5.3 Water as a Solvent
A major reason we must consume water is that it is an excellent solvent for many of the chemicals that make up our bodies, as well as for a wide variety of other substances. In this capacity, water acts as a solvent, a substance capable of dissolving other substances. Solutes are those substances that dissolve in a solvent. The resulting mixture is called a solution, a homogeneous mixture of uniform composition. Furthermore, aqueous solutions are solutions in which water is the solvent. Later in this chapter we will examine why certain kinds of substances dissolve in water and others do not. For now, we simply note that a remarkable variety of substances can dissolve in water and that this has important consequences for living organisms as well as for the environment. Table 5.1 summarizes some examples of water acting as a solvent.
Because water is such a good solvent, drinking water is rarely, if ever, just "pure" water. You can be assured that it almost certainly contains other substances. Municipal water companies provide information about the dissolved mineral content, the solutes, for tap water. An analysis of tap water in a Midwest home revealed the information in Table 5.2.
Table 5.1: Importance of Water as a Solvent
In our bodies:
• Blood plasma is an aqueous solution containing a variety of life-supporting substances.
• Inhaled oxygen dissolves in blood plasma in the lungs, allowing O2 to combine with hemoglobin.
• Blood plasma carries dissolved CO2 to the lungs to be exhaled.
• Blood plasma transports nutrients into all the cells and organs.
• Water helps to maintain a chemical balance by carrying wastes away.
In the environment:
• Water can transport toxic substances into, within, and out of living organisms.
• Water-soluble toxic substances, such as some pesticides, lead ions, and mercury ions, can be widely distributed.
• Water may reduce the concentrations of pollutants to safe levels by dilution or by carrying them away (or both).
• Rainwater carries substances, including those responsible for acid rain, from the atmosphere down to Earth.
Similar information can often be found on the labels or at the Web sites for commercial bottled water. For example, a label on Evian bottled water includes the information in Table 5.3.
Most of the solutes in Tables 5.2 and 5.3 will be discussed in this chapter. The number given with each dissolved ion indicates how much of that substance (in milligrams) is present in 1 L of water. This raises a reasonable question: Should we be concerned about the amounts of any of these substances? Calcium ions, for example, have a definite health benefit in producing stronger bones. Milk and milk products, not Evian water, is the preferred source for calcium ion; you would have to drink 4 L of Evian water to get the same amount of calcium ion as that in one 8-oz glass of milk. In contrast, the nitrate ion, depending on its concentration, can be dangerous, especially for infants. The other substances listed for Evian bottled water are not likely to cause a health problem. Elsewhere on the label it is noted that sodium (sodium ions), a health concern for some people, is present at less than 5 mg per 500-mL bottle.
Perhaps you have never considered drinking a glass of water as a risk-benefit act, yet it is. We usually consider water that has been chemically analyzed and treated to have important benefits with very low risk. Overwhelmingly, this is a valid assumption. But, however useful each tap water analysis or bottled water label may be, it necessarily is incomplete. As already noted, no information appears about health risks for the given concentrations of solutes. The information indicates nothing about whether other substances, if any, are present in the water. It does not indicate how much of each of the other substances is present or whether the substance might be harmful. For example, even though a tiny amount of lead is found in almost all water samples, it is usually in such low amount as not to be a health problem. If the water has been chlorinated to purify it, the water almost certainly has trace amounts of some chlorination by-products. Indeed, we rarely stop to think about what trace amounts of substances may be in the water, because we tend to assume that the water is safe to drink. In part, this is because extensive federal and state regulations and standards govern municipal water quality to protect the public. Most bottled water is regulated as well, often by self-imposed industry standards.
In assessing the health-giving or risk-taking aspects of drinking water, it is not sufficient to only know what substances are present in the water and how toxic they are. We also need to know how much of each substance is present in a particular amount of the water. In other words, we need to be able to understand what is meant by the concentration of a solute and the usual ways of expressing it. We now turn to these topics.
Solute Concentration in Aqueous Solutions
The concept of concentration was first introduced in Chapter 1 in relation to the composition of air. We used concentration units again in Chapters 2 and 3, looking at concentrations of chlorine compounds in the stratosphere or greenhouse gases accumulating in the troposphere. Now we will revisit this concept in terms of substances dissolved in water.
Although concentrations of components found in air might be a bit hard to visualize, solute concentrations in aqueous solutions are more familiar and therefore more easily imagined. For example, if you were asked to dissolve I teaspoon of an ingredient in 1 cup of water, a solution of a specific concentration would result: 1 tsp per cup (1 tsp/cup). Note that you would have the same 1 tsp/cup concentration if you also dissolved 2 tsp of the ingredient in 2 cups of water, 4 tsp in 4 cups, or 1/2 tsp in 1/2 cup. Even though you used larger or smaller quantities of the ingredient, the number of cups of water increased or decreased proportionally. Therefore, the concentration, the ratio of amount of ingredient to amount of water solution, would be the same in each case: 1 tsp per 1 cup (1 tsp/cup). Expressing solute concentrations in aqueous solutions follow the same pattern, but are often expressed in different units. We will use four ways of expressing concentration: percent; parts per million; parts per billion; and molarity. Three of these units are familiar to you from their use in earlier chapters, and molarity uses the mole concept introduced in Chapter 3. Each unit has particular application in various circumstances.
Percent: The most familiar way of expressing concentration is percent, defined in Chapter 1 as parts per hundred. For example, a solution containing 5 g of sodium chloride (NaCl) in 100 g of solution would be a five percent (5%) solution by weight. Hydrogen peroxide (H202) solutions, often found as an antiseptic in medicine cabinets, are usually 3% H202, indicating that they contain 3 g of H202 in 100 g of solution (or 6 g in 200 g of solution, etc.).
Ppm and ppb: Concentrations of dissolved substances in drinking water are normally far lower than 1% (1 part per hundred, pph). Correspondingly, different units are used to express such low concentrations. Parts per million (ppm) is the most common way of expressing the concentration of a solute in drinking water. A 1-ppm solution of calcium ion in drinking water contains 1 g of calcium ion in 1 million (1,000,000, or 106) g of that sample of drinking water, actually a dilute solution. The same concentration, 1 ppm, could be applied to a solution with 2 g of calcium ion in 2 106 g of water, 5 g in 5 106 g of water, or 5 mg (5 10–3 g) in 5000 (5 103) g of water. Although parts per million is a very useful concentration unit, measuring one million grams of water is not very convenient. Therefore, we look to find an easier but equivalent way to establish ppm. We find it using the unit liter (L), the volume occupied by 1000 g of water at 4 C. It is far easier to measure 1 L, a volume unit, rather than 1 106 g of water. Now we can say that 1 ppm of any substance in water equals 1 mg of that substance per 1 L of water.
1 ppm = (1 g solute)/(1,000,000 g water) = (1 mg solute)/(1,000 g water) = (1 mg solute)/(1 L water)
Drinking water contains substances naturally present at concentrations in the parts per million range, as illustrated on the Evian bottled water label. Toxic water pollutants also may be present in the parts per million concentration range. For example, the acceptable limit for nitrate ion, often found in well water in some agricultural areas, is 10 ppm; and the limit for the fluoride ion is 4 ppm.
Some pollutants are of concern at concentrations much lower than even parts per million so are reported as parts per billion (ppb). One part per billion of mercury (Hg) in water means 1 g Hg in 1 billion (1 109) g of water. In more convenient terms, this means 1 microgram (1 10–6 g, abbreviated as 1 g) Hg in 1 L (1 103 g) of water. For example, the acceptable limit for mercury in drinking water is 2 ppb.
2 ppb Hg = (2 g Hg)/(1 109 g H2O) = (2 10–6 g Hg)/(1 103 g H2O) = (2 g Hg)/(1 L H2O)
One part per million is a tiny concentration. Several analogies to a concentration of 1 ppm were given in Section 1.2, including that 1 ppm corresponds to 1 second in nearly 12 days. A similar analogy can be offered for parts per billion: 1 ppb corresponds to 1 second in 33 years, or approximately 1 inch on the circumference of the Earth.
Molarity: Another useful concentration unit in chemistry is molarity, which is based on the unit of the chemical mole. Molarity (M) is defined as the number of moles of solute present in one liter of solution.
Molarity (M) = (moles of solute)/(liter of solution)
The great advantage of molality is that a 1 molar (1 M) solution of any solute contains exactly the same number of chemical units (atoms or molecules) as any other 1 molar solution. The mass of solute may vary depending on the molar mass, but the number of chemical units will be the same for all 1 M solutions. Methods of chemical analysis of water (Section 5.15) frequently use molarity to express concentration. For now, we simply want to develop some familiarity with molarity itself.
As an example, consider a solution of NaCl in water. The molar mass of NaCl is 58.5 g; therefore, 1 mol of NaCl weighs 58.5 g. If we were to dissolve 58.5 g of NaCl in some water and then enough water to make exactly 1.000 L of solution, we would have a 1.00 M NaCl solution (Figure 5.6). Note the use of a volumetric flask, a type of glassware that contains a precise amount of solution when filled to the mark on its neck. But, there are many ways to make a 1 M NaCl solution. Another possibility, among many others, would be to use 0.500 mol NaCl (29.2 g) in 0.500 L of solution. This would require the use of a 500.0-mL volumetric flask, rather than the 1.000-L flask shown in Figure 5.6.
1 M NaCl = (1 mol NaCl)/(1 L solution) or (0.5 mol NaCl)/(0.5 L solution), etc.
So far in this chapter we have developed some ideas about drinking water, some of the substances that may be present in it, and how to express the concentrations of those substances. We shift now to a more detailed examination of water at the molecular level. Our aim is to understand water's unique properties, including its excellence as a solvent.
5.5 Water's Molecular Structure and Physical Properties
This section brings us to try and answer this important question—what is water? It is clear that water is essential to our lives and that water is an excellent solvent. What may not be as clear is that our dependence on water is possible only because water has a number of unusual properties. In fact, the physical properties of water are quite peculiar, and we are very fortunate that they are. If water were a more conventional compound we would be very different creatures.
This most common of liquids is full of surprises not the least of which is its physical state. Water is a liquid and not a gas at room temperature (about 25 C) and normal atmospheric pressure. This is surprising because almost all other compounds with similar molar masses to water's 18.0 g/mol are gases under similar conditions of temperature and pressure. Consider three common atmospheric gases (N2, O2, and CO2) whose molar masses are 28, 32, and 44 g/mol, respectively. All have molar masses greater than that of water, yet they are gases to breathe rather than liquids to drink.
Not only is water a liquid under these conditions, it also has an anomalously high boiling point of 100 C. This temperature is one of the reference points for the Celsius temperature scale. The other is the freezing point of water, 0 C. And when water freezes, it exhibits another somewhat bizarre property—it expands. Most liquids contract when they solidify. These and other unusual properties derive from water's chemical composition and molecular structure. We will continue to explore the reasons for water's unusual behaviors as this section continues.