When finely divided clay particles are dispersed throughout water, they do not remain suspended but eventually settle out of the water because of the gravitational pull. The dispersed clay particles are much larger than molecules and consist of many thousands or even millions of atoms. In contrast, the dispersed particles of a solution are of a molecular size. Between these extremes is the situation in which dispersed particles are larger than molecules but not so large that the components of the mixture separate under the influence of gravity. These intermediate types of dispersions or suspensions are called colloidal dispersions, or simply colloids. Colloids are on the dividing line between solutions and heterogeneous mixtures. Like solutions, colloids can be gases, liquids, or solids. Examples of each are listed in Table 13.6.
The size of the dispersed particle is the property used to classify a mixture as a colloid. Colloid particles range in diameter from approximately 10 to 2000 Å. Solute particles are smaller. The colloid particle may consist of many atoms, ions, or molecules, or it may even be a single giant molecule. For example, the hemoglobin molecule, which carries oxygen in blood, has molecular dimensions of 65 Å 55 Å 50 Å and a molecular weight of 64,500 amu.
Although colloid particles may be so small that the dispersion appears uniform even under a microscope, they are large enough to scatter light very effectively. Consequently, most colloids appear cloudy or opaque unless they are very dilute. (Homogenized milk is a colloid.) Furthermore, because they scatter light, a light beam can be seen as it passes through a colloidal suspension, as shown in Figure 13.27. This scattering of light by colloidal particles, known as the Tyndall effect, makes it possible to see the light beam coming from the projection housing in a smoke-filled theater or the light beam from an automobile on a dusty dirt road.
The most important colloids are those in which the dispersing medium is water. Such colloids are frequently referred to as hydrophilic (water loving) or hydrophobic (water fearing). Hydrophilic colloids are most like the solutions that we have previously examined. In the human body the extremely large molecules that make up such important substances as enzymes and antibodies are kept in suspension by interaction with surrounding water molecules. The molecules fold in such a way that the hydrophobic groups are away from the water molecules, on the "inside" of the folded molecule, while the hydrophilic, polar groups are found on the surface, interacting with the water molecules. These hydrophilic groups generally contain oxygen or nitrogen. Some examples are shown in Figure 13.28.
Figure 13.28 Examples of hydrophilic groups on the surface of a giant molecule (macromolecule) that help to keep the molecule suspended in water.
Hydrophobic colloids can be prepared in water only if they are stabilized in some way. Otherwise, their natural lack of affinity for water causes them to separate from the water. Hydrophobic colloids can be stabilized by adsorption of ions on their surface, as shown in Figure 13.29. (Adsorption refers to adherence to a surface. It differs from absorption, which means passage into the interior, as when water is absorbed by a sponge.) These adsorbed ions can interact with water, thereby stabilizing the colloid. At the same time, the mutual repulsion between colloid particles with adsorbed ions of the same charge keeps the particles from colliding and so getting larger.
Figure 13.29 Schematic illustration of the stabilization of a hydrophobic colloid in water by adsorbed ions.
Hydrophobic colloids can also be stabilized by the presence of other hydrophilic groups on their surfaces. For example, small droplets of oil are hydrophobic. They do not remain suspended in water; instead, they separate, forming an oil slick on the surface of the water. Addition of sodium stearate, whose structure is shown below, or any similar substance having one end that is hydrophilic (polar, or charged) and one that is hydrophobic (nonpolar), will stabilize a suspension of oil in water, as shown in Figure 13.30. The hydrophobic ends of the stearate ions interact with the oil droplet, and the hydrophilic ends point out toward the water with which they interact.
Figure 13.30 Stabilization of an emulsion of oil in water by stearate ions.
These concepts have an interesting application in our own digestive system. When fats in our diet reach the small intestine, a hormone causes the gallbladder to excrete a fluid called bile. Among the components of bile are compounds that have chemical structures similar to sodium stearate; that is, they have a hydrophilic (polar) end and a hydrophobic (nonpolar) end. These compounds emulsify the fats present in the intestine and thus permit digestion and absorption of fat-soluble vitamins through the intestinal wall. The term emulsify means "to form an emulsion," a suspension of one liquid in another (Table 13.6). A substance that aids in the formation of an emulsion is called an emulsifying agent. If you read the labels on foods and other materials, you will observe that a variety of chemicals are used as emulsifying agents. These chemicals typically have a hydrophilic end and a hydrophobic end.
Colloidal particles frequently must be removed from a dispersing medium, as in the removal of smoke from stacks or butterfat from milk. Because colloidal particles are so small, they cannot be separated by simple filtration. The colloidal particles must be enlarged, a process called coagulation. The resultant larger particles can then be separated by filtration or merely by allowing them to settle out of the dispersing medium.
Heating or adding an electrolyte to the mixture may bring about coagulation. Heating the colloidal dispersion increases the particle motion and so the number of collisions. The particles increase in size as they stick together after colliding. The addition of electrolytes causes neutralization of the surface charges of the particles, thereby removing the electrostatic repulsions that inhibit their coming together. The effect of electrolytes is seen in the depositing of suspended clay in a river as it mixes with salt water. This results in the formation of river deltas wherever rivers empty into oceans or other salty bodies of water.
Semipermeable membranes can also be used to separate ions from colloidal particles, because the ions can pass through the membrane but the colloid particles cannot. This type of separation is known as dialysis. This process is used in the purification of blood in artificial kidney machines. Our kidneys are responsible for removing the waste products of metabolism from blood. In a kidney machine, blood is circulated through a dialyzing tube immersed in a washing solution. That solution is isotonic in ions that must be retained by the blood but is lacking the waste products. Wastes therefore dialyze out of the blood, but the ions do not.
A 0.100 L solution is made by dissolving a sample of CaCl2(s) in water. (a) Is CaCl2 an electrolyte or a nonelectrolyte? (b) The solution has an osmotic pressure of 3.55 atm at 27°C. What is the approximate molarity of CaCl2 in the solution? (c) The van't Hoff factor (see the "Closer Look" box in Section 13.5 of the textbook) for CaCl2 is 2.6 in the concentration range 0.04 - 0.12 M. Using this value, calculate the molarity of the CaCl2. (d) The enthalpy of solution for CaCl2 is H = -81.3 kJ/mol. If the final temperature of the solution was 27.0°C, what was its initial temperature? (Assume that the density of the solution is 1.0 g/mL, that its specific heat is 4.18 J/g-K, and that the solution loses no heat to its surroundings.)
SOLUTION (a) Soluble ionic compounds are strong electrolytes. CaCl2 consists of metal cations (Ca2+) and nonmetal anions (Cl–) and is hence a strong electrolyte.
(b) Solving Equation 13.13 for molarity, M, we have
This concentration is the total effective concentration of particles in the solution. Because each CaCl2 unit ionizes to form three ions (one Ca2+ and two Cl–), the concentration of CaCl2 is approximately 0.144 M/3 = 0.048 M. This concentration will be approximate because of ion pairing as described in the "Closer Look" box in Section 13.5 of the textbook.
(c) In the concentration range of 0.4-1.2 M CaCl2, the ion pairing reduces the effective number of free ions per CaCl2 unit from the ideal or limiting value of 3 to an actual value of 2.6. This number is the van't Hoff factor, i. Using this value, we calculate that the concentration of CaCl2 equals 0.144 M/2.6 = 0.055 M, just a bit higher than what we estimated in part (b).
(d) If the solution is 0.055 M in CaCl2 and has a total volume of 0.100 L, the number of moles of solute is (0.100 L)(0.055 mol/L) = 0.0055 mol. Hence the quantity of heat generated in forming the solution is (0.0055 mol)(-81.3 kJ/mol) = -0.45 kJ. The solution absorbs this heat, causing its temperature to increase. The relationship between temperature change and heat is given by Equation 5.19:
The heat absorbed by the solution is q = +0.45 kJ = 450 J. The mass of the 0.100 L of solution is (100 mL)(1.0 g/mL) = 100 g (to 2 significant figures). Thus the temperature change is
A kelvin has the same size as a degree Celsius. Because the solution temperature increases by 1.1°C, the initial temperature was 27.0°C - 1.1°C = 25.9°C.