|Like Dissolves Like||Hydrophilic and Hydrophobic Molecules|
|Soaps, Detergents, and Dry-Cleaning Agents||Units of Concentration|
By convention, we assume that one or more solutes dissolve in a solvent to form a mixture known as the solution. The photographs that accompany this section illustrate what happens when we add a pair of solutes to a pair of solvents.
|Solutes:||I2 and KMnO4|
|Solvents:||H2O and CCl4|
The solutes have two things in common. They are both solids, and they both have a deep violet or purple color. The solvents are both colorless liquids, which do not mix.
The difference between the solutes is easy to understand. Iodine consists of individual I2 molecules held together by relatively weak intermolecular bonds. Potassium permanganate consists of K+ and MnO4- ions held together by the strong force of attraction between ions of opposite charge. It is therefore much easier to separate the I2 molecules in iodine than it is to separate the ions in KMnO4.
There is also a significant difference between the solvents: CCl4 and H2O. The difference between the electronegativities of the carbon and chlorine atoms in CCl4 is so small (EN = 0.56) that there is relatively little ionic character in the CCl bonds.
Even if there was some separation of charge in these bonds, the CCl4 molecule wouldn't be polar, because it has a symmetrical shape in which the four chlorine atoms point toward the corners of a tetrahedron, as shown in the figure below. CCl4 is therefore best described as a nonpolar solvent.
The difference between the electronegativities of the hydrogen and oxygen atoms in water is much larger (EN = 1.24), and the HO bonds in this molecule are therefore polar. If the H2O molecule was linear, the polarity of the two OH bonds would cancel, and the molecule would have no net dipole moment. Water molecules are not linear, however, they have a bent, or angular shape. As a result, water molecules have distinct positive and negative poles, and water is a polar molecule, as shown in the figure below. Water is therefore classified as a polar solvent.
|Because water molecules are bent, or angular, they have distinct negative and positive poles. H2O is therefore an example of a polar solvent|
Because the solvents do not mix, when water and carbon tetrachloride are added to a separatory funnel, two separate liquid phases are clearly visible. We can use the relative densities of CCl4 (1.594 g/cm3) and H2O (1.0 g/cm3) to decide which phase is water and which is carbon tetrachloride. The denser CCl4 settles to the bottom of the funnel.
When a few crystals of iodine are added to the separatory funnel and the contents of the funnel are shaken, the I2 dissolves in the CCl4 layer to form a violet-colored solution. The water layer stays essentially colorless, which suggests that little if any I2 dissolves in water.
When this experiment is repeated with potassium permanganate, the water layer picks up the characteristic purple color of the MnO4- ion, and the CCl4 layer remains colorless. This suggests that KMnO4 dissolves in water but not in carbon tetrachloride. The results of this experiment are summarized in the table below.
Solubilities of I2 and KMnO4 in CCl4 and Water
This table raises two important questions. Why does KMnO4 dissolve in water, but not carbon tetrachloride? Why does I2 dissolve in carbon tetrachloride, but not water?
It takes a lot of energy to separate the K+ and MnO4- ions in potassium permanganate. But these ions can form weak bonds with neighboring water molecules, as shown in the figure below.
|KMnO4 dissolves in water because the energy released when bonds form between the K+ ion and the negative end of the neighboring water molecules and between the MnO4- ion and the positive end of the solvent molecules compensates for the energy it takes to separate the K+ and MnO4- ions.|
The energy released when these bonds form compensates for the energy that has to be invested to rip apart the KMnO4 crystal. No such bonds can form between the K+ or MnO4- ions and the nonpolar CCl4 molecules. As a result, KMnO4 can't dissolve in CCl4.
The I2 molecules in iodine and the CCl4 molecules in carbon tetrachloride are both held together by weak intermolecular bonds. Similar intermolecular bonds can form between I2 and CCl4 molecules in a solution of I2 in CCL4. I2 therefore readily dissolves in CCl4. The molecules in water are held together by hydrogen bonds that are stronger than most intermolecular bonds. No interaction between I2 and H2O molecules is strong enough to compensate for the hydrogen bonds that have to be broken to dissolve iodine in water, so relatively little I2 dissolves in H2O.
We can summarize the results of this experiment by noting that nonpolar solutes (such as I2) dissolve in nonpolar solvents (such as CCl4), whereas polar solutes (such as KMnO4) dissolve in polar solvents (such as H2O). As a general rule, we can conclude that like dissolves like.
|Practice Problem 1:
Elemental phosphorus is often stored under water because it doesn't dissolve in water. Elemental phosphorus is very soluble in carbon disulfide, however. Explain why P4 is soluble in CS2 but not in water.
|Practice Problem 2:
The iodide ion reacts with iodine in aqueous solution to form the I3-, or triiodide, ion.
I-(aq) + I2(aq) I3-(aq)
What would happen if CCl4 was added to an aqueous solution that contained a mixture of KI, I2, and KI3?
The family of compounds known as the hydrocarbons contain only carbon and hydrogen. Because the difference between the electronegativities of carbon and hydrogen is small (EN = 0.40), hydrocarbons are nonpolar. As a result, they do not dissolve in polar solvents such as water. Hydrocarbons are therefore described as immiscible (literally, "not mixable") in water.
When one of the hydrogen atoms in a hydrocarbon is replaced with an -OH group, the compound is known as an alcohol, as shown in the figure below. As might be expected, alcohols have properties between the extremes of hydrocarbons and water. When the hydrocarbon chain is short, the alcohol is soluble in water. Methanol (CH3OH) and ethanol (CH3CH2OH) are infinitely soluble in water, for example. There is no limit on the amount of these alcohols that can dissolve in a given quantity of water. The alcohol in beer, wine, and hard liquors is ethanol, and mixtures of ethanol and water can have any concentration between the extremes of pure alcohol (200 proof) and pure water (0 proof).
|The structure of the alcohol known as ethanol.|
As the hydrocarbon chain becomes longer, the alcohol becomes less soluble in water, as shown in the table below.
Solubilities of Alcohols in Water
|Formula||Name||Solubility in Water (g/100 g)|
|CH3(CH2)9OH||decanol||insoluble in water|
One end of the alcohol molecules has so much nonpolar character it is called hydrophobic (literally, "water-hating"), as shown in the figure below. The other end contains an -OH group that can form hydrogen bonds to neighboring water molecules and is therefore said to be hydrophilic (literally, "water-loving"). As the hydrocarbon chain becomes longer, the hydrophobic character of the molecule increases, and the solubility of the alcohol in water gradually decreases until it becomes essentially insoluble in water.
|One end of this alcohol molecule is nonpolar, and therefore hydrophobic. The other end is polar, and therefore hydrophilic.|
People encountering the terms hydrophilic and hydrophobic for the first time sometimes have difficulty remembering which stands for water-hating and which stands for water-loving. If you can remember that Hamlet's girlfriend was named Ophelia (not Ophobia), you might be able to remember that the prefix philo- is commonly used to describe love for example, in philanthropist, philharmonic, philosopher, and so on.
The data in the table above show one consequence of the general rule that like dissolves like. As molecules become more nonpolar, they become less soluble in water. The table below shows another example of this rule. NaCl is relatively soluble in water. As the solvent becomes more nonpolar, the solubility of this polar solute decreases.
Solubility of Sodium Chloride in Water and in Alcohols
|Formula of Solvent||Solvent Name||Solubility of NaCl (g/100 g solvent)|
The chemistry behind the manufacture of soap hasn't changed since it was made from animal fat and the ash from wood fires almost 5000 years ago. Solid animal fats (such as the tallow obtained during the butchering of sheep or cattle) and liquid plant oils (such as palm oil and coconut oil) are still heated in the presence of a strong base to form a soft, waxy material that enhances the ability of water to wash away the grease and oil that forms on our bodies and our clothes.
Animal fats and plant oils contain compounds known as fatty acids. Fatty acids, such as stearic acid (see figure below), have small, polar, hydrophilic heads attached to long, nonpolar, hydrophobic tails.
Fatty acids are seldom found by themselves in nature. They are usually bound to molecules of glycerol (HOCH2CHOHCH2OH) to form triglycerides, such as the triglyceride known as trimyristin, which can be isolated in high yield from nutmeg, shown in the figure below.
These triglycerides break down in the presence of a strong base to form the Na+ or K+ salt of the fatty acid, as shown in the figure below. This reaction is called saponification, which literally means "the making of soap."
|The saponification of the trimyristin extracted from nutmeg.|
Part of the cleaning action of soap results from the fact that soap molecules are surfactants they tend to concentrate on the surface of water. They cling to the surface because they try to orient their polar CO2- heads toward water molecules and their nonpolar CH3CH2CH2... tails away from neighboring water molecules.
Water can't wash the soil out of clothes by itself because the soil particles that cling to textile fibers are covered by a layer of nonpolar grease or oil molecules, which repels water. The nonpolar tails of the soap molecules on the surface of water dissolve in the grease or oil that surrounds a soil particle, as shown in the figure below. The soap molecules therefore disperse, or emulsify, the soil particles, which makes it possible to wash these particles out of the clothes.
Most soaps are more dense than water. They can be made to float, however, by incorporating air into the soap during its manufacture. Most soaps are also opaque; they absorb rather than transmit light. Translucent soaps can be made by adding alcohol, sugar, and glycerol, which slow down the growth of soap crystals while the soap solidifies. Liquid soaps are made by replacing the sodium salts of the fatty acids with the more soluble K+ or NH4+ salts.
Forty years ago, more than 90% of the cleaning agents sold in the United States were soaps. Today soap represents less than 20% of the market for cleaning agents. The primary reason for the decline in the popularity of soap is the reaction between soap and "hard" water. The most abundant positive ions in tap water are Na+, Ca2+, and Mg2+ ions. Water that is particularly rich in Ca2+, Mg2+, or Fe3+ ions is said to be hard. Hard water interferes with the action of soap because these ions combine with soap molecules to form insoluble precipitates that have no cleaning power. These salts not only decrease the concentration of the soap molecules in solution, they actually bind soil particles to clothing, leaving a dull, gray film.
One way around this problem is to "soften" the water by replacing the Ca2+ and Mg2+ ions with Na+ ions. Many water softeners are filled with a resin that contains -SO3- ions attached to a polymer, as shown in the figure below. The resin is treated with NaCl until each -SO3- ion picks up an Na+ ion. When hard water flows over this resin, Ca2+ and Mg2+ ions bind to the -SO3- ions on the polymer chain and Na+ ions are released into solution. Periodically, the resin becomes saturated with Ca2+ and Mg2+ ions. When this happens, it has to be regenerated by being washed with a concentrated solution of NaCl.
|When a water softener is "charged," it is washed with a concentrated NaCl solution until all of the -SO3- ions pick up an Na+ ion. The softener then picks up Ca2+ and Mg2+ ions from hard water, replacing these with Na+ ions.|
There is another way around the problem of hard water. Instead of removing Ca2+ and Mg2+ ions from water, we can find a cleaning agent that doesn't form insoluble salts with these ions. Synthetic detergents are examples of such cleaning agents. Detergent molecules consist of long, hydrophobic hydrocarbon tails attached to polar, hydrophilic -SO3- or -OSO3- heads, as shown in the figure below.
|The structure of one of the components of a synthetic detergent.|
By themselves, detergents don't have the cleaning power of soap. "Builders" are therefore added to synthetic detergents to increase their strength. These builders are often salts of highly charged ions, such as the triphosphate (P3O105-) ion.
Cloth fibers swell when they are washed in water. This leads to changes in the dimensions of the cloth that can cause wrinkles -- which are local distortions in the structure of the fiber or even more serious damage, such as shrinking. These problems can be avoided by "dry cleaning," which uses a nonpolar solvent that does not adhere to, or wet, the cloth fibers. The nonpolar solvents used in dry cleaning dissolve the nonpolar grease or oil layer that coats soil particles, freeing the soil particles to be removed by detergents added to the solvent, or by the tumbling action inside the machine. Dry cleaning has the added advantage that it can remove oily soil at lower temperatures than soap or detergent dissolved in water, so it is safer for delicate fabrics.
When dry cleaning was first introduced in the United States between 1910 and 1920, the solvent was a mixture of hydrocarbons isolated from petroleum when gasoline was refined. Over the years, these flammable hydrocarbon solvents have been replaced by halogenated hydrocarbons, such as trichloroethane (Cl3C-CH3), trichloroethylene (Cl2C=CHCl), and perchloroethylene (Cl2C=CCl2).
The concentration of a solution is defined as the amount of solute dissolved in a given amount of solvent or solution.
|Concentration||=||amount of solute|
|amount of solvent or solution|
There are many ways in which the concentration of a solution can be described.
|M||=||moles of solute|
|liters of solution|
|Practice Problem 3:
At 25oC, a saturated solution of chlorine in water can be prepared by dissolving 5.77 grams of Cl2 gas in enough water to give a liter of solution. Calculate the molarity of this solution.
Mass percent is literally the percentage of the total mass of a solution that is due to the solute.
|Mass percent||=||mass of solute||x||100%|
|mass of solution|
A 3.5% solution of hydrochloric acid, for example, has 3.5 grams of HCl in every 100 grams of solution. The concentration of a solution in units of moles per liter can be calculated from the mass percent and density of the solution.
It is also possible to describe the concentration of a solution in terms of the volume percent. This unit is used to describe solutions of one liquid dissolved in another or mixtures of gases. Wine labels, for example, describe the alcoholic content as 12% by volume, because 12% of the total volume is alcohol.
|Volume percent||=||volume of solute||x||100%|
|volume of solution|
Molarity is the concentration unit most commonly used by chemists. It has one disadvantage. It tells us how much solute we need to make a solution, and it gives us the volume of the solution produced, but it doesn't tell us how much solvent will be required to prepare the solution. We can make a 0.100 M solution of CuSO4, for example, by dissolving 0.100 mole of CuSO4 5 H2O in enough water to give one liter of solution. But how much water is enough? Because the CuSO4 5 H2O crystals occupy some volume, it takes less than a liter of water, but we have no idea how much less.
When it is important to know how much solute and solvent are present in a solution, chemists use two other concentration units: molality and mole fraction.
The molality (m) of a solution is defined as the number of moles of solute in the solution divided by the mass in kilograms of the solvent used to make the solution.
|Molality (m)||=||moles of solute|
|kilograms of solvent|
A 0.100 m solution of CuSO4, for example, can be prepared by dissolving 0.100 mole of CuSO4 in 1 kilogram of water. Because the density of water is about 1 g/cm3, or 1 g/mL, the volume of water used to prepare this solution will be approximately one liter. The total volume of the solution, however, will be larger than 1 liter because the CuSO4 5 H2O crystals will undoubtedly occupy some volume. As a result, a 0.100 m solution is slightly more dilute than a 0.100 M solution of the same solute.
|Practice Problem 4:
A saturated solution of hydrogen sulfide in water can be prepared by bubbling H2S gas into water until no more dissolves. Calculate the molality of this solution if 0.385 grams of H2S gas dissolve in 100 grams of water at 20oC and 1 atm.
Molality has an important advantage over molarity. The molarity of an aqueous solution changes with temperature, because the density of water is sensitive to temperature. Because molality is defined in terms of the mass of the solvent, not its volume, the molality of a solution does not change with temperature.
The ratio of solute to solvent in a solution can also be described in terms of the mole fraction of the solute or the solvent in a solution. By definition, the mole fraction of any component of a solution is the fraction of the total number of moles of solute and solvent that come from that component. The symbol for mole fraction is a Greek capital letter chi: C. The mole fraction of the solute is defined as the number of moles of solute divided by the total number of moles of solute and solvent.
|Csolute||=||moles of solute|
|moles of solute + moles of solvent|
Conversely, the mole fraction of the solvent is the number of moles of solvent divided by the total number of moles of solute and solvent.
|Csolvent||=||moles of solvent|
|moles of solute + moles of solvent|
In a solution that contains a single solute dissolved in a solvent, the sum of the mole fraction of the solute and the solvent must be equal to 1.
Csolute + Csolvent = 1
|Practice Problem 5:
Calculate the mole fractions of both the solute and the solvent in a saturated solution of hydrogen sulfide in water at 20oC and 1 atm.