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.
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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.
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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.
5.4
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.