Solutions of Molecular Substances in Liquids

The solubilities of nonpolar gases in water generally increase as the molecular mass of the gas increases, as shown in Table 14.1(1). This is precisely the trend expected: as the gas molecules become larger, the strength of the solvent–solute interactions due to London dispersion forces increases, approaching the strength of the solvent–solvent interactions.

Virtually all common organic liquids, whether polar or not, are miscible. If the predominant intermolecular interactions in two liquids are very different from one another, however, they may be immiscible. For example, organic liquids such as benzene, hexane, CCl₄, and CS₂ (S=C=S) are nonpolar and have no ability to act as hydrogen bond donors or acceptors with hydrogen-bonding solvents such as H₂O, HF, and NH₃; hence they are immiscible in these solvents. When shaken with water, they form separate phases or layers separated by an interface (Figure 14.1(d)), the region between the two layers.

The solubilities of simple alcohols in water are given in Table 14.1(2) "Solubilities of Straight-Chain Organic Alcohols in Water at 20°C". Only the three lightest alcohols (methanol, ethanol, and n-propanol) are completely miscible with water. As the molecular mass of the alcohol increases, so does the proportion of hydrocarbon in the molecule. Correspondingly, the importance of hydrogen bonding and dipole–dipole interactions in the pure alcohol decreases, while the importance of London dispersion forces increases, which leads to progressively fewer favorable electrostatic interactions with water. Organic liquids such as acetone, ethanol, and tetrahydrofuran are sufficiently polar to be completely miscible with water yet sufficiently nonpolar to be completely miscible with all organic solvents.

The same principles govern the solubilities of molecular solids in liquids. For example, elemental sulfur is a solid consisting of cyclic S₈ molecules that have no dipole moment. Because the S₈ rings in solid sulfur are held to other rings by London dispersion forces, elemental sulfur is insoluble in water. It is, however, soluble in nonpolar solvents that have comparable London dispersion forces, such as CS₂ (23 g/100 mL). In contrast, glucose contains five –OH groups that can form hydrogen bonds. Consequently, glucose is very soluble in water (91 g/120 mL of water) but essentially insoluble in nonpolar solvents such as benzene. The structure of one isomer of glucose is shown here.

Low-molecular-mass hydrocarbons with highly electronegative and polarizable halogen atoms, such as chloroform (CHCl₃) and methylene chloride (CH₂Cl₂), have both significant dipole moments and relatively strong London dispersion forces. These hydrocarbons are therefore powerful solvents for a wide range of polar and nonpolar compounds. Naphthalene, which is nonpolar, and phenol (C₆H₅OH), which is polar, are very soluble in chloroform.

Hydrophilic and Hydrophobic Solutes

A solute can be classified as hydrophilicA substance attracted to water. Hydrophilic substances are polar and can form hydrogen bonds to water. (literally, “water loving”), meaning that it has an electrostatic attraction to water, or hydrophobicA substance that repels water. Hydrophobic substances do not interact favorably with water. (“water fearing”), meaning that it repels water. A hydrophilic substance is polar and often contains O–H or N–H groups that can form hydrogen bonds to water. For example, glucose with its five O–H groups is hydrophilic. In contrast, a hydrophobic substance may be polar but usually contains C–H bonds that do not interact favorably with water, as is the case with naphthalene and n-octane. Hydrophilic substances tend to be very soluble in water and other strongly polar solvents, whereas hydrophobic substances are essentially insoluble in water and soluble in nonpolar solvents such as benzene and cyclohexane.

The difference between hydrophilic and hydrophobic substances has substantial consequences in biological systems. For example, vitamins can be classified as either fat soluble or water soluble. Fat-soluble vitamins, such as vitamin A, are mostly nonpolar, hydrophobic molecules. As a result, they tend to be absorbed into fatty tissues and stored there. In contrast, water-soluble vitamins, such as vitamin C, are polar, hydrophilic molecules that circulate in the blood and intracellular fluids, which are primarily aqueous. Water-soluble vitamins are therefore excreted much more rapidly from the body and must be replenished in our daily diet. A comparison of the chemical structures of vitamin A and vitamin C quickly reveals why one is hydrophobic and the other hydrophilic.

Because water-soluble vitamins are rapidly excreted, the risk of consuming them in excess is relatively small. Eating a dozen oranges a day is likely to make you tired of oranges long before you suffer any ill effects due to their high vitamin C content. In contrast, fat-soluble vitamins constitute a significant health hazard when consumed in large amounts. For example, the livers of polar bears and other large animals that live in cold climates contain large amounts of vitamin A, which have occasionally proven fatal to humans who have eaten them.

Solution Concentration

One important concept of solutions is in defining how much solute is dissolved in a given amount of solvent. This concept is called concentrationHow much solute is dissolved in a given amount of solvent.. Various words are used to describe the relative amounts of solute. DiluteA solution with very little solute. describes a solution that has very little solute, while concentratedA solution with a lot of solute. describes a solution that has a lot of solute. One problem is that these terms are qualitative; they describe more or less but not exactly how much.

Rather than just qualitative terms, we also need quantitative ways to express the amount of solute in a solution; that is, we need specific units of concentration. Here we will introduce several common and useful units of concentration.

MolarityThe number of moles of solute divided by the number of liters of solution. (M) is defined as the number of moles of solute divided by the number of liters of solution:

$$\text{molarity =}\frac{\text{moles of solute}}{\text{liters of solution}}$$

which can be simplified as

$$\text{M}=\frac{\text{mol}}{\text{L}}\text{, or mol/L}$$

As with any mathematical equation, if you know any two quantities, you can calculate the third, unknown, quantity.

For example, suppose you have 0.500 L of solution that has 0.24 mol of NaOH dissolved in it. The concentration of the solution can be calculated as follows:

$$\text{molarity}=\frac{\text{0}\text{.24 mol NaOH}}{\text{0}\text{.500 L}}=0.48\text{M NaOH}$$

The concentration of the solution is 0.48 M, which is spoken as "zero point forty-eight molarity" or "zero point forty-eight molar." If the quantity of the solute is given in mass units, you must convert mass units to mole units before using the definition of molarity to calculate concentration. For example, what is the molar concentration of a solution of 22.4 g of HCl dissolved in 1.56 L? First, convert the mass of solute to moles using the molar mass of HCl (36.5 g/mol):

$$\text{22}\text{.4}\cancel{\text{g HCl}}\times \frac{\text{1 mol HCl}}{\text{36}\text{.5}\cancel{\text{g HCl}}}=0.614\text{mol HCl}$$

Now we can use the definition of molarity to determine a concentration:

$$\text{M}=\frac{\text{0}\text{.614 mol HCl}}{\text{1}\text{.56 L}}=0.394\text{M}$$

Related concentration units are parts per thousand (ppth)Ratio of mass of solute to total mass of sample times 1,000., parts per million (ppm)Ratio of mass of solute to total mass of sample times 1,000,000., and parts per billion (ppb)Ratio of mass of solute to total mass of sample times 1,000,000,000.. Parts per thousand is defined as follows:

$$\text{ppth}=\frac{\text{mass of solute}}{\text{mass of sample}}\times \mathrm{1,000}$$

There are similar definitions for parts per million and parts per billion:

$$\text{ppm}=\frac{\text{mass of solute}}{\text{mass of sample}}\times \mathrm{1,000,000}$$ and $$\text{ppb}=\frac{\text{mass of solute}}{\text{mass of sample}}\times \mathrm{1,000,000,000}$$

Each unit is used for progressively lower and lower concentrations. The two masses must be expressed in the same unit of mass, so conversions may be necessary.

As with molarity, algebraic rearrangements may be necessary to answer certain questions.

For ionic solutions, we need to differentiate between the concentration of the salt versus the concentration of each individual ion. Because the ions in ionic compounds go their own way when a compound is dissolved in a solution, the resulting concentration of the ion may be different from the concentration of the complete salt. For example, if 1 M NaCl were prepared, the solution could also be described as a solution of 1 M Na⁺(aq) and 1 M Cl^-(aq) because there is one Na⁺ ion and one Cl^- ion per formula unit of the salt. However, if the solution were 1 M CaCl₂, there are two Cl^-(aq) ions for every formula unit dissolved, so the concentration of Cl^-(aq) would be 2 M, not 1 M.

In addition, the total ion concentration is the sum of the individual ion concentrations. Thus for the 1 M NaCl, the total ion concentration is 2 M; for the 1 M CaCl₂, the total ion concentration is 3 M.