Water chemistry and weathering regimes

Comparison of dissolved major ion compositions in four large rivers draining very different crustal areas (Table 5.2) shows the dominance of calcium (Ca), magnesium (Mg), sodium (Na) and potassium (K). Overall, however, the chemistry of each river is different and weathering regimes control most of these variations. The dissolved ion composition of freshwater depends upon:

1 the varying composition of rainfall and atmospheric dry deposition;

2 the modification of atmospheric inputs by evapotranspiration;

Table 5.2 Dissolved major ion composition (mmoll ') of some major rivers. Data from Meybeck (1979); except Rio Grande from Livingston (1963).

Mackenzie (1)

Orinoco (2)

Ganges (3)

Rio Grande (4)

Ca2+

0.82

0.08

0.61

2.72

Mg2+

0.43

0.04

0.20

0.99

Na+

0.30

0.06

0.21

5.10

K+

0.02

0.02

0.08

0.17

Cl4

0.25

0.08

0.09

4.82

SO44

0.38

0.03

0.09

2.48

hco3

1.82

0.18

1.72

3.00

SiO2

0.05

0.19

0.21

0.50

Drainage basin characteristics: (1) northern arctic Canada; (2) tropical northern South America; (3) southern Himalayas; (4) arid southwestern North America.

Drainage basin characteristics: (1) northern arctic Canada; (2) tropical northern South America; (3) southern Himalayas; (4) arid southwestern North America.

3 the varying inputs from weathering reactions and organic matter decomposition in soil and rocks;

4 differential uptake by biological processes in soils.

Where crystalline rock or highly weathered tropical soils are present (i.e. where weathering inputs are low or exhausted), dissolved freshwater chemistry is most influenced by natural atmospheric inputs, for example sea spray and dust, as well as by anthropogenic gases, for example SO2.

Sea-salt inputs are common in coastal regions. These salts have been introduced into the marine atmosphere from bubble bursting and breaking waves and are deposited on land with rain and dust fall. Small amounts of sea-salts are, however, also present in rainwater of central continental areas, thousands of miles from the sea. Sea-salt inputs have broadly similar, predominantly sodium chloride (NaCl), chemistry to the seawater from which they were derived. Thus, sodium or chloride ions can be used as a measure of sea-salt inputs to rainwater. Chloride concentrations in rain falling on oceanic islands are around 200 |mmoll-1, rain within 100 km of coastal continental areas contains around 10-100 |mmoll-1, while further inland chloride concentrations fall below 10 |mmoll-1, but not to zero.

The importance of a seawater source of ions other than sodium and chloride in rainwater can be assessed by computing their relative abundance with respect to sodium and comparing this with the same ratio in seawater. This comparison can be extended to freshwater, although here there is the complication that some of the ions could be derived from weathering. If we overlook this complication initially, then, where rainwater inputs make a large contribution to the chemistry of freshwater, the dominant cation is likely to be Na+. If weathering reactions are important, then the major dissolved cations will be those soluble elements derived from local rock and soil. In the absence of evaporite minerals, which are a minor component of continental crust (see Fig. 4.1), the most weatherable rocks are limestones (CaCO3). The calcium ion, liberated by limestone dissolution, is an indicator of this weathering process. This is clearly demonstrated by comparing the Ca2+ concentration in groundwater from a limestone aquifer, with groundwater from granites or metamorphic schists (Table 5.3).

Table 5.3 Chemical analyses of US groundwater from various rock types (mmoll '). Adapted from Todd (1980). This material is used by permission of John Wiley & Sons, Inc.

South Carolina

Metamorphic schist Georgia

Limestone Texas

Ca2+

0.3

0.7

1.8

Mg2+

0.2

0.2

0.4

Na+

0.4

0.7

1.0

K+

0.1

0.2

0.2

Cl-

0.1

0.1

0.7

SO2-

0.1

0.1

0.2

hco3

1.2

2.3

4.5

Al

0.004

0

Fe

0.003

0.002

0.001

SiO2

0.6

0.35

0.18

Na+/(Ca2++ Na+)

0.57

0.50

0.36

The ratio of Na+: (Na+ + Ca2+) can therefore be used to discriminate between rainwater and weathering sources in freshwaters. When sodium is the dominant cation (sea-salt contribution important), Na+: (Na+ + Ca2+) values approach 1. When calcium is the dominant ion (weathering contribution important), Na+: (Na+ + Ca2+) values approach 0.

The composition of dissolved ions in riverwater can be classified by comparing Na+: (Na+ + Ca2+) values with the total number of ions present in solution (Fig. 5.3). Note that the total dissolved ions or salts can also be expressed as the ionic strength of the water (Box 5.1). Data which plot in the bottom right of Fig. 5.3 represent rivers with low ion concentrations and sodium as the dominant cation. These rivers flow over crystalline bedrock (low weathering rates) or over extensively weathered, kaolinitic, tropical soils (low weathering potential, chemical index of alteration (CIA) c. 100 (see Table 4.9)). The Rio Negro, a tributary of the Amazon (Fig. 5.4), draining the highly weathered tropical soils of the central Amazonian region, has low ionic strength (Box 5.1) with weathering-derived sodium as the major cation. The Onyx River in the dry valleys of Antarctica is a better example of a low-ionic-strength, sea-salt-sodium-dominated river. This river has its source as glacial melt water and has a starting chemistry almost totally dominated by marine ions. As it flows over the igneous and metamorphic rocks of the valley floor, its composition evolves to higher ionic strength with an increasing proportion of calcium (Fig. 5.3).

Major river systems flow over a wide range of rock types, acquiring the dissolved products of weathering reactions. Freshwaters originating in areas with active weathering processes will have higher ion concentrations and an increasing predominance of calcium over sodium. These rivers plot along a trend from A to A' on Fig. 5.3. The Mackenzie and Ganges (Table 5.2) fall within this group, despite very different geomorphological settings.

The Amazon and its tributaries are a good example of a river system where the chemistry of the lower reaches integrates the products of differing soil and bedrock weathering regimes (Fig. 5.4). Rivers draining the intensely weathered soils and sediments of the central Amazonian region, such as the Rio Negro, have low total cation concentrations of less than 200 |meql-1 (i.e. sum of all major cations concentrations X charge; see also footnote to Table 4.10). The Rio Negro has water relatively enriched in sodium, silica, iron, aluminium and hydrogen ions, because of the limited supply of other cations from weathering reactions. By contrast, rivers draining easily erodible sedimentary rocks (including carbonates) of the Peruvian Andes are characterized by high total cation concentrations of 450-3000 |meql-1, including abundant calcium, magnesium, alkalinity (see below and Box 5.2) and sulphate. Between these two extremes in water composition are rivers with quite low total cation concentrations, with sodium enriched relative to calcium and magnesium, but also with high concentrations of silica, consistent with the weathering of feldspars (e.g. albite (see eqn. 4.14)). These rivers drain areas without large amounts of easily weatherable rock, but drain soils not so completely degraded as the lowest concentration group characterized by the Rio Negro.

In arid areas, evaporation may influence the major dissolved ion chemistry of rivers. Evaporation concentrates the total amount of ions in riverwater.

10 000

1000

Major oceans / .A

A Black

Major oceans / .A

A Black

Rio Grande '. Volga

* • • Congo ""Niger s ° ° Lake Superior o \

Rainfall dominance

Rio Grande '. Volga

* • • Congo ""Niger s ° ° Lake Superior o \

Rainfall dominance

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 Weight ratio of Na/(Na+Ca)

Fig. 5.3 Variation of the weight ratio of Na/(Na + Ca) as a function of the total dissolved salts and ionic strength for surface waters. Arrows represent chemical evolution of rivers from source downstream. Modified from Science 170, 1088-1090, Gibbs, R.J. Copyright (1970) by the AAAS.

Evaporation also causes CaCO3 to precipitate from water before NaCl, the latter being a more soluble salt. The formation of CaCO3 removes calcium ions from the water, increasing the Na+: (Na+ + Ca2+) value. Data for rivers influenced by evaporation plot along a diagonal B to B' on Fig. 5.3, evolving toward a sea-

Palaeozoic and older metamorphic rock ' y'' y' _ with highly weathered soils

■ ■ ■ ■

Rocks in which evaporite minerals (NaCl and CaSO4) occur

Unshaded areas mainly Cenozoic sedimentary rock

Fig. 5.4 Simplified geological map of the Amazon Basin showing major river systems. After Stallard and Edmond (1983), copyright (1983) American Geophysical Union modified by permission of American Geophysical Union.

Fig. 5.4 Simplified geological map of the Amazon Basin showing major river systems. After Stallard and Edmond (1983), copyright (1983) American Geophysical Union modified by permission of American Geophysical Union.

water composition in the upper right (see also Table 6.1). The Colarado River in the arid southwest of the USA is an interesting example. Construction of the Hoover Dam on the Colorado created Lake Mead upstream of the dam, increasing evaporation from the catchment. The concentration of dissolved solids increased by 10-20% in outflowing water below the dam, while CaCO3 precipitation occurred in Lake Mead. A more extreme example is seen in the Salton Sea

Box 5.1 Ionic strength

The concentration of an electrolyte solution can be expressed as ionic strength (/), defined as:

where c is the concentration of ion (i) in mol l-1, z is the charge of ion (i) and S represents the sum of all ions in the solution.

As a measure of the concentration of a complex electrolyte solution, ionic strength is better than the simple sum of molar concentrations, as it accounts for the effect of charge of multivalent ions.

For example, the Onyx River in Antarctica (Section 5.3), 2.5 km below its glacier source, has the following major ion composition (in mmoll-1): Ca2+ = 55.4; Mg2+ = 44.4; Na+ = 125; K+ = 17.6; H+= 10-3; Cl- = 129; SO42- = 32.2, HCO-= 136 and OH- = 10 (rivers in this region have pH around 9, which means that HCO3- is the dominant carbonate species). Putting these ions in equation 1 gives:

/ = X ScCa.4 + cMg.4 + cNa.1 + cK.1 +cH.1 + cCl.1 + CSO4.4

Substituting the mmoll-1 values (and correcting to moll-1 with the final 10-6 term in eqn. 3) gives:

/ = X[(55.4 ¥ 4) +(44.4 x 4) +125 +17.6 +10-3 +129 + (32.2 x 4) +136 +10] x 10-6

Freshwaters typically have ionic strengths between 10-3 and 10-4moll-1, whereas seawater has a fairly constant ionic strength of 0.7 mol l-1.

Box 5.2 Measuring alkalinity

Alkalinity is measured by adding acid to a water sample until the pH falls to 4. At this pH, HCO3- and CO32- alkalinity (known as the carbonate alkalinity (Ac)) will have been converted to carbon dioxide (CO2), i.e.:

Note that twice as much H+ is used up neutralizing the CO|- (eqn. 2) relative to the HCO3- (eqn. 1). This is expressed in the formulation that expresses carbonate alkalinity, by the 2 in front of the cCO|-term, i.e.:

This formulation is expressed in concentration because the carbonate species are measured values (see Section 2.6). The volume of acid used is a measure of the alkalinity, which is usually expressed as milliequivalents per litre (see footnote to Table 4.10). These units account for the difference in H+ neutralizing power between COf- and HCO3-. Note, however, that at pH around 7.5-8, monovalent HCO3- accounts for almost all alkalinity (Fig. 5.5) such that at these pH values milliequivalents are essentially the same as millimoles.

on the border between the USA and Mexico. This very large lake was created in 1905 by floodwaters of the Colorado, forming a closed basin lake with no outflow to the sea. The Salton Sea began as a freshwater lake with total dissolved salt concentrations around 3.5 g l-1 in 1907. Evaporation since this time means that the lakewaters are now saltier than seawater. In 1997, total dissolved salt concentra-

tions stood at 47 gl-1. Rivers draining into the Salton Sea have a Na+/Ca2+ ratio of about 5 : 1 (on an atomic basis) but as a result of evaporation and CaCO3 precipitation, this ratio increases to 27 :1 in the Salton Sea itself. Other examples of rivers in which evaporation plays an important role include the Jordan and Rio Grande. In all these arid areas, dissolution of evaporite minerals in the catchments may also contribute to the increasing dominance of NaCl.

The classification of riverwater composition in Fig. 5.3 is simplified and does not always work. For example, weathering of feldspars (see Section 4.4.4) can produce solutions of low ionic strength, but rich in sodium and silica, which plot in the bottom right of Fig. 5.3. This effect probably influences the classification of the Rio Negro. Weathering of evaporite minerals will also affect the composition of rivers. For example, in the Amazon catchment there are a small number of tributaries draining areas of predominantly evaporite rock. These have very high total cation concentrations and are characterized by high sodium, chloride, calcium and sulphate concentrations from the weathering of the evaporite minerals, halite and gypsum. Despite these complications, Fig. 5.3 remains a useful way to compare factors controlling riverwater chemistry. Indeed, it is remarkable that most of the world's major rivers can be rationalized in this straightforward way.

5.3.1 Alkalinity, dissolved inorganic carbon and pH buffering

In Section 4.4 we saw that most soilwaters that feed rivers and groundwater have near-neutral pH, with HCO3- as the major anion. This results from the dissolution of CO2 in water (see eqn. 4.7) and from the acid hydrolysis of silicates and carbonates. The total concentration of weak acid anions like HCO3- in water is referred to as alkalinity. These anions are available to neutralize acidity (H+) in natural waters, consequently it is important to understand their chemical behaviour.

In continental waters, bicarbonate (HCO3-) and carbonate (CO32-) ions are the most important components of alkalinity, although in seawater other ions also contribute to alkalinity. The relative importance of HCO3- and CO32- depends on the pH of the solution and can be calculated from the known dissociation constants (see Box 4.5) of these ions and the solution pH.

The first dissociation of dissolved carbon dioxide (expressed here as carbonic acid),

has a dissociation constant, aH aHCO-

aH2CO3

The '«' denotes activity, the formal thermodynamic representation of concentration (see Section 2.6).

Similarly, for the second dissociation of carbonic acid,

the dissociation constant is, aH +. aCO3-

2 aHCO-

It is simple to demonstrate that the alkalinity in most continental waters is dominated by HCO3- by rearranging the above equations at a typical pH value for these waters. For example, many mature rivers have pH values around 8. Rearranging equation 5.9 to solve for aHCO- gives, aH +. aCO2-

And substituting values for pH 8 (pH = -log10aH+, see Box 3.5) and K2 (eqn. 5.9) gives,

This shows that at typical pH values for continental waters the HCO3- anion is 200 times more abundant than the CO2- anion. Repeating this exercise for a range of pH values results in the graphical relationship shown in Fig. 5.5. Note that when pH falls below 5 on Fig. 5.5, almost all of the weak acid anions (HCO- and CO32-) have disappeared and at pH of 4 only undissociated acid (H2CO3) remains. This relationship is used as the basis for measuring alkalinity (Box 5.2).

The pH of natural waters is in fact controlled mainly by the relative concentrations of dissolved inorganic carbon (DIC) species, i.e. H2CO3, HCO3- and

Fig. 5.5 Relationship between the total dissolved inorganic carbon species (TDIC; i.e. H2CO3 + HCO- + COl-) and pH. Most natural waters have pH between 7 and 9 where the HCO3-anion is abundant (>80% of the TDIC). In highly alkaline waters (pH > 10.3) the COf- anion becomes dominant, while in acidic waters (pH <6.4) the undissociated acid (H2CO3) is the dominant TDIC species.

Fig. 5.5 Relationship between the total dissolved inorganic carbon species (TDIC; i.e. H2CO3 + HCO- + COl-) and pH. Most natural waters have pH between 7 and 9 where the HCO3-anion is abundant (>80% of the TDIC). In highly alkaline waters (pH > 10.3) the COf- anion becomes dominant, while in acidic waters (pH <6.4) the undissociated acid (H2CO3) is the dominant TDIC species.

CO2- (Fig. 5.5). These species react to maintain the pH within relatively narrow limits. This is known as buffering the pH and the principles, using worked examples, are demonstrated in Box. 5.3. The relationship between pH, CaCO3 weathering and alkalinity is nicely illustrated by real data (Fig. 5.6) from a small stream in North Yorkshire (UK). The stream begins as drainage from an organic peat bog. The initial pH is about 4 and hence the alkalinity is zero (Fig. 5.6). The stream then flows from the bog over siliceous mudrocks with very limited potential for weathering such that the water chemistry changes little. After this the stream flows onto limestone, the acidic water reacting rapidly with the CaCO3 to release Ca2+ and HCO- ions:

CaCO3(S) + H2CO3( aq) ^ Ca2+q) + 2hco3(aq) eqn. 5.12

In less than 100 m of distance flowing on the limestone the pH rises sharply to values of 6 or 7 and it continues to increase to a value of about 8 at which it stabilizes due to the buffering action of the alkalinity (Fig. 5.6). The alkalinity is strongly buffering the system at values above 2 meq l-1 where the relationship between pH and alkalinity is asymptotic (Fig. 5.6). The relationship between Ca2+

Fig. 5.6 Relationship between pH, alkalinity and dissolved calcium ions in stream waters in the Malham Tarn area of northern England, flowing from bog on siliceous mudrock to limestone. Note that pH is buffered around 8 once limestone weathering begins. Data from Woof and Jackson (1988).

Box 5.3 Worked examples of pH buffering

The principle of pH buffering can be illustrated by considering the simple case of acetic acid, CH3COOH (abbreviated here to HA) and sodium acetate, CH3COONa (abbreviated here to NaA). Acetic acid partially dissociates in water (H2O) while the sodium salt completely dissociates.

Rearranging gives:

For a 0.1 molar (m) solution of HA and NaA (for simplicity we will assume that activity (a) and concentration (c) are the same) and assuming very little HA dissociates, then:

We know that pH = -log10 aH+ (see Box 3.5), so in our example:

To illustrate the principle of buffering, consider what happens when 0.005moles of NaOH (sodium hydroxide, a strong base) are added to 1 litre of 0.1 M HA and NaA. The added base reacts with the hydrogen ions (H+), causing an amount of the HA equivalent to the added NaOH to dissociate (eqn. 1); the HA concentration decreases and the A- concentration increases by this amount.

aH+ = 10-475 x (°.1-°.°05) = 1.61 x 10-5moll-i (0.1 + 0.005)

Now pH becomes:

added OH-. Buffering will continue if an excess of HA is available. If acid is added to the solution, H+ will react with the excess of A- to increase the HA and decrease the A-concentration by an amount equivalent to the added H+, resulting in a similar buffering effect. The HA and NaA solution is an effective buffer because it can react to neutralize either added acid or base.

By contrast, if 0.005moles of NaOH are added to 1 litre of water, the pH will rise to 11.7, as illustrated below.

Thus:

eqn. 10

So, if 0.005 moles of OH- are added to 1 litre of water (again assuming that activity and concentration are the same), then:

In natural waters the buffering system involves the weak acid, carbonic acid (H2CO3), and the associated anions, bicarbonate (HCO-) and carbonate (CO32-). At pH 4-9, HCO- is the major anion. In the following example we ignore CO3- (and again assume that activity and concentration are the same). First, we can rewrite equation 4 for the HCO-system:

eqn. 13

and define the terms used (eqns 14-17). At 25°C the equilibrium constant for equation 1 in Box 5.2 is defined as:

eqn. 14

(see also Box 3.7), whilst the relationship between partial pressure of carbon dioxide (pCO2) and H2CO3 is:

The pH is barely altered because the excess of undissociated HA dissociates to neutralize the

hH2O,

eqn. 15

(continued)

3H2CO3 = 10-14 x pCO2 Now equation 13 can be rewritten:

3H2CO3 = 10-14 x pCO2 Now equation 13 can be rewritten:

eqn. 16

eqn. 20

Consider the case of the Mackenzie River, where HCO3-= 1.8mmoll-1 (Table 5.2) and atmospheric pCO2 = 3.6 x 10-4atm:

Although this treatment is simplified, eqn. 18 it serves to illustrate the way in which

Consider the case of the Mackenzie River, where HCO3-= 1.8mmoll-1 (Table 5.2) and atmospheric pCO2 = 3.6 x 10-4atm:

pH can be calculated. In practice the pH of most natural water containing HCO3

nd and CO32- is buffered between pH 7

and alkalinity is, in this example, approximately linear, although for every mmol of Ca2+, 2mmol of HCO- are released, as predicted in equation 5.12.

0 0

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