How To Homemade Iron Can Be Converted To Dichromate Ion And Vice Versa

Solving for the carbonate concentration, we obtain fco32-]

and hence

S =/[C032! = (fKsp)1

= {Ksp( 1 + Kb/[OH-])}1/2

Thus, as expected by application of Le ChStelier's principle to the reactions, the solubility of calcium carbonate decreases as the (fixed) hydroxide concentration increases, the limit at high levels of OH" being (Ksp)^2, the value we obtained assuming no reaction of carbonate ion with water. In contrast, for water that is neutral or acidic and therefore low in hydroxide ion, the CaC03 solubility is much larger than this value, as illustrated in Figure 13-8, where the logarithm of S is plotted against the pH of the water body.

Solubility Iron Carbonate
FIGURE 13-8 Molar solubility, 5, Iri units of

KSp, of CaCOj in C02-free water versus pH.

Water in Equilibrium with Both CaC03 and Atmospheric C02

The systems discussed above are somewhat unrealistic since they fail to consider the other important carbon species in water—namely, carbon dioxide and carbonic acid—and the reactions that involve them. These reactions will now be considered in the context of a body of water that is also in equilibrium with solid calcium carbonate, i.e., the three-phase system illustrated in Figure 13-7.

At first sight, it might seem that since reaction (1) provides another source of bicarbonate ion, then by Le Châtelier's principle the production of bicarbonate from the reaction (5) of carbonate with water should be suppressed. However, a more important consideration is that reaction ( 1 ) produces hydrogen ion, which combines with the hydroxide ion that is produced in reaction (4) by the interaction of carbonate ion with water:

H+ + OH" HzO(aq) (7)

Consequently, the equilibrium positions of both reactions that produce bicarbonate ion are shifted to the right due to the disappearance of one of their products by the above reaction.

If reactions (1), (3), (4), (5), and (7) are all added together to deduce the net process, then after canceling common terms the net result is

CaC03(s) + C02(g) + H20(aq) ^^ Ca2+ + 2 HCQ3" (8)

In other words, combining equimolar amounts of solid calcium carbonate and atmospheric carbon dioxide yields aqueous calcium bicarbonate, Ca(HC03)2, without any apparent production or consumption of acidity or alkalinity:

calcium carbonate (rock) + carbon dioxide (air) = calcium bicarbonate

(in solution)

Natural waters in which this overall process occurs can be viewed as the site of a giant titration of an acid that originates with C02 from air with a base that originates with carbonate ion from rocks. [Note that we need not consider reaction (2) in this analysis, since reaction (5) of the conjugate base of bicarbonate with water was included.] In the ocean neutral scheme of sequestration (Chapter 7), bulk carbon dioxide is reacted with solid calcium carbonate or with some other calcium-containing salt; the resulting slurry of calcium bicarbonate (or other salt) is then transported to and deposited in the ocean. This technique avoids the pH-lowering side effect of the ocean acidic scheme in which COz is directly dissolved in the ocean water.

It should be noted that each of the individual reactions added together is itself an equilibrium that does not lie entirely to the right. Since the reactions differ in their extent of completion, it is an approximation to state that the overall reaction shown above is the only resulting reaction. Nevertheless, it is the dominant process, and it is mathematically convenient to first consider this process alone in estimating the extent to which CaC03 and C02 dissolve in water when both are present.

Since reaction (8) equals the sum of reactions (1), (3), (4), (5), and (7), its equilibrium constant Kg is the product of their equilibrium constants;

Kg = KspKbKHKal/Kw

Here Kai is the first acid dissociation constant for carbonic acid. KH is the Henry's law constant (see Chapter 3) for reaction (3). Since Kw is the ion product for water, the equilibrium constant for reaction (7) is 1/KW. The other constants in the equation for Kg have been defined previously (see Table 13-3). Thus at 25°C for the overall reaction (8), it follows that

From the balanced equation for the reaction, it follows that the expression for Kg is

K8 = [Ca2+] [HC0312/PC02

If the calcium concentration again is called S, then from the stoichiom-etry of reaction (8) the bicarbonate concentration must be twice as large, equal to 2S; after substitution for the concentrations in the equation for Ks and rearrangement, we obtain

4s3 = k8pco2

Thus the solubility of calcium carbonate increases as the cube root of the partial pressure of carbon dioxide to which the water is exposed:

S = (K8PCo2/4)1/3

Substituting the current partial pressure of C02 in the atmosphere, 0.00036 atm, corresponding to an atmospheric concentration of C02 of 365 ppm (Chapter 6), and the numerical value of K8 into this equation yields

S = 5.1 X 10~4 mol/L-1 = [Ca2+]

and thus

The amount of C02 dissolved is also equal to S and is 35 times that which dissolves without the presence of calcium carbonate (see results of Problem 3-10). Furthermore, the calculated calcium concentration is four times that calculated without the involvement of carbon dioxide. Thus the acid reaction of dissolved C02 and the base reaction of dissolved carbonate have a synergistic effect on each other that increases the solubilities of both the gas and the solid (see Table 13-4). In other words, water that contains carbon dioxide more readily dissolves calcium carbonate. In fact, groundwater may

Calculated Ion Concentrations for Aqueous Equilibrium Systems

Ion

CaCOj

C02 and CaCO,

[HCOjl

9.9 X 10~5

M

1.0 X 10"

"3 M

ICO,2 ]

8.8 X 10

M

[Caz+]

9.9 X 10"5

M

5.2 X 10"

4 M

[OH"]

9.9 X 10^5

M

1.8 X 10"

"6M

[Hi

1.0 X 10"1C

1M

5.6 X 10"

"9 M

pH

10.0

8.3

become supersaturated with carbon dioxide as a result of biological decomposition processes; and in that case, the calcium carbonate solubility increases even more—at least until the water reaches the surface, when degassing of the excess C02 would occur.

PROBLEM 13-15

Repeat the above calculation for the solubility of CaC03 in water that is also in equilibrium with atmospheric C02 for a water temperature of 5°C. At this temperature, KH = 0.065 for C02 and for H2C03 is 3.0 X 10"7; see Problem 13-13 for other necessary data.

Finally, the residual concentrations of C03 , of H+, and of OH" in the system can be deduced from equilibrium constants for reactions (4), (5), and (7), since equilibria in these processes are in effect, notwithstanding the overall reaction (8). Thus from reaction (4),

[C032~] = Ksp/[Caz+] = 4.6 X 10"9/5.1 X 10"4 = 9.0 X 10"6M

From reaction (5),

[OH"] = kb [C032"]/[HC03-]

= (2.1 X 10"4) .X (9.0 X 10"6)/1.0 X 10"3 = 1.9 X 10"6

and finally from reaction (7)

[H+] = Kw/[OH~] = 1.0 X 10"14/l-9 X 10"6 = 5.3 X 10"9

From this value for the hydrogen ion concentration, we conclude that according to this calculation, river and lake water at 25 °C whose pH is determined by saturation with C02 and CaC03 should be slightly alkaline, with a pH of about 8.3.

Typically, the pH values of calcareous waters lie in the range from 7 to 9, in reasonable agreement with our calculations. Due to the smaller amount of bicarbonate in noncalcareous waters, their pH values are usually close to 7. Of course, if natural waters are subject to acid rain, the pHs can become substantially lower since there is little HC03~ or C032" readily available with which to neutralize the acid.

About 80% of natural surface waters in the United States have pH values between 6.0 and 8.4. Lakes and rivers into which acid rain falls will have elevated levels of sulfate ion and perhaps of nitrate ion since the principal acids in the precipitation are H2S04 and HN03 (see Chapter 4).

PROBLEM 13-16

In waters subject to acid rain, the pH is determined not by the C02~carbonate system but rather by the strong acid from the precipitation. Assuming that equilibrium with atmospheric carbon dioxide is in effect, calculate the concentration of HCOj in natural waters with pH = 6, 5, and 4 at 25°C. (See Table 13-3 for data.)

PROBLEM 13-17

Using algebraic expressions and numerical values for Kal and jfC,2 of H2C03 (Table 13-3), calculate the pi 1 values for which [H2C03] = [HC03~] and for which [HC03~] = [C032"].

Ion Concentrations in Natural Waters and Drinking Water

The Abundant Ions in Fresh Water

As is evident from Table 13-5, the most abundant ions found in samples of unpolluted fresh calcareous water usually are calcium and bicarbonate, as expected from our previous analysis. Commonly, such water also contains magnesium ion, Mg2^, principally from the dissolution of MgC03; plus some sulfate ion, S042 ; smaller amounts of chloride ion, CI , and sodium ion, Na~; and even smaller levels of fluoride ion, F~, and potassium ion, K7. The overall reaction (8) of carbon dioxide and calcium carbonate implies that the ratio of bicarbonate ion to calcium ion should be 2:1, and this is indeed a rule that is closely obeyed on average in river water in North America and Europe. The calculated calcium ion concentration, 5.1 X M, agrees well with the North American river-water average value of 5.3 X 1CT4 M, and similarly for the bicarbonate ion data. The close agreement between the calculated and the experimental results is somewhat fortuitous because river-water temperatures on average lie below 25 °C—which results in a higher C02 solubility than has been assumed— and because several minor factors have been oversimplified in the calculation. In fact, even calcareous river water is usually unsaturated with respect to CaC03.

Water in rivers and lakes that is not in contact with carbonate salts contains substantially fewer dissolved ions than are present in calcareous waters. The concentration of sodium and potassium ions may be as high as those of calcium, magnesium, and bicarbonate ions in these fresh waters. Even in areas with no limestone in the soil, the waters contain some bicarbonate ion due to the weathering of aluminosilicates in submerged soil and rock in the j River Water Concentrations and Drinking Water I Standards for Ions

River Water Molar Concentration

Average for World

Average for U.S.

Drinking Water Concentration in ppm

Maximum Recommended Concentration

Average U.S.

Canada

*hco3~

9.2 X 10"4

9.6 X 10~4

60

Ca2+

3.8 X 1CT4

3.8 X 10~4

15

Mg2+

1.6 X 10"4

3.4 X 10"4

8

Na+

3.0 X 10"4

2.7 X 10~4

6

200

cr

2.3 X 10"4

2.2 X 10"4

8

250

250

so42-

1.1 X 10"4

1.2 X 10"4

12

250

500

K+

5.4 X 10~5

5.9 X 10~5

2

F"

5.3 X 10~6

0.1

0.8-2.4

1.5

no3~

1.4 X 10~5

Fe3+

7.3 x 10~6

*Note: The value for bicarbonate is actually the total alkalinity.

Sources: World data from R. A. Larson and E. J. Weber, Reaction Mechanisms in Environmental Organic Chemistry (Boca Raton, FL: Lewis Publishers).

*Note: The value for bicarbonate is actually the total alkalinity.

Sources: World data from R. A. Larson and E. J. Weber, Reaction Mechanisms in Environmental Organic Chemistry (Boca Raton, FL: Lewis Publishers).

presence of atmospheric carbon dioxide. The weathering reaction can be written in general terms as

M+(Al-silicate~)(s) + COz(g) + H20->M+ + HC03~ + H4Si04

Here M is a metal such as potassium, and the anion is one of the many alumi-nosilicate ions found in rocks (see Chapter 16). The weathering of potassium feldspar is an example of one of the most important sources of potassium ion in natural waters:

3 KAlSi3Og(s) + 2 C02(g) + 14 HzO-»

2 K+ + 2 HC03~ + 6 H4Si04 + KAl3Si3O10(OH)2(s)

Thus bicarbonate normally is the predominant anion in both calcareous and noncalcareous waters since it is produced by the dissolution of limestone and aluminosilicates, respectively.

The average compositions of river water in the United States and in the world as a whole are given in Table 13-5. As discussed, the values for the calcium and magnesium ion concentrations vary significantly from place to place, depending upon whether or not the underlying soil is calcareous.

Fresh water in which the concentration of ions is abnormally high is called saline water and is usually unsuitable for drinking. Most saline water is the result of irrigation, in which water is transported into land where little rainfall occurs. The water largely evaporates if the climate is hot and dry, leaving behind salts of the ions that were present in the irrigation water. The runoff water from rainfall and irrigation into water supplies consequently is saline. If the irrigation water is recycled, it becomes more and more saline as time goes on. The wintertime de-icing of roads in northern climates also contributes salinity to water bodies.

Fluoride Ion in Water

The level of fluoride ion, F , in water also displays substantial variations, from less than 0.01 ppm to more than 20 ppm, in different regions of the world. The source of most F~ is weathering of the mineral fluorapatite, Ca5(P04)3F.

In Mexico and some European countries, sodium fluoride, NaF, is added to table salt. In many communities of English-speaking countries—including the United States (about half the population), Canada, Australia, and New Zealand in which the F~ concentration in the drinking-water source is low— a soluble fluoride compound such as fluorosilicic acid, H2SiF6, or its sodium salt, both of which react with water to release fluoride ion, is often added in order to bring the fluoride level up to about 1 ppm, i.e., 5 X 10~5 M. This value was considered at least in the past to be optimum in strengthening children's teeth against decay while providing a margin of safety. If the fluoride level is in excess of this value, as it is in some natural waters, deleterious effects on teeth such as mottling can occur. The maximum contaminant level (MCL) of fluoride in U.S. drinking water is 4 ppm. Almost all brands of toothpaste available in developed countries contain added fluoride in the form of sodium fluoride, stannous fluoride, SnF2, or sodium monofluorophosphate. Most children in North America also receive topical fluoride not only from their toothpaste, but in some cases by application from their dentists.

The addition of fluoride ion to public supplies of drinking water continues to be a controversial subject because at high concentrations fluoride is known to be poisonous and perhaps carcinogenic, and because some people feel that it is immoral to force everyone to drink water to which a substance has been added. In fact, for many people, the total amount of fluoride ion ingested from food and beverages (especially tea) exceeds that from water.

Bottled Drinking Waters

The maximum concentration of ions recommended for drinking water in the United States and in Canada is also listed in Table 13-5. The concentration of sodium ion, Na+, in water is of interest since high consumption of it from water and salted food is believed to increase blood pressure, which may lead to cardiovascular disease. Excessive sulfate, beyond 500 mg/L, causes a laxative effect in some people. It is interesting to note that some varieties of bottled drinking water, which people presumably drink in preference to tap water due to health concerns about the latter, exceed the recommended values for some ions. Several well-known bottled waters exceeded the drinking water standards for arsenic and/or fluoride in a 1999 survey by the U.S. National Resources Defense Council. However, they were remarkably free of chloroform, a substance that plagues municipal water supplies, as discussed in Chapter 14. A more recent survey found that most brands would meet the new 10-ppm standard for arsenic (Chapter 15) but that bisphenol-A (Chapter 12) is leached from the plastic into the water contained in most large polycarbonate jugs.

Suppliers of bottled water sometimes advertise their products as having "zero" concentrations of fluoride and/or sodium ions or as being sodium-free or fluoride-free. These are misleading statements since in reality the actual concentrations are not zero but below the level of detection in the analytical method used by the bottler or below a threshold specified by government. "Zero" is not a meaningful chemical concept to answer the question of "how much" of a substance is present in a sample.

Seawater

The total concentration of ions in seawater is much higher than that in fresh water since it contains large quantities of dissolved salts. The predominant species in seawater are sodium and chloride ions, which occur at about 1000 times their average concentration in fresh water. Seawater also contains

some Mg and S04 and lesser amounts of many other ions. If seawater is gradually evaporated, the first salt to precipitate is CaC03 (present to the extent of 0.12 g/L), followed bv CaSQ4 • H20 (1.75 g/L), then NaCl (29.7 g/L), MgSQ4 (2.48 g/L), MgCl2 (3.32 g/L), NaBr (0.55 g/L), and finally KC1 (0.53 g/L). Thus "sea salt" is a mixture of all these salts, which together constitute about 3.5% of the mass of seawater. Due primarily to the operation of the C02-bicarbonate—carbonate equilibrium system discussed previously for fresh water, the average pH of surface ocean water is about 8.1. Seawater has a low organic content, its DOC value being about 1 mg/L.

Alkalinity Indices for Natural Waters

The actual concentrations of the cations and anions in a real water sample cannot simply be assumed to be the theoretical values calculated above for calcium, carbonate, and bicarbonate for two reasons:

• the water may not be in equilibrium with either solid calcium carbonate or with atmospheric C02; and

• other acids or bases may also be present.

The index devised by analytical chemists to represent the actual concentration in water of the anions that arc basic is provided by the alkalinity value for the sample. Alkalinity is a measure of the ability of a water sample to act as a base by reacting with hydrogen ions. In practical use, the alkalinity of a body of water is a handy measure of the capacity of the water body to neutralize acids and hence to resist acidification when acid rain falls into it. (Alkalinity differs from [OH-]. or pOH, in that the latter equals only the hydroxide concentration a water sample has at a particular moment and does not include its ability to generate additional OH" if acid is added to it.) From an operational viewpoint, alkalinity, more properly termed total alkalinity, is the number of moles of H+ required to titrate 1 liter of a water sample to the end point. For a solution containing carbonate and bicarbonate ions, as well as OH~ and H+, by definition

(total) alkalinity = 2 |(XY J + [HCO, 1 + [OH ] - [H+]

The factor of 2 appears in front of carbonate ion concentration since in the presence of H+ it is first converted by the ion to bicarbonate ion, which is then converted by a second hydrogen ion to carbonic acid:

CO32" + H+ ^^ HC03" 1 ICO, + H+ ^^ H2C03

Minor contributors to the alkalinity of fresh-water systems can include dissolved ammonia and the anions oiphosphoric, boric, and silicic acids, and H2S, as well as natural organic matter.

The alkalinities of natural waters range from less than 5 X 10 M (50 fxM) to more than 2 X 10" 3 M (2000 ¿uM), compared to the value of about 3 X 10^4 M that corresponds to our estimate for water in contact with atmospheric C02 (see Problem 13-19). Lakes having alkalinities less than about 200 fiM are considered "high" in sensitivity to acid rain, those from 200 to 400 fiM are classified as "moderate" in sensitivity, and those with alkalinity greater than 400 fxM are considered to have "low" sensitivity. Alkalinity values are sometimes reported as milligrams of CaC03 equivalent, rather than moles of H+, per liter in a manner similar to that explained below for the hardness concept.

By convention in analytical chemistry, methyl orange is used as the indicator in titrations by which total alkalinity is determined. Methyl orange is chosen because it does not change color until the solution is slightly acidic (pH = 4); under such conditions, not only has all the carbonate ion in the sample been transformed to bicarbonate, but virtually all the bicarbonate ion has been transformed to carbonic acid (see Problem 13-18 and Figure 13-6).

Another index encountered in the analysis of natural waters is the phe-nolphthalein alkalinity (also called carbonate alkalinity), which is a measure of the concentration of the carbonate ion and of other similarly basic anions.

In order to titrate only C032" and not HCO, as well, the indicator Phenolphthalein or one with similar characteristics is used. Phenolphthalein changes color in the pH range 8 to 9, so it provides a fairly alkaline end point. At such pH values, only a negligible amount of the bicarbonate ion has been converted to carbonic acid, but the majority of CO;" has been converted to HC03" (Figure 13-6). Thus,

Phenolphthalein alkalinity = [C032~]

PROBLEM 13-18

Calculate the value of the ratios IHCOj j/[C032"] and [H2C03]/[HC03~] at pH values of 4 and 8.5 to confirm the statements made above concerning the nature of the species present at the methyl orange and Phenolphthalein end points of the titrations. [Hint: Use the equilibrium constant expressions and K values for reactions (I) and (2).]

PROBLEM 13-19

Calculate the value expected for the total alkalinity and for the Phenolphthalein alkalinity of a 25QC saturated solution of calcium carbonate in water that is also in equilibrium with atmospheric carbon dioxide. Use the concentrations quoted in the last column of Table 13-4.

PROBLEM 13-20

Calculate the total alkalinity for a sample of river water whose Phenolphthalein alkalinity is known to be 3.0 X 10~5 M, whose pH is 10.0, and whose bicarbonate ion concentration is 1.0 X 10"4M.

The alkalinity value for a lake is sometimes used by biologists as a measure of its ability to support aquatic plant life, a high value indicating a high potential fertility. The reasons for such a situation are often the following ones. Algae extract the carbon dioxide they need for photosynthesis from bicarbonate ion, which is plentiful in calcareous waters, by a reversal of the C02-CaC03 reaction discussed previously:

Ca2+ + 2 HCO," (aq)->C02 + CaC03(s) + H20

Indeed, small crystals of calcium carbonate are sometimes observed in lakes where there is active photosynthesis.

C02 + H20 + sunlight-* CH20 polymer + | 02

(as algae)

In noncalcareous waters, which have low alkalinity and low calcium content, dissociation of the bicarbonate ion in the water forms not only carbon dioxide but also hydroxide ion:

The algae readily exploit this C02 for their photosynthetic needs, at the cost of allowing a buildup of hydroxide ion to such an extent that the lake water becomes quite basic—with a pH as high as 12.3 in some cases.

The Hardness Index for Natural Waters

As a measure of certain important cations present in samples of natural waters, analytical chemists often use the hardness index, which measures the total concentration of the ions Ca2+ and Mg2+, the two species that are principally responsible for hardness in water supplies. Chemically, the hardness index is defined in this way:

Experimentally, hardness can be determined by titrating a water sample with ethylenedkminetetraacetic acid (EDTA), a substance that forms very strong complexes with metal ions other than those of the alkali metals (see Chapter 15 for details). Traditionally, hardness is expressed not as a molar concentration of ions but as the mass in milligrams (per liter) of calcium carbonate that contains the same total number of (¡¿positive (2+) ions, For example, a water sample that contains a total of 0,0010 mole of Ca2+ + Mgz+ per liter would possess a hardness value of 100 mg ofCaC03, since the molar mass of CaCOj is 100 g and 0.0010 mole of it weighs 0.1 g, or 100 mg.

Most calcium enters water from either CaC03 in the form of limestone or from mineral deposits of CaS04. The source of much of the magnesium is dolmitic limestone, CaMg(C03)2. Hardness is an important characteristic of natural waters, since calcium and magnesium ions form insoluble salts with the anions present in soaps, thereby forming a scum in wash water. Water is termed "hard" if it contains substantial concentrations of calcium and/or magnesium ions; thus calcareous water is "hard." Some scientists define water as being hard if its hardness index exceeds 150 mg/L.

Many areas possess soils that contain little or no carbonate ion, and thus its dissolution and reaction with C02 to produce bicarbonate do not occur. Such "soft" water typically has a pH much closer to 7 than does hard water, since it contains few basic anions. However, there are lakes with little dissolved calcium or magnesium but relatively high concentrations of dissolved sodium carbonate, Na2C03; such lakes have a very low degree of hardness but are high in alkalinity.

Interestingly, people who live in hard-water areas are found to have a lower average death rate from ischemic heart disease than do people living in areas with very soft water. Recent research in rural Finland found the risk of heart attacks decreased continuously as the magnesium concentration in the local water supply increased.

PROBLEM 13-21

What is the value of the hardness index in milligrams CaC03/L for a 500-mL sample of water that contains 0.0040 g of calcium ion and 0.0012 g of magnesium ion?

PROBLEM 13-22

Calculate the hardness, in milligrams CaC03/L, of water that is in equilibrium at 25°C with carbon dioxide and calcium carbonate, using results in the last column of Table 13-4- Assume the water is free of magnesium. Is the calculated value greater or less than the median hardness value found for surface waters of 37 mg/L?

Aluminum in Natural Waters

The concentration of aluminum ion, Al3+, in natural waters normally is quite small, typically about 10 6 M. This low value is the consequence of the fact that in the typical pH range for natural waters (6 to 9), the solubility of the aluminum contained in rocks and soils to which the water is exposed is very small. The solubility of aluminum in water is controlled by the insolubility of aluminum hydroxide, Al(OH)3. Given that the Ksp of the hydroxide is about lO"33 at usual water tertiperatures, then for the reaction

it follows that

[Al3+] [OH"]3 = 10"33

Take, for instance, a sample of water whose pH is 6. Since the hydroxide concentration in such water is 10"8 M, it follows that

[Al3+] = 10"33/(10"8)3 = 10"9M

Although this value is very small, for every one-unit decrease of the pH, the concentration of aluminum ion increases by a factor of 1000, so it reaches 10 6 M at pH = 5 and 10^3 M at a pH of 4- Thus aluminum is much more soluble in highly acidified rivers and lakes than in those where pH values do not fall below 6 or 7. Indeed, Al3+ is usually the principal cation in waters whose pH is less than 4-5, exceeding even the concentrations of Ca2+ and Mg"~, which are the dominant cations at pH values greater than 4-5.

In the recent past, fears arose that human ingestion of aluminum from drinking water and from the use of aluminum cooking pots was a major cause of Alzheimer's disease; however, the research upon which this conclusion was reached could not be reproduced. Today many neuroscientists do not believe that there is a strong connection between the disease and intake of the metal, since past epidemiological studies on this matter have not been definitive or consistent. However, Canadian and Australian research reported in the mid-1990s indicates that consumption of drinking water with greater than 100 ppb aluminum—not an uncommon level in drinking water purified by aluminum sulfate (see Chapter 14)—can lead to neurological damage such as memory loss and perhaps to a small increase in the incidence of Alzheimer's disease.

It is thought that the principal deleterious effect of acid waters upon fish arises from the solubilization of aluminum from soil and its subsequent existence as a free ion in the acidic water, as discussed in Chapter 4. Unfortunately, the Al(OH)3 then precipitates as a gel on contact with the less acidic gills of the fish, and the gel prevents the normal intake of oxygen from water, thus suffocating the fish.

It is also believed that aluminum mobilization in soils is one of the stresses that acid rain places on trees, which in turn results in the dieback of forests. Soils that contain limestone are usually considered to be buffered against much change in pH, due to the ability of carbonate and bicarbonate ion to neutralize H+, but over a period of decades, surface soil may gradually lose its carbonate content due to a continual bombardment by acid rain. Thus soils receiving acid rain eventually become acidified. When the pH of the soil drops below about 4-2, aluminum leaching from soil and rocks becomes particularly appreciable. Such acidification has occurred in some regions of central Europe, including Poland, the former Czechoslovakia, and eastern Germany, and the resulting solubilization of aluminum may have contributed to the forest diebacks observed there in the 1980s.

PROBLEM 13-23

What is the concentration, in grams per liter, of dissolved aluminum in water having a pH of 5.5?

PROBLEM 13-24

Calculate the pH value at which the aluminum ion concentration dissolved in water is 0.020 M, assuming that it is controlled by the equilibrium with solid aluminum hydroxide.

Review Questions

1. Write the balanced half-reaction involving 02 when it oxidizes organic matter in acidic waters.

2. How does temperature affect the solubility of 02 in water? Explain what is meant by thermal pollution.

3. Define BOD and COD, and explain why their values for the same water sample can differ slightly. Explain why natural waters can have a high BOD.

4. What do the acronyms TOC and DOC stand for, and how do they differ in terms of what they measure?

5. Write the half-reaction, used in the COD titration, which converts dichromate ion to Cr3+ ion, and balance it.

6. Write the balanced chemical reaction by which organic carbon, represented as CH20, is dispropor-tionated by bacteria under anaerobic conditions.

7. Draw a labeled diagram classifying the top and bottom layers of a lake in summer as either oxidizing or reducing in character, and show the stable forms of carbon, sulfur, nitrogen, and iron in the two layers.

8. What are some examples of highly reduced and highly oxidized sulfur in environmentally important compounds? Write the balanced reaction by which sulfate can oxidize organic matter.

9. Explain the phenomenon of acid mine drainage, writing balanced chemical equations as appropriate. How does Fe3+ also act as an oxidizing agent here?

10. What is meant by the pE of an aqueous solution? What does a low (negative) pE value imply about the solution? What species determines the pE value in aerated water?

11. What are the acid and the base that dominate the chemistry of most natural water systems and whose interaction produces bicarbonate ion?

12. What is the source of most of the carbonate ion in natural waters? What name is given to waters that are exposed to this source?"

13. Write the approximate net reaction between carbonate ion and water in a system that is not also exposed to atmospheric carbon dioxide. Is the resulting water acidic, alkaline, or neutral?

14. Write the approximate net reaction between carbonate ion and water in a system that is exposed to atmospheric carbon dioxide. Is the resulting water mildly acidic or mildly alkaline? Explain why the production of bicarbonate ion from carbonate ion does not inhibit its production from carbon dioxide, and vice versa.

15. If two equilibrium reactions are added together, what is the relationship between the equilibrium constants for the individual reactions and that for the overall reaction?

16. What is the natural source of fluoride ion in water? How and why is the fluoride level in drinking water artificially increased to about 1 ppm in many municipalities?

17. Define the total alkalinity index and the phenolphthalein alkalinity index for water.

18. Define the hardness index for water.

19. Which are the most abundant ions in clean, fresh water?

20. Explain why aluminum ion concentrations in acidified waters are much greater than those in neutral water. How does the increased aluminum ion level affect fish and trees?

Green Chemistry Questions

See the discussion of focus areas and the principles of green chemistry in the Introduction before attempting these questions.

1, What takes place during the scouring of cotton, and why is this process necessary for the production of finished cotton fibers?

2. BioPreparation (an enzymatic process) replaced the use of large amounts of sodium hydroxide in the scouring of cotton.

(a) Describe any environmental problems or worker hazards that would be associated with the use of sodium hydroxide solutions in the scouring of cotton.

(b) Would these same environmental problems or worker hazards be eliminated by the use of BioPreparation?

3. The development of BioPreparation by Novozymes won a Presidential Green Chemistry Challenge Award.

(a) Into which of the three focus areas for these awards does this award best fit?

(b) List at least three of the twelve principles of green chemistry that are addressed by the green chemistry developed by Novozymes.

Additional Probl ems

See Table 13-3 for data.

1. Over a period of several days, estimate your approximate daily water usage in the categories of showering/bathing, clothes washing, toilets, dishwashing, and cooking. (Many flush toilets display volume-per-flush data. Washing machine water volume capacity can be estimated from the dimensions of the washer cavity: 1 L = 10 cm X 10 cm X 10 cm. Using a measuring cup, discover how long it takes your shower to deliver 1 liter of water, and adjust the data accordingly for the length of your average shower.)

2, The TOC parameter for water samples is mea sured by oxidizing the organic material to carbon dioxide and then measuring the amount of this gas evolved from the solution. If a 5.0-L sample of wastewater produced 0.25 mL of carbon dioxide gas, measured at a pressure of 0.96 atm and a tem perature of 22°C, calculate the TOC for the sam ple. Assuming the average composition of the organic matter to be CH20, calculate what the chemical oxygen demand for the water sample would be due to its organic content. [The gas con stant R = 0.082 L atm moP1 KT1.]

3. (a) Balance the reduction half-reaction that converts S04" to H2S under acidic conditions.

(b) Deduce the expressions relating pE to pH, the concentration of sulfate ion, and the partial pressure of hydrogen sulfide gas, given that for the half-reaction, pE° = —3.50 V when the pH is 7.0.

(c) Deduce the partial pressure of hydrogen sulfide when the sulfate ion concentration is 1Q^5 M and the pH is 6.0 for water that is in equilibrium with atmospheric oxygen.

4. Calculate the solubility of lead(II) carbonate, PbCOj (Ksp = 1.5 X 10"I3) in water, given that most of the carbonate ion it produces subsequently reacts with water to form bicarbonate ion. Recalculate the solubility assuming that none of the carbonate ion reacts to form bicarbonate.

Is your result significantly different from that calculated assuming complete reaction of carbonate with water?

5. The bicarbonate ion, 11CC)3 , can potentially act as an acid or as a base in water. Write the chemical equations for these two processes, and from the information given in this chapter, determine the corresponding acid and base dissociation constants. Given the relative magnitudes of the dissociation constants, decide whether the dominant reaction of bicarbonate in water will be as an acid or as a base. Calculate the pH of an aqueous

0.010.M solution of sodium bicarbonate in water using the dominant reaction alone and assuming that the amounts of carbonate ion and carbonic acid from other sources are negligible in this case.

6. A sample of lake water at 25 °C is analyzed and the following parameters are found:

total alkalinity = 6.2 x 10 4 M phenolphthalein alkalinity = 1.0 x 10 5 M pH = 7.6

hardness = 30.0 mg/L [Mg2+] = 1.0 X 10"4M

Extract all possible single-ion concentrations that you can by combining one or more of these data. Also determine whether or not the water is at equilibrium with respect to the carbonate—bicarbonate system and whether or not it is saturated with calcium carbonate.

Further Readings

1. T. Oki and S. Kanae, "Global Hydrological Cycles and World Water Resources," Science 313(2006): 1968.

2. W. Stumm and J.J. Morgan, Aquatic Chemistry; Chemical Equilibria and Rates in Natural Waters, 3rd ed. (New York: Wiley-Interscience, 1996).

3. A. Kousa et al., "Calcium: Magnesium Ratio in Local Groundwater and Incidence of Acute Myocardial Infarction Among Males in Rural Finland," Environmental Health Perspectives 114 (2006): 730.

Websites of Interest

7. The 02 concentration of a water sample can be determined using the so-called Winkler titration method. In it, the oxygen in a small sample of the water is reacted with MnSO^ in a basic solution. The reaction precipitates the manganese as MnOz, which converts added I" to I2. Molecular iodine is then quantitatively determined by titrating it in acidic solution with a standardized solution of sodium thiosulfate, Na2S203. The set of equations for the reactions is:

Mn02(s) + 4 H+ + 2 I"-» Mn24 + I2 + 2 H20

I2(aq) + 2 S2032"-> S406'" + 21"

In determining the BOD of a sample of water, a chemist tested two 10.00-mL samples of the water, one before and one after the five-day incubation period. They required 10.15 and 2.40 niL of a 0.00100 M standard solution of K2S203. Calculate the BOD, in units of milligrams per liter, of this water sample. On the basis of these results, would you consider this water to be polluted?

4. "What's in That Bottle?" Consumer Reports (January 2003): 38.

5. B. Hileman, "Fluoridation of Water," Chemical and Engineering News (August 1, 1998): 26.

6. F. M. M. Morel and J. G. Hering, Principles and Applications of Aquatic Chemistry (New York: Wiley, 1993).

7. G. Sposito, ed., The Environmental Chemistry of Aluminum, 2nd ed. (Boca Raton, FL: Lewis Publishers, 1996).

Log on to www.whfreeman.com/envchem4/ and click on Chapter 13.

Continue reading here: The Pollution And Purification Of Water

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  • Ugo
    What is the TOC of a 5L sample of wastewater which produced 0.25ml of CO2 gas at 0.96atm and 22℃?
    5 years ago
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    How to homemade iron can be converted to dichromate ion and vice versa?
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