonia is sp3 hybridized. Boron trifluoride is an acid with an incomplete octet of electrons. The boron is sp2 hybridized with an empty p orbital. The reaction occurs when an sp3 orbital of ammonia overlaps with the empty p orbital of boron trifluoride. In the process, the boron becomes sp3 hybridized. With this overlap the two molecules form a new bond.

 

Density

For liquids, density is usually measured as the number of grams per milliliter—expressed as g/mL. Density is a useful tool in the laboratory. Its most common use is in extractions, the transfer of a solute from one solvent to a better one. To do an extraction, choose a solvent that is better than the one the solute is currently dissolved in. The two solvents must be insoluble in one another. Then shake the two together and separate them. Because you know the density of each solvent and which solvent best dissolves the solute, you know which one to keep. You may also use density, as you do boiling or melting points, to help identify unknown liquids.

Density of Solids
You can easily estimate the relative density of many organic liquids, but you cannot do so as readily with solids because the density of a solid depends on too many factors. Two generalizations that relate to the density of solids are: organic solids typically have lower densities than inorganic solids, and nonionic solids typically have lower densities than ionic solids. Beyond these generalizations, this book does not consider the density of solids.

The density of an organic liquid depends on three factors: the molecular weight of the substance, the ratio of the number of heavy atoms to the number of carbons in the molecule, and how well the molecules pack together. The first two factors are very closely related because, when determining the density of a liquid, there is more to the picture than just the molecule's actual weight. Of greater importance is the ratio of the number of heavy atoms to carbon atoms in the molecule. For example, the van der Waals radius of a methyl group is

200 pm, and the van der Waals radius of a bromine atom is 195 pm. The formula weight of a methyl group is 15, and the atomic weight of the bromine is 80. Bromine weighs more than five times as much as the methyl group, yet it takes up a smaller volume of space. This leads to a dramatic increase in density for bromine compared to a methyl group. Thus, when you compare the two in compounds pentane and 1-bromobutane, for example, you find that both compounds have nearly the same molecular volume, but the density of pentane is 0.62 g/mL, and that of 1-bromobutane is 1.27 g/mL.

This type of difference holds true for any set of molecules with similar sizes—the molecule with the heavier atoms is the molecule with the higher density. For example, cyclohexane and cyclohexane-d12 (in which all of the hydrogens have been replaced with deuterium, 2H) have densities of 0.78 and 0.89 g/mL respectively.
The third factor that determines a liquid's density is how efficiently its molecules pack together. The closeness of packing

depends on how readily each molecule fits into an aggregation of molecules. For example, hexane, with a density of 0.66 g/mL, has many more stable conformations than does cyclohexane, with a density of 0.78 g/mL. Cyclohexane packs more efficiently because it is more symmetrical than hexane.
Intermolecular attractions affect the packing efficiency of a molecule. The greater these attractions, the more readily the molecules pack together. Dipolar attractions and, especially, hydrogen bonding increase the density of a liquid. For example, compare the isomers diethyl ether and 1-butanol. Diethyl ether, which has van der Waals and dipolar attractions, has a density of 0.71 g/mL. 1-Butanol, which has hydrogen bonding in addition to van der Waals and dipolar attractions has a density of 0.81 g/mL.

Chemists often perform extractions using a two-phase liquid system consisting of water and some organic liquid. They use an estimate of the density of the organic liquid relative to that of water. They work with estimates rather than precise densities because that is all they need to know to decide which layer is which. Here is a rule for estimating the densities of molecules relative to water. Any molecule containing bromine, iodine, or multiple chlorine atoms (e.g., CHCl3) has a density greater than that of water (1 g/mL). All other organic solvents generally have a density less than that of water. This rule, of course, is less effective for large molecules. In a large molecule, the functional group becomes less important, and all the physical properties tend to converge to the same values. Table 4.3 illustrates how chain length affects densities. Note that the density changes rapidly for molecules containing between two and ten carbons, but above ten carbons, the density does not change very much.

Density

Density is the mass of some substance per unit volume of that substance.

emulsion

An emulsion is a suspension of small droplets of one liquid in another liquid in which it is normally insoluble.

Surfactants

When chemists hear the word “surfactant,” they first think of soaps or detergents. Although soaps and detergents are the most familiar of the surfactants, they are not the only ones. There are different kinds of surfactants—many of which have commercial uses. Even your body produces surfactants.

A surfactant is one of a class of chemical structures with a “dual personality” in respect to solubility. Surfactants are usually large molecules. One end is soluble in water; the other end is soluble in the typical organic solvents. This dual solubility is due to a highly polar segment at one end of the molecule and a nonpolar segment at the other end.
Now look at soaps and detergents. A typical soap is the sodium or potassium salt of a long-chain carboxylic acid. Potassium stearate, CH3(CH2)16COOK, is a soap molecule. The “head” of the molecule, the ionic portion, is water-soluble; the “tail” of the molecule, the hydrocarbon chain, is water-insoluble. Chemists call the head the hydrophilic portion of the molecule and the tail the hydrophobic portion.

Adding a soap molecule to water results in the formation of small agglomerations of molecules with long hydrophobic tails. These tails are dissolved in one another and the polar portions pointing outward into the water. These agglomerations are called micelles. Figure 4.8 illustrates the structure of a micelle.

When you wash your skin or clothing, you want to remove two basic types of materials: dry particles and grease. Usually, you can readily rinse dry particles off. Washing grease off is much harder because it is insoluble in water. Here the surfactants in soap help. The tails of the soap molecules are a solvent for the grease. So they dissolve the grease. The heads of the soap molecules project out of the grease particle. Because the heads are soluble in water, they dissolve in the water, and the water carries the soap and the grease away. The interaction of the ionic ends of the surfactant molecules with the surrounding water holds the micelles suspended in the water (see Figure 4.8). The suspended micelles form an emulsion of the grease-soap solution and water.

A difficulty with soaps is that they don't work well in hard water. The most common cations in hard water are calcium, magnesium, and iron ions. When you put soap into hard water, a precipitate of calcium, magnesium, or iron salts forms from an interaction of the cations with the carboxylate ions in the soap. This precipitate is the “bathtub ring” that you must scrub from your tub after you bathe using soap in hard water.

To solve this problem, chemists have developed a variety of surfactants with more soluble calcium or magnesium salts. One is sodium dodecanyl sulfate (sodium lauryl sulfate), CH3(CH2)11OSO3- Na⊕. The textile industry uses sodium lauryl sulfate as a detergent. It is also used as a wetting agent in photography and toothpaste. In an aqueous solution, a wetting agent works by lowering the surface tension of the solution below that of pure water.

A very important surfactant that lowers the surface tension of a solution is found in the transfer of oxygen to the bloodstream. This surfactant is necessary for life itself. A complex mixture of lipids and water coats the interior of the lungs. Simply described, this mixture consists of an aqueous solution of diapalmitoylphosphatidylcholine (DPPC), whose molecules have a hydrophilic head and a hydrophobic tail.

DPPC lowers the surface tension of the interior of the lungs, thereby increasing the rate of oxygen absorption. Polluted air and, particularly, cigarette smoke reduce the ability of the lungs to produce DPPC and thus inhibit the transfer of oxygen into the body.
DPPC also keeps your lungs inflated. DPPC coats the inside of the alveoli in your lungs and this coating lowers the surface tension on the inside of the alveoli. Outside the alveoli is the blood, which has a much higher surface tension. This difference in surface tension pulls the alveoli into a spherical shape.
Before birth, the fetus makes some respiratory-type movements, but its lungs remain collapsed. Immediately after birth the infant makes several strong inspiratory movements, and the lungs expand. The surfactant in the lining of the lungs keeps them inflated. However, when an infant is born before the surfactant system is functional, their lungs will not properly inflate or remain inflated. This problem, called hyaline membrane disease, is especially serious for premature infants who don't have a functional surfactant system. To help fight this condition, premature infants are given surfactants until their own surfactant system begins to function properly.

micelle

A micelle is an agglomeration of molecules that contain polar and nonpolar portions with one set of portions “dissolved” in one another and the other set of portions “dissolved” in the solvent.

Hydrophilic and Hydrophobic

Hydrophilic means “water-loving.” Hydrophobic means “water-hating.”

Surfactants

A surfactant is a molecule where one end of the molecule is soluble in one solvent and the other end of the same molecule is soluble in a different solvent. Surfactant is an acronym derived from the words surface active agent.

Recrystallization

A common laboratory purification technique is recrystallization. To begin, dissolve a solid compound in a minimum quantity of a hot solvent. On cooling, crystals of the original molecule form. The ideal solvent is one that does not dissolve the solute very well at low temperatures, but dissolves it readily at its boiling point. In general, the best solvent has a slightly lower polarity than the solute and thus the intermolecular interactions of solute with solvent are weaker than the intermolecular interactions of the solute. At higher temperature the solute becomes more soluble. Another useful characteristic for the solvent is a low boiling point, which makes it easier to remove solvent traces from the purified solute crystals. Finally, a good solvent should be a better solvent for any impurities than it is for the desired compound.

Solubility

Most chemical reactions take place in solution. That is, the solutes, or the reagents that you want to react, are uniformly mixed with, or dissolved in, the solvent. There are two types of solutions: single-phase solutions and multiple-phase solutions. In a single-phase, or homogeneous, solution all the solutes are soluble in the same solvent. Reactions take place best in single-phase solutions. This

uniform distribution allows a higher rate of reaction between the reactants because they have a greater amount of contact with each other. Single-phase solutions also allow you to control the solution concentrations and reaction conditions. Multiple-phase mixtures are mixtures in which the individual solutes are not soluble in the same solvent, thus, you must use more than one solvent. Each solvent then forms a different layer in the mixture. In multiple-phase mixtures, the reactants have less contact with each other so that the rate of reaction is generally slower than in a single-phase solution.

One compound is soluble in another as a result of the various intermolecular forces present in both compounds. These intermolecular forces are van der Waals forces, hydrogen bonding, and polarity. Molecular weight also plays a part in the solubility of a compound. The higher the molecular weight the lower the solubility. A compound dissolves in a solvent when the interactions between the compound and the solvent are either similar in strength or stronger than those interactions between molecules in the compound to be dissolved. There are also correlations between the molecular structure of the compounds and their solubility. A simple rule of thumb for estimating solubility is: A compound dissolves most easily in a solvent that is structurally similar to itself. The phrase “structurally similar” means that the solvent and the solute are similar types of molecules. For example, polar solutes or solvents mix with other polar solutes or solvents and nonpolar solutes or solvents mix with other nonpolar solutes or solvents, but polar solutes or solvents do not generally mix with nonpolar solutes or solvents.
Water is highly polar making it an excellent solvent for ionic compounds and small polar organic molecules. If an ethyl group

replaces one of the hydrogens in water, the result is ethanol, a solvent that has a substantially reduced polarity. Ethanol actually has both a polar and a nonpolar end. Although ethanol participates in hydrogen bonding and thus is a moderate solvent for salts and highly polar molecules, it also has van der Waals forces, making it a good solvent for a variety of organic molecules as well. Ethanol is too polar, however, to readily dissolve many low-polarity substances. Replacing both hydrogens in water with ethyl groups produces diethyl ether. Diethyl ether is only moderately polar. It acts as a hydrogen bond acceptor and so is a poor solvent for salts. However, it is a good solvent for a variety of both polar and nonpolar organic molecules.
Water is a poor solvent for low polarity compounds, such as gasoline. Gasoline, a mixture of alkanes, will not dissolve in water. All attempts to dissolve gasoline in water, or water in gasoline, result in a two-phase solution with the gasoline floating on the water. Gasoline won't dissolve in water for several reasons. (1) The molecular structures of the two substances are different. Gasoline, a mixture of hydrocarbons, is nonpolar, and water is highly polar. (2) The

molecular weights of the hydrocarbons are much larger in comparison to water. (3) The van der Waals forces between the molecules of gasoline are much stronger than they are in water. With water, primarily hydrogen bonding and secondarily dipolar attractions determine its intermolecular interactions. Both hydrogen bonding and dipolar attractions are stronger than van der Waals forces. Thus, gasoline does not dissolve in water because such a solution would reduce the number of stronger interactions. With gasoline and water, there is an unfavorable solvent-solute interaction.
Solubility is dependent on the temperature of the solution. In general, the higher the temperature the higher the solubility of a solute in a given solvent. For example, the solubility of benzoic acid in water is 1.7 g/L at 0oC; it increases to 68.0 g/L at 95oC. Figure 4.7 shows a plot of the solubility of benzoic acid in water versus temperature. Note that the solubility of benzoic acid increases rapidly as the temperature increases above75oC.

Boiling Points

An important physical property of a liquid is its boiling point, which is the transformation point of a compound from its liquid phase to its gas phase. The definition of boiling point is the temperature at which the vapor pressure of a liquid equals the external (atmospheric) pressure above that liquid.

When considering the boiling points for a homologous series of molecules, molecular weight is an important factor. Table 4.1 lists the boiling points for several common homologous series. As you look at the table, notice that within a series the boiling points quickly in-crease. For each series, the boiling point shows a fairly regular increase of 20-30oC with each additional CH2 group in the chain. Also, the boiling points vary drastically based on the functional group that the molecule contains.

The molecules in the liquid phase are still close enough to each other that their physical interactions are similar to the physical interactions that occur between the same molecules in their solid phase. Although the interactions in the liquid phase are far more random, and thus generally weaker, than in the solid phase, they still occur and are important. In the gas phase, the molecules are so far apart that the intermolecular interactions are much weaker and are generally of little importance. Except for highly polar substances, such as carboxylic acids, there are essentially no interactions between most molecules in the gas phase.
The interactions between molecules are the result of attractions between those molecules. These attractions fit into three different categories: van der Waals forces, dipolar attractions, and hydrogen bonding. The energies associated with these interactions are small compared to those associated with chemical bonds, but for a collection of molecules, they are significant.
When a pair of molecules approach each other, the nonbonded electrons on one molecule tend to attract the partially positive atoms on the other molecule. These attractive forces increase until they reach a maximum at intermolecular distances between 200 and 400 pm. At distances closer than this, the molecules tend to repel one another because their electrons repel one another. The actual distances at which the molecules begin to repel one another is the sum of the van der Waals radii of the two groups. The average distance between molecules in the liquid phase is in the range of 200-400 pm.
A complementary polarization occurs when a temporary dipole in one molecule induces a similar dipole in another molecule.
To understand what happens with these attractive forces, consider two nonpolar molecules, X and Y. Keep in mind that the distribution of electrons in these molecules is continually fluctuating. As the two molecules approach each other, they experience a mutual attraction, and any polarization in molecule X induces a complementary polarization in molecule Y. Chemists call this at-traction van der Waals forces or London dispersion forces.

 

van der Waals forces, which are sometimes called induced polarizations or induced dipoles, are only temporary and are constantly changing because the electron distribution within each molecule rapidly fluctuates. When the polarization in one molecule changes, it influences a neighboring molecule, which in turn influences another neighboring molecule. The net effect is that all neighboring molecules are attracted to each other. The magnitude of van der Waals forces is based on the number of electrons in the molecules and how many of those electrons participate in these induced dipole-dipole interactions.
For a low polarity liquid to boil, it must overcome the van der Waals forces. The major factor in the magnitude of these forces is the shape of the molecule. Highly branched molecules have a more spherical shape and smaller van der Waals attractions. Unbranched molecules have more surface area that can be involved in intermolecular interactions and higher van der Waals attractions because they can pack closer. You can see this effect in the boiling points of the following three isomers: pentane, 2-methylbutane, and 2,2-dimethylpropane.

Individual van der Waals forces are very weak. However, a typical molecule can participate in so many polarization interactions that the van der Waals forces are among the most important of the intermolecular forces in the liquid phase. They are the only forces possible for nonpolar molecules.

The second category of attractions that occurs between molecules is dipolar attractions. Molecules with permanent dipoles have dipolar attractions because of the charge polarization in their bonds. The interactions between molecules with permanent dipoles are similar to the van der Waals interactions between molecules with induced dipoles. The only difference is that the dipoles are permanent. Methyl fluoride illustrates the interaction of molecules with a permanent dipole. Methyl fluoride has a very polar C—F bond with a partial positive charge on the carbon and a partial negative charge on the fluorine atom:

 

In the liquid form, many other molecules of methyl fluoride surround each individual molecule of methyl fluoride. All these molecules tend to line up with the negative end of one dipole associated with the positive end of another:

As with van der Waals forces, molecules with dipole attractions require energy to overcome these forces. The dipolar forces raise the boiling point of methyl fluoride above that of a comparable compound without electronegative substituents. For example, methyl fluoride and ethane, have similar molecular weights, but methyl fluoride boils at –78oC whereas ethane boils at –89oC.
The third category of interactions that affects the boiling point is hydrogen bonding. Hydrogen bonding is a type of weak bonding interaction that involves a hydrogen bond donor and a hydrogen bond acceptor. A hydrogen bond donor is a molecule containing a hydrogen attached to an electronegative atom. The most common electronegative atoms in organic molecules are oxygen and nitrogen. A hydrogen bond acceptor is a molecule containing an atom with a nonbonding pair of electrons. The best hydrogen bond acceptors in organic molecules are also oxygen and nitrogen.

 

 

The strongest hydrogen bonds are with the O—H group. Weaker hydrogen bonds form with N—H bonds. Much weaker still are the hydrogen bonds formed with S—H and P—H bonds. The strength of an individual hydrogen bond is roughly 5 kcal/mole, much smaller than the typical covalent bond strengths of 80-100 kcal/mole. Hydrogen bonds are stronger than dipolar interactions, which are about 1-2 kcal/mole.
Of the three types of attractive forces, hydrogen bonding is the strongest. Hydrogen bonding substantially raises the boiling points of the compounds in which it occurs. For example, the isomeric compounds dimethyl ether and ethanol have widely different boiling points due to hydrogen bonding in ethanol.

Dimethyl ether has no hydrogen atoms attached to the oxygen, so no hydrogen bonding is possible. However, ethanol has a hydrogen attached to the oxygen, so hydrogen bonding occurs.
Chemists consider hydrogen bonding a very weak or partial bonding between an oxygen of one molecule and a hydrogen of another. This bonding causes aggregations, or groupings, of molecules much like those resulting from dipolar attractions. However, these molecular aggregations possess much more stability than those resulting from dipolar interactions.

 

Similar aggregations of molecules occur with amines. However, the boiling point differences of isomeric amines are less dramatic than for isomeric compounds of oxygen. For example, 3-methyl-1-butanamine boils at 95-96oC whereas N,N-dimethylpropanamine boils at 65oC.

 

The smaller difference in boiling points suggests that hydrogen bonds with N—H bonds are weaker than hydrogen bonds with O—H bonds. The N—H bonds are less polar because nitrogen has a lower electronegativity than oxygen. The hydrogen bonds are weaker because the hydrogen end of the dipole in the N—H bond is less positive than that in the O—H bond.

The previous discussion considers intermolecular hydrogen bonding. An additional factor comes into effect when two functional groups in one molecule participate in hydrogen bonding. The resultant intramolecular hydrogen bond is much more important than an intermolecular hydrogen bond in determining the properties of the molecule. For example, 2-nitrophenol has a much lower boiling point than either of its isomers, 3-nitrophenol or 4-nitrophenol.

2-Nitrophenol forms an intramolecular hydrogen bond between the hydrogen of the O—H bond and one of the oxygens in the NO2 group. This intramolecular hydrogen bond prevents an intermolecular hydrogen bond from forming. Thus, boiling requires much less energy for the 2-nitrophenol isomer than for the 4-nitrophenol isomer because there are no strong intermolecular forces to overcome in going from the liquid phase to the gas phase.

 

2-Nitrophenol represents a category of compounds in which intramolecular hydrogen bonding forms either a stable five- or six-membered ring. Section 3.5, page 000 discusses the relative stability of rings of various sizes. Because of this stability, they form more readily than rings of any other size. Any process that can result in a five- or six-membered ring is favored over one yielding another ring size. Expect to see this pattern repeatedly in your study of organic chemistry.

Melting Points

Melting occurs as a result of applying sufficient heat to a solid so that it moves from its solid phase to its liquid phase. The temperature at which this transformation takes place is the compound’s melting point. Each compound has a specific melting point; thus, when working with an unknown solid compound or a known solid compound whose purity is in doubt, one important measurement that chemists often take is the compound's melting point. Fortunately, taking a melting point is relatively easy to do. They then use this temperature to help identify the unknown solid compound or to verify that the compound is impure.

Four factors that influence a compound’s melting point are symmetry, polarity, hydrogen bonding, and molecular weight. Before looking at those factors, here is some information that you need to keep in mind about solids. As you learned in Section 4.1, a crystalline solid is composed of molecules arranged in a regular pattern. To arrange themselves in this pattern, the molecules normally have the same conformation and the same specific orientation relative to each other. Figure 4.6 shows a hypothetical crystalline solid. Above the solid are molecules A and B. They are both the same as the molecules in the solid, but neither molecule A nor molecule B can fit into the crystalline lattice. Molecule A has an incorrect conformation; and although molecule B has the correct conformation, it has an incorrect orientation.

The more symmetrical a molecule is, the better it fits into a crystalline structure, and the higher is its melting point. Because molecules must have a similar conformation to form a crystalline lattice, any molecules with the “wrong” symmetry do not easily fit into the growing crystal. For example, consider the following isomeric

compounds—pentane with a melting point of –130oC and 2,2-dimethylpropane with a melting point of –17oC. Pentane exists in a large variety of conformations that are similar in energy to each other. Below are three of its many possible staggered conformations.

Because the rotational energy barrier for the various conformations of pentane is very low, pentane exists as a mixture of conformations except at a very low temperature. Thus, pentane does not readily form a solid. It favors the liquid phase over the solid phase in the solid-liquid equilibrium.
In contrast to the conformational mobility of pentane, 2,2-dimethylpropane has only one conformation of the carbon skeleton:

All the hydrogen atoms of 2,2-dimethylpropane are equivalent, and the molecule has a very high level of symmetry. This symmetry means that each molecule has the “right” conformation to form a crystal. Also, this symmetry increases the probability that each molecule has the correct orientation to fit into a growing crystal. Thus, 2,2-dimethylpropane readily forms a crystalline lattice. It favors its solid form in the solid-liquid equilibrium and requires a higher temperature than pentane to convert from its solid form to its liquid form.
Another important factor contributing to the melting point of a compound is its polarity. The more polar a compound is, the stronger are its intermolecular attractions and the higher is its melting point. For example, benzyl alcohol melts at –15oC whereas benzoic acid melts at 122oC.

Both compounds are polar, but the intermolecular attractions of benzoic acid are much stronger than those of benzyl alcohol, as is shown by the fact that benzoic acid forms a relatively stable dimer. This dimer is held together by attractions, called hydrogen bonds, between the polar oxygens and the acidic hydrogens of the carboxylic acid functional groups. Hydrogen bonds influence the melting point of a solid, but they influence the boiling point of a liquid even more

The molecular weight of a compound influences its melting point. As the molecular weight increases in a homologous series of compounds, so does the melting point. The reason the melting point increases with the weight is that it takes more energy to separate larger molecules from a crystalline structure than it takes to separate smaller ones. Although the increase in melting point holds true over a long series, other factors at times override the effects of molecular weight for molecules of similar size. Two of these factors are symmetry and hydrogen bonding. For example, because methane molecules are symmetrical and pack into a crystalline lattice better than propane molecules do, methane’s melting point (–182.5 oC) is seven degrees higher than propane's (–189.7 oC) despite its lower molecular weight.