modern hydrology and sustainable water development (gupta/modern hydrology and sustainable water...

P1: SFK/UKS P2: SFK

c06 BLBK276-Gupta August 28, 2010 15:23 Trim: 246mm X 189mm Printer Name: Yet to Come

6Aqueous chemistry and

human impacts on waterquality

Why is water chemistry, in general, and ground-water chemistry, in particular, important? Waterquality is an important aspect for all uses of wa-ter, be it for drinking, irrigation, industrial, or otherpurposes. Some of the quality considerations werediscussed in Section 1.2. Aqueous chemistry is cen-tral to understanding: (i) sources of chemical con-stituents in water; (ii) important natural chemicalprocesses occurring in groundwater; and (iii) vari-ations in chemical composition of groundwater inspace and time. Aqueous chemistry is also impor-tant to estimate the fate of contaminants, both insurface waters and groundwaters, and for remedia-tion of contamination.

As mentioned in Chapter 1, water is called the‘universal solvent’ because it dissolves more sub-stances than any other liquid. This is due to the po-lar nature of the water molecule, which makes it agood solvent of ionic as well as polar molecules. Asa result, chemical composition of natural water isderived from multiple sources of solutes, includinggases and aerosols from the atmosphere, weather-ing and erosion of rocks and soil, solution or precip-itation reactions occurring below the land surface,and those resulting from human activities. Some ofthe common inorganic solutes in water are listedin Table 6.1. In this chapter, the focus is on under-standing the broad interrelationships between thevarious processes and their effects that make up thechemical character of natural waters. These include

principles of chemical thermodynamics, processesof dissolution and/or precipitation of minerals, andthe laws of chemical equilibrium including the lawof mass action, etc.

Basic data used in the determination of waterquality are obtained by the chemical analysis ofwater samples in the laboratory or onsite measure-ment/sensing in the field. Most of the measuredconstituents are reported in gravimetric units, inmilligrams per litre (mg/L or mg l−1) or milli-equivalents per litre (meq/L or meq l−1). However,in chemical calculations, it is common to use mo-lar concentration units, which give moles of soluteper unit volume of solution (mol/L or mol l−1), de-noted by M (e.g. a 2.5 M Ca2+ solution contains 2.5moles of Ca2+ per litre. A mole is the amount ofsubstance containing N atoms or molecules, whereN = 6.022 × 1023 is the Avogadro’s number,rounded off to 4 significant digits. The mass/weightof one mole of atoms is called the formulamass/weight (also called molecular mass/weight).For example, the atomic mass of oxygen is 16.00g, and the formula mass of CO2 is (12.01 + 2 ×16.00) = 44.01 g. For calculations involving chemi-cal reactions, it is useful to employ moles and molarconcentrations, because chemicals react in directproportion to the number of respective moleculespresent. Conversion from molar mass per volumeto units (i.e. mg l−1 or µg l−1) can be done by usingthe formula weight of the solute below:

Modern Hydrology and Sustainable Water Development S. K. Gupta

© 2011 S. K. Gupta. ISBN: 978-1-405-17124-3

P1: SFK/UKS P2: SFK


136 MODERN HYDROLOGY AND SUSTAINABLE WATER DEVELOPMENT

Table 6.1 Some common inorganic solutes in water.

Cations Anions Other

Major ConstituentsCalcium (Ca2+) Bicarbonate

(HCO3−)

Dissolved CO2

(H2CO3)

Magnesium(Mg2+)

Chloride (Cl−) Silica (SiO2(aq))

Sodium (Na+) Sulphate (SO42−)

Potassium (K+)

Minor ConstituentsIron (Fe2+, Fe3+) Carbonate (CO3

2−) Boron (B)

Strontium (Sr2+) Fluoride (F−)Nitrate (NO3

−)

mol

L× formula weight

( g

mol

)× 1000 mg

g= mg

L

(6.1)

Chemical analyses may be grouped and statisticallyevaluated by parameters such as mean, median, fre-quency distribution, or ion correlations to charac-terize large volumes of data. Drawing graphs ofanalyses, or of groups of analyses, aids in showingchemical relationships between various dissolvedconstituents of water and identifying their prob-able sources, regional water-quality relationships,temporal and spatial variation of water quality, andassessment of water resources potential. Graphsmay show water types based on chemical composi-tion, relationships between ions (or groups of ions)in individual waters, or water from many sourcesconsidered simultaneously. Relationship betweenwater quality and hydrogeological characteristics,such as the stream discharge rate or groundwaterflow patterns, can be shown by appropriate math-ematical equations, graphs, and maps.

6.1 Principles and processes controllingcomposition of natural waters

The fundamental concepts relating to chemical pro-cesses, which are useful in developing a unifiedapproach to understand the chemistry of naturalwaters, are mainly related to chemical thermody-

namics as well as to rates of reactions and theircausative mechanisms.

Thermodynamic principles, applied to chemicalenergy transfers, form a basis for quantitatively eval-uating the feasibility of various possible chemicalprocesses occurring in natural water systems, forpredicting the direction in which chemical reac-tions may proceed, and in many cases for predict-ing the actual dissolved concentrations of reactionproducts that may be present in a given water body.

6.1.1 Thermodynamics of aqueous systems

A review of fundamental relationships and princi-ples relating to chemical energy is helpful in un-derstanding how thermodynamic concepts maybe used.

The law of conservation of energy states that al-though its form may change, the total amount of en-ergy in any system remains constant. This principleis also known as the first law of thermodynamics.

The chemical energy stored in a substance atconstant temperature and pressure is termed ‘en-thalpy’ and is represented by �H. The ‘delta’ indi-cates that it represents a departure from an arbitrarystandard state taken as zero point. For chemical el-ements, this standard reference state is that of 1mole (an amount equal to the molecular weight ofthe constituent expressed in grams) at 25◦C tem-perature and 1 atmosphere pressure. Enthalpy maybe thought of as having two components: an inter-nal component that is termed ‘entropy’ (�S) anda component that is or can become available ex-ternally, termed ‘free energy’ (�G) or Gibbs freeenergy.

The concept of entropy is implicit in the secondlaw of thermodynamics, which can be stated as aspontaneously occurring process in an isolated sys-tem that will tend to convert a less probable stateto one or more probable states. There is a finiteprobability that such a system tends to favour agenerally random or disordered condition, or even-tually a state of relative chaos. Entropy may thusbe considered a measure of the degree of disor-ganization or disorder within a system. However,entropy is more difficult to estimate quantitativelythan its corollary, free energy, which is released ina spontaneous chemical process.

P1: SFK/UKS P2: SFK


AQUEOUS CHEMISTRY AND HUMAN IMPACTS ON WATER QUALITY 137

The second law of thermodynamics can also berephrased to the effect that in a closed system thereaction affinities tend to reach their minimumvalues. At the point of equilibrium for a specificreaction, the value of reaction affinity is zero.The relationship governing these chemical energymanifestations is:

�H = �G + T �S (6.2)

where T is temperature on the Kelvin scale. Thisis a general statement of the third law of thermo-dynamics, which also may be paraphrased ‘the en-tropy of a substance at absolute zero temperature(T = 0K) is zero’.

Enthalpy, entropy, and free energy values are ex-pressed in units of heat.

6.1.2 Chemical reactions

Chemical reactions in which various elements par-ticipate involve changes in the arrangement and as-sociation of atoms and molecules and interactionsamongst electrons that surround the atomic nuclei.The field of natural water chemistry is concernedprincipally with reactions that occur in relatively di-lute aqueous solutions, although some natural wa-ters have rather high solute concentrations. Thereacting systems of interest are generally heteroge-neous, that is, they involve either a liquid phase,a solid or a gaseous phase, or all the three phasescoexist together.

6.1.2.1 Reversible and irreversible reactionsin water chemistry

In the strict sense, an irreversible process is onein which reactants are completely converted intoproducts. Thus a reversible process is one in whichboth reactants as well as products can be presentwhen the reaction affinity is zero or nearly zero. Itis inferred that to achieve and sustain this condi-tion, both the forward and reverse reactions occursimultaneously, at least on the micro-scale and atcomparable rates when reaction affinities are small.

Some reactions, even though favoured thermo-dynamically, do not take place to a significantextent owing to energy barriers in some of thereaction pathways. If such a condition applies

to one of the reactions in a reversible process asdefined above, the process apparently behaves asirreversible. Similarly, in open systems in whichreactants and/or products may enter and leave,irreversible behaviour may be expected. There-fore, in natural water systems, the reversible orirreversible nature of a chemical reaction is depen-dent on kinetic factors and on some of the physicalfeatures of the system of interest as well as on ther-modynamic considerations. Therefore, it has beenfound more convenient in natural-water systems toconsider chemical reactions of interest on the basisof the ease with which these can be reversed. Onthis basis, three general types of processes are: (1)readily reversible processes; (2) processes whosereversibility is retarded; and (3) processes thatare irreversible in a fundamental (thermodynamic)sense. Specific processes in natural-water systemsrepresent a continuum from type 1 to type 3.An example of an easily reversible process is theformation of complex ions or similar homogeneous(single-phase) reactions. Dissolved carbon dioxide,represented as H2CO3, dissociates reversibly as:

H2CO3(aq) ↔ HCO−3 + H+ ↔ CO2−

3 + 2H+

(6.3)

The symbol (aq) indicates the aqueous phase.Other symbols commonly used to indicate phaseof reactants or products on chemical reactions are(s) for solid, (l) for liquid, and (g) for gas. An ex-ample of reactions in which reversibility is severelyaffected can be seen in the weathering of albite, acommon feldspar mineral, in which a solid productkaolinite is formed:

2Na Al Si3O8 + 2H+ + 9 H2O → 2Na+

+ 4 H2SiO4 (aq) + Al2Si2O5 (OH)4(s) (6.4)

Kaolinite dissolves reversibly as:

Al2Si2O5 (OH)4(s) + 7H2O ↔ 2 Al(OH)4

+ 2Si (OH)4(aq) + 2H+ (6.5)

Here dissolution of kaolinite is essentially irre-versible because it cannot be reconstituted to asignificant extent without subjecting it to temper-atures and pressures that differ greatly from thoseprevailing in normal weathering regimes.

P1: SFK/UKS P2: SFK



A process that is thermodynamically irreversibleis that of altering the crystal structure of a solidto a more stable form during its aging. For exam-ple, when a ferric hydroxide amorphous precipi-tate changes to goethite with aging:

Fe(OH)3(s) → FeOOH(c) + H2O (6.6)

the reaction affinity will be greater than zero aslong as any Fe(OH)3 is available.

6.1.3 Chemical equilibrium – the law ofmass action

Study of chemical equilibria is based on the lawof mass action, according to which the rateof a chemical reaction is proportional to theactive masses of the participating substances. Ahypothetical reaction between two substances Aand B producing products C and D, in a closedsystem, can be written as:

aA + bB ↔ cC + dD (6.7)

where the lower case letters represent coefficientsrequired to balance the equation. The rates of for-ward and reverse reaction, according to the law ofmass action are, respectively:

R1 = k′1[A]a[B]b (6.8)

and

R2 = k′2[C ]c[D]d (6.9)

where bracketed terms represent active massesor activities of reactants/products. The quantitiesk′

1 and k′2 are proportionality constants for the

forward and reverse reactions, respectively. WhenR1 = R2, the system is in a state of dynamic equi-librium and no change in active concentrations(represented by quantities in the square brackets)occurs. This leads to the equation:

[C ]c[D]d

[A]a[B]b= k′

1

k′2

= K (at equilibrium) (6.10)

The quantity, K, is referred to as the equilibriumconstant. Activities are dimensionless quantitiesand, therefore, the equilibrium constant is alsodimensionless. It has a characteristic value forany given set of reactants and products and manyexperimentally determined values are available in

published chemical literature. The value of theequilibrium constant is influenced by temperatureand pressure. Standard thermodynamic conditions(25◦C temperature and 1 atmosphere pressure) areusually specified, but K values for many reactionshave been determined over a wide temperaturerange. This form of the mass law (Eqn 6.10) isa statement of final conditions in a system atequilibrium. If a reaction is not in equilibrium, theright-hand side of Eqn 6.10 is called the reactionquotient or the ion activity product (IAP):

Q = IAP = [C ]c[D]d

[A]a[B]b(6.11)

When IAP <K, the reaction proceeds to the rightand the concentrations of A and B fall while thoseof C and D rise; when IAP >K, the reaction pro-ceeds to the left and concentrations change in theopposite direction.

Example 6.1. The following reaction describesdissociation of bicarbonate ion in water (see alsoEqn 6.3):

HCO−3 ↔ CO2−

3 + H+

The equilibrium constant for this reaction at 15◦Cis 10−10.43, i.e.:

KHCO−3

= [H+][CO2−

3

][HCO−

3

] = 10−10.43

If the water has a pH of 5.9 (i.e. [H+] = 10−5.9 (seeSection 1.2.2.5 in Chapter 1) and

[HCO−

3

] = 2.43 ×10−3, assuming equilibrium, the

[CO2−

3

]is given by:

[CO2−

3

] = KHCO−3

[HCO−

3

]

[H+]

=[10−10.43

] [2.43 × 10−3

][10−5.9

] = 7.2 × 10−8

The equilibration equations such as Eqn 6.10 can bewritten for reactions involving a variety of phasesassociated with groundwater: solute-solute, solute-water, solute-solid, and solute-sorbed. By conven-tion activities of water, solid phases or non-aqueousliquid phases in contact with water are set equal to1.0 in the equilibrium equation. This is becausethere is essentially an unlimited supply of thesesubstances in contact with reactants/products and

P1: SFK/UKS P2: SFK



equilibrium concentrations are independent of theamount of these substances. For example, dissolu-tion reaction of the mineral calcite (calcium car-bonate) in groundwater is:

CaCO3(s) = Ca2+ + CO2−3 (6.12)

The corresponding equilibrium equation is writtenas:

KCaCO3 =[Ca2+] [

CO2−3

][CaCO3

] = [Ca2+] [

CO2−3

]

(6.13)

Thus, activity of solid phases such as calcite isusually omitted from the equilibrium equation, asin Eqn 6.13.

In dilute solutions with low ion concentrations,the electrostatic repulsions between differentions and their ability to collide and react witheach other are not drastically hampered, butthis is not the case with water having a highion concentration. In the latter case, it becomesnecessary to use a correction factor known as theactivity coefficient to get the effective activity foruse in the equilibrium equation (Eqn 6.10). For achemical, the activity [D], concentration (D), andthe activity coefficient, γ D, are related as:

[D] = γD(D) (6.14)

In Eqn 6.14, activity [D] is dimensionless and con-centration (D) is in molar (M, mol l−1) units. Fordilute solutions γ ≈ 1 in magnitude, and the ac-tivity is essentially equal to that corresponding tothe magnitude of the molar concentration. In moreconcentrated solutions, such as γ �= 1, the activitydeviates from that corresponding to the magnitudeof the molar concentration.

Several mathematical models for activity coef-ficients have been developed based on electro-static theory and empirical observations. The mostwidely used model for ions in low- to moderately-concentrated- solutions is the extended Debye-Huckle equation:

log γi = −A Z2i

√I

1 + B a√

I(I < 0.1) (6.15)

where A and B are constants for a given solventthat depend on pressure and temperature and are

Table 6.2 Values of constants in the Debye-Huckle equationfor activity coefficients A and B (Eqn 6.15). After Manovet al. (1943).

T (◦C) A B

0 0.4883 0.324110 0.4960 0.325820 0.5042 0.327225 0.5085 0.328130 0.5130 0.329040 0.5221 0.3302

given in Table 6.2 for water. Zi is the charge of theion under consideration; I is the ionic strength ofthe solution in moles; and a is related to the sizeof the hydrated ion. Its values for various ions ofinterest are given in Table 6.3.

For brackish water with ionic strength I >0.1M,the Davies equation (Eqn 6.16) is a better approxi-mation of an ion’s activity coefficient (Stumm andMorgan 1996):

log γi = −A Z2i

( √I

1 + √I

− 0.3I

)(I <0.5)

(6.16)

Table 6.3 Values of the parameter ‘a’ in the Debye-Huckleequation (Eqn 6.15). After Kielland (1937) and Butler (1998).

Ions Charge a

Th4+, Sn4+ 4 11

Fe(CN )4−6 4 5

Al3+, Fe3+, Cr3+ 3 9

PO3−4 3 4

Mg2+, Be2+ 2 8

Ca2+, Cu2+, Zn2+, Sn2+, Mn2+, Fe2+,Ni2+, CO2+

2 6

Sr2+, Ba2+, Cd2+, Hg2+, S2−, CO2−3 ,

SO2−3 , MoO2−

4 , Pb2+, Ra2+2 5

CrO2−4 , HPO2−

4 , SO2−4 , Hg2+

2 , SeO2−4 2 4

H+ 1 9

Li+ 1 6

HCO−3 , H2PO−

4 ,HSO−3 , Na+ 1 4

Cl−, Br−, I−, F−, OH−, HS−, NO−3 ,

NH+4 , K−, Ag+, CNS, CNO−, ClO−

4 ,NO−

2 , Rb+, Cs+, CN−

1 3

P1: SFK/UKS P2: SFK



10-5 10-4 10-3 10-2 10-10

0.1

0.2

0.3

0.1

0.5

0.6

0.7

0.8

0.9

1.0

Act

ivit

y C

oef

fici

ent

Ionic Strength (M)

Charge-4

Cahrge-3

Charge-2

Charge-1

Fig. 6.1 Relationship between activity coefficients (γ ) for dissolved ions and ionic strength (I), in molar units, of a solution atone atmospheric pressure and at 25◦C temperature. Redrawn from Hem (1985). © U.S. Geological Survey.

where A is the same constant that appears in theprevious equation.

Activity coefficients versus ionic strength curves,obtained using the Debye-Huckle and Daviesmodel, are shown in Fig. 6.1. In general, the ac-tivity coefficient decreases with increasing ionicstrength and is smaller for ions with higher charge.

6.1.4 Reaction rates and deviation fromthe equilibrium

A chemical reaction can occur spontaneously ina closed system when the total free energy ofreactants exceeds that of the reaction products.The chemical equation representing such a pro-cess and standard free energy data for participatingspecies can be used to determine if the reactioncan be spontaneous. However, the knowledge thata given reaction is thermodynamically favourablegives only a limited amount of information that canbe used to predict how fast the reaction will pro-ceed. In fact, many reactions that are feasible donot occur at significant rates and some considera-tion of reaction rate theory is, therefore, necessaryto study natural-water chemistry.

In an irreversible chemical process:

αA → products (6.17)

The rate of reaction of A can be written as:

−d[A]

dt= k[A]α (6.18)

The rate constant, k, and the exponent, α, are pro-portionality factors that must be determined ex-perimentally. When the value of the exponent, α,is unity, the process is termed ‘first order’ in A, im-plying that the rate of the process is a function onlyof the concentration of A. It is implicitly assumedthat temperature and pressure remain constant dur-ing the process and that the ions or molecules ofA do not interact with each other. It is possible fora process to be of second order in A (for α = 2),implying that two ions of A must interact to createa product. Another type of second-order process isone involving two reactants:

αA + βB → products (6.19)

The rate of such a reaction, where α and β are bothunity, would be:

−d[A]

dt= −d[B]

dt= k[A][B] (6.20)

Higher-order reactions occur when α and/or β inthe above schematic reaction is greater than 1.

P1: SFK/UKS P2: SFK



Zero-order reactions occur when the rate is inde-pendent of concentration of the reactant, i.e.:

d[A]/dt = k

Such reactions are of considerable interest in sometypes of natural water systems and might occurwhen concentration of A is much smaller comparedto that of another reacting substance. Similarly, forprocesses whose rate is controlled by availabilityof reaction sites on a solid surface, zero-order ki-netics will be applicable if the number of reactionsites is large compared to the concentration of A.Reactions of fractional orders can also be observedfor some processes. Generally these involve com-binations of several reactions or effects, such asdiffusion or mixing, that are not entirely chemicalin nature. In evaluating the kinetic properties ofany chemical process it is important to considerthe effects of processes like these. Generally, onlyone step in the process controls the rate and de-termines the order of the whole reaction. This isknown as the rate-determining step.

Integration of the first-order rate equation leadsto:

[At] = [A0] e−kt (6.21)

where [A0] is the concentration of the reactant A att = 0 and [At] is the concentration at time t. Notethat Eqn 6.21 is similar to the radioactive decayequation (Eqn. 7.1). Therefore, one can define thehalf-life of the reaction as the time required forhalf the amount of A present at any moment todisappear:

t1/2 = ln 2

k= 0.693

k(6.22)

Reaction rates vary widely; some reactions are sorapid that they can result in runaway situationsleading to explosions, while others are so slowthat it requires geological times to measure theirkinetics. A general range of half-lives for severalcommon types of aqueous reactions is shown inTable 6.4.

Rates of some reactions are limited by factorsother than their corresponding theoretical chem-ical reaction rate. Some sorption-desorption re-actions are limited by molecular diffusion thattransports solute to sorption sites that are not in

Table 6.4 Approximate range of reaction half-lives for somecommon types of aqueous reactions. After Langmuir andMahoney (1984).

Type of Reaction Typical half-life

Solute-solute Fraction of seconds to minutesSorption-desorption Fraction of seconds to daysGas-solute Minutes to daysMineral-solute Hours to millions of years

direct contact with flowing water. Other reactions,including many redox reactions, involve micro-organisms and are, therefore, affected by the popu-lation density of micro-organisms as well as concen-tration of various nutrients used by organisms.

The use of equilibrium calculations is valid only ifa chemical system remains essentially closed longenough for the reaction of interest to approachequilibrium. In this sense, it is reasonable to applyequilibrium calculations for many solute-solute re-actions (Table 6.4). Disequilibrium prevails whentransport and other agents of change are rapid com-pared to the reaction rate, as for example in somemineral-solute reactions. Even in such cases, equi-librium calculations may still be useful to showwhere the system is heading in the long run.

A strictly chemical parameter for evaluatingdeparture from equilibrium is the saturation index,SI. This is the difference between the logarithmsof the activity quotient, Q, and the equilibriumconstant, K:

S I = log Q − log K = log(Q/K )) (6.23)

where Q is the reaction quotient or ionic activityproduct, defined by Eqn 6.11, obtained by usingobserved activities of participating substances inan actual system; K being defined by Eqn 6.10 forequilibrium situation.

6.1.5 Mineral dissolution and precipitation

For a mineral–water system, a positive value of SIindicates that the solution is supersaturated and thereaction should proceed in the direction that willcause more solute to precipitate out of the solution.A negative value of SI indicates under-saturation of

P1: SFK/UKS P2: SFK



Table 6.5 Reactions and solubility products for some common minerals for standard conditions (25◦C temperature; 1atmosphere pressure). After Morel and Hering (1993); Nordstrom et al. (1990).

Mineral Reaction Log(Kso) Ref.

Salts:Halite NaCl = Na+ + Cl− 1.54 (1)Sylvite K Cl = K + + Cl− 0.98 (1)Fluorite CaF2 = Ca2+ + 2F − −10.6 (2)

Sulphates:Gypsum CaSO4 . 2H2 O = Ca2+ + SO2−

4 + 2H2 O −4.58 (2)Anhydride CaSO4 = Ca2+ + SO2−

4 −4.36 (2)Barite BaSO4 = Ba2+ + SO2−

4 −9.97 (2)Carbonates:

Calcite CaC O3 = Ca2+ + C O2−3 −8.48 (2)

Aragonite CaC O3 = Ca2+ + C O2−3 −8.34 (2)

Dolomite CaMg (C O3)2 = Ca2+ + Mg2+ + 2C O2−3 −17.1 (2)

Siderite F eC O3 = F e2+ + C O2−3 −10.9

Hydroxides:Gibbsite Al (OH)3 = Al3+ + 3OH− −33.5 (1)Goethite α · F eOOH + H2 O = F e3+ + 3OH− −41.5 (1)

Silicates:Quartz Si O2 + 2H2 O = Si (OH)4 −3.98 (2)Chalcedony Si O2 + 2H2 O = Si (OH)4 −3.55 (2)Amorphous silica Si O2 + 2H2 O = Si (OH)4 −2.71 (2)

the solution, and a zero value of SI indicates thatthe system is in equilibrium.

The equilibrium constant for a mineral–solute re-action is called the solubility product and is de-noted by Kso. Like all equilibrium constants, thesolubility product is defined by an equation suchas Eqn 6.10; for example, for the dissolution/precipitation reaction for the mineral anhydrite:

CaSO4 ↔ Ca2+ + SO2−4

would be:

Kso =[Ca2+] [

SO2−4

]

[CaSO4]= [

Ca2+] [SO2−

4

](6.24)

where

[Ca2+]

and[SO2−

4

]

are activities of the two ions. Under equilibriumconditions the activity of the solid phase equals 1so that:

[CaSO4] = 1

Solubility products for standard conditions of tem-perature and pressure for some common mineralsare given in Table 6.5.

The amount of mineral that can be dissolvedin a given volume of water depends on theinitial concentration of dissolution products inwater. More dissolution occurs if water has a lowinitial concentration than if it has a high initialconcentration of the minerals to be dissolved. Thisexplains why caves develop at shallow depthsin recharge areas in limestone (calcite) country.The freshly infiltrated pore waters in these areashave low concentrations of dissolution products,namely:

Ca2+, CO2−3 , and HCO−

3

Therefore, limestone dissolves rapidly creatingvoids that eventually give rise to caves along frac-tures through which water seepage occurs. As wa-ter seeps down, concentration of the dissolutionproducts keeps increasing, which results in pro-gressively less dissolution of the limestone. It isalso interesting to note that solubility of solid salts

P1: SFK/UKS P2: SFK



in water, and in most other solvents, increases withtemperature while that of gases decreases.

6.1.6 Gas–water partitioning

Due to surface tension of water, the surface of awater body in contact with the atmosphere rathertends to be maintained. However, water moleculesare able to cross the surface and escape into theatmosphere and gas molecules from the air can dif-fuse into the water body. Both processes operatingsimultaneously tend to produce saturation near theair–water interface. Rates of absorption of gases bywater or rates of evaporation of water are functionsof: (i) the total surface area of the air–water inter-face; (ii) the degree of departure from saturation inthe layer immediately below the interface (humid-ity in the case of water vapour); and (iii) the rateat which the molecules of the dissolved or vapourphase are transported away from the interface. Thetransport rate would be slow if it were controlledsolely by molecular diffusion. In most natural sys-tems, however, motion of the gas or liquid phasehelps to transport the evaporated or dissolved ma-terial away from the interface.

The processes by which gases from the atmo-sphere dissolve in water are of direct relevanceto water quality. As the dissolved oxygen is es-sential to aquatic organisms, the occurrence, dis-solution, and transport of oxygen are importantin the study of biochemical processes relating towater pollution. The process of photosynthesis isa major source of atmospheric oxygen. Langbeinand Durum (1967) reviewed some properties ofstream-channel geometry and stream flow rates asapplicable to the role of rivers in the uptake ofoxygen from the atmosphere. Understanding sys-tems of this kind entails consideration of rates andthe manner in which the rate of one process may af-fect the rate of another. Some gases, notably carbondioxide, react with water and their rate of assim-ilation is affected by subsequent changes in theirform. Carbon dioxide is an important constituentin many geochemical processes.

Gas partitioning at the free air–water interfacecan be reasonably well described by Henry’s lawfor solubility of a gas into a liquid. According tothis law, concentrations in the two phases are di-

rectly proportional to each other (see Section 7.5 inChapter 7; Eqn. 7.18). If same concentration units(e.g. mol l−1) are used for both phases, the Henrycoefficient, ki, is ‘dimensionless’. It depends ontemperature, T , and concentrations of all dissolvedspecies, cj (Ballentine and Burnard 2002). In thisform, Henry coefficients (ki) are the inverse of ‘sol-ubility of a gas in water’. Poorly soluble gases havelarge values of ki and vice versa.

Generally Eqn. 7.18 is formulated for each gas interms of its partial pressure, pi, (expressed in theunits of atmosphere, bar, . . .) and correspondingequilibrium concentration of gas Ci,eq (cm3 STP g−1

in water, mol l−1, mol kg−1, mol mol−1). STP de-notes Standard Temperature (T = 0◦C = 273.15 K)and Pressure (P = 760 mmHg = 1 atm). To dif-ferentiate between different units that are in use,Henry’s law is written with coefficient, Hi, (in ap-propriate units) as:

Hi = pi

Ci,eq(6.25)

Partial pressure of a gas A in a gas mixture is de-fined as the pressure that gas A would exert, if itwas the only gas occupying the same volume asthe total gas mixture. The sum of partial pressuresof each component gas equals the total pressureof a gas mixture (Dalton’s Law). For example, inthe Earth’s atmosphere, about 21% of moleculesare O2 and 0.036% are CO2. Assuming a sample ofthe atmosphere at 1 atm pressure, the partial pres-sure of O2 is 0.21 atm and the partial pressure ofCO2 is 0.00036 atm. The atmospheric CO2 partialpressure is, therefore, equivalent to 360 ppmv.

Sometimes, instead of Henry’s coefficient, theequilibrium concentration (Ci,eq) for pi = 1 atm isalso reported. This last formulation of Henry’s Lawconstant is in a way the inverse of Eqn 6.25 (orEqn. 7.18), and is related to solubility of a givengas expressed in appropriate units (e.g. mol l−1

atm−1). To differentiate this formulation from theprevious two formulations of Henry’s constant, thecoefficient, KH, is used:

KH = Ci,eq

pi(6.26)

Equilibrium constants (KH) for dissolution of someimportant natural gases are listed in Table 6.6.

P1: SFK/UKS P2: SFK



Table 6.6 Dissolution equilibrium constants of somenatural gases at 25◦C. After Stumm and Morgan (1996).

Dissolution ReactionLog(KH)(mol l−1 atm−1)

C O2(g) + H2 O ↔ H2C O∗3 −1.47

C O(g) ↔ C O(aq) −3.02O2(g) ↔ O2(aq) −2.90O3(g) ↔ O3(aq) −2.03N2(g) ↔ N2(aq) −3.18C H4(g) ↔ C H4(aq) −2.98

∗ Represents the sum of [CO2(aq)] and [H2CO3] formeddue to hydration of a small part of [CO2(aq)]. For all valuesof pH, [CO2(aq)] [H2CO3].

Example 6.2. Calculate the dissolved concentra-tion of O2 in water that is in equilibrium with theatmosphere at 25◦C.

Solution. Partial pressure of oxygen is 0.21 atm(21 000 ppm). Using Eqn 6.26, and dissolution co-efficient of O2 from Table 6.6:

O2(aq) = pO2 .KH

= (0.21 atm)(10−2.90 M.l−1 atm−1

)

= 2.6 × 10−4 M.l−1

6.1.7 Carbonate reactions, alkalinity,and hardness

In most natural waters, acid–base reactions andpH are dominated by the interaction of carbondioxide and the aqueous carbonate compoundsH2C O∗

3 (dissolved CO2), HC O−3 (bicarbonate),

and C O2−3 (carbonate). The dissolution reaction of

atmospheric CO2(g) and the associated acid–basereactions between the carbonate compoundsare:

CO2(g) + H2O ↔ H2CO∗3 (6.27)

H2CO∗3 + H2O ↔ H3O+ + HCO−

3 (6.28)

HCO−3 + H2O ↔ H3O+ + CO2−

3 (6.29)

The equilibrium equations and the constants forthese reactions at 1 atm pressure and 25◦C tem-

perature are:

KCO2 =[H2CO∗

3

]

pCO2

= 10−1.47 atm−1 (6.30)

KH2CO3 = [H+][HCO−

3

][H2CO∗

3

] = 10−6.35 (6.31)

KHCO−3

= [H+][CO2−

3

][HCO−

3

] = 10−10.38 (6.32)

The constant defined in Eqn 6.30 is the same as thefirst constant in Table 6.6, except that the constanthere is defined with respect to activity

[H2CO∗

3

]in-

stead of concentration(H2CO∗

3

). Eqn 6.31 and Eqn

6.32 represent two equations with four unknowns,namely:

[H+],[H2CO∗

3

],[HCO−

3

], and

[CO2−

3

]

When pH (the negative logarithm of hydrogen ionconcentration) is known, [H+] is known and, there-fore, the ratios:

[HCO−

3

] / [H2CO∗

3

]and

[HCO−

3

] / [CO2−

3

]

are also known, without knowing the magnitudeof individual carbonate activities. The relative dis-tribution of dissolved carbonate species as a func-tion of pH is shown in Fig. 6.2. It is seen thatH2CO∗

3 is the dominant species when pH <6.35;HCO−

3 is dominant when 6.35 < pH < 10.38; andCO2−

3 is dominant when pH >10.38. The pH ofmost natural waters falls in the range 6.5 < pH< 10, in which HCO−

3 is the dominant carbonatespecies.

In the saturated zone, the carbonate system isclosed as there is no direct contact with the gasphase and Eqn 6.30 does not apply. The assumptionof equilibrium with atmospheric CO2 is reasonablein open systems such as streams and some unsat-urated pore waters. The partial pressure of CO2 insoil gas is often higher than in the atmosphere dueto plant decay and root respiration. In open sys-tems, H2CO∗

3 is independent of pH, fixed by theatmospheric pCO2 . Given the pH of water and pCO2 ,the equilibrium activities of all carbonate speciesmay be calculated for an open system, as shownbelow.

P1: SFK/UKS P2: SFK



Fig. 6.2 Distribution of the dissolved carbonatespecies as function of pH for a closed natural watersystem. Adopted from Fitts (2002). © Academic Press.

Example 6.3. Determine the activities of all car-bonate species and the molar concentration of bi-carbonate for an open equilibrium system of waterwith ionic strength I = 0.01 M, pH = 5.7, and anatmospheric CO2 concentration of 360 ppm.

Solution. Atmospheric CO2 concentration of 360ppm = pCO2 = 3.60 × 10−4 atm. Using Eqn 6.30,the activity of H2CO∗

3 is calculated as:

[H2CO∗

3

] = KCO2 pCO2 = (10−1.47 atm−1

)

× (3.60 × 10−4 atm

) = 1.22 × 10−5

Next, HCO−3 is now calculated using Eqn 6.31:

[HCO−

3

] = KH2CO3

[H2CO∗

3

]

[H+]

= (10−6.35) (1.22 × 10−5)

10−5.7= 2.73 × 10−6

Finally, using Eqn 6.32,[CO2−

3

]is calculated as:

[CO2−

3

] = KHCO−3

[HCO−

3

]

[H+]

= (10−10.38) (2.73 × 10−6)

10−5.7= 5.71 × 10−11

In this case, about 80% of the dissolved carbonateis H2CO∗

3; about 20% is HCO−3 ; and a negligible

fraction is CO2−3 . This is consistent with Fig. 6.2

for pH 5.7.

The bicarbonate molar concentration is calcu-lated from its activity, and activity coefficient for

ionic strength I = 0.01 M:

(HCO−

3

) =[HCO−

3

]

γHCO−3

= 2.73 × 10−6

0.90

= 3.03 × 10−6 M

6.1.7.1 Alkalinity

Total alkalinity is a parameter that defines the ca-pacity of water to neutralize acid added to it. It isdefined as:

Alk = (HCO−3 ) + 2

(CO2−

3

) + (OH−) − (H+)

(6.33)

Total alkalinity is measured in units of equivalentcharge, generally expressed as milli-equivalents perlitre (meq l−1), which is related to concentrationas:

meq

L= mg

L.

Z

gram formula weight= mmol

L.Z

(6.34)

where Z is the charge on the ion, which is also itsvalency. Sometimes alkalinity is also expressed asmg CaCO3 l−1, mg HCO3

− l−1, or mg Ca l−1.Equivalent charge units are used because hydro-

gen ions are neutralized by their charge rather thantheir mass. Thus a carbonate ion CO2−

3 can neu-tralize two hydrogen ions and, therefore, [H+] =2[CO2−

3 ], with concentration expressed in molarunits. The standard adopted for alkalinity is purewater in which only CO2 is dissolved. In this state,only ions on the right-hand side of Eqn 6.33 are in

P1: SFK/UKS P2: SFK



solution and the charge balance requires Alk = 0.There are other definitions of alkalinity, dependingon the reference state chosen, but pure water iscommonly used as the standard for natural waters.

Adding a base to water is equivalent to addingcations as well as OH− ions to the solution, whileadding acid generally adds anions plus H+ ions tothe solution. Therefore, adding a base increases[Alk] and adding an acid decreases [Alk]. Waterwith high level of alkalinity is more effective in neu-tralizing acid that is added to it. In acidic waters,Alk <0. Dissolution of basic minerals, especiallycarbonate minerals, in water tends to increase itsalkalinity. Therefore, as groundwater moves froma recharge area to a discharge area, its alkalinitytypically changes from low to high.

Caustic alkalinity is the amount of H+ requiredto decrease the pH to ∼10.8, which equals thestoichiometric amount of H+ required to completeonly the reaction:

H+ + O H− = H2O

In this case, all the carbonate species will bepresent entirely as CO3

2−. The amount of acid re-quired to reach this end point cannot be deter-mined readily because of the poorly defined endpoint, caused by the masking effect of the bufferingof water. But this can be calculated if the amountsof carbonate and the total alkalinity are known.

Carbonate alkalinity, also called phenolph-thalein alkalinity, is the amount of H+ required tolower the pH to 8.3 (phenolphthalein end point).It is often determined by the colour change of thephenolphthalein indicator. It equals the stoichio-metric amount of H+ required to complete the tworeactions: (i) H+ + OH− = H2O; and (ii) H+ +CO3

2− = HCO3−.

Therefore, total alkalinity, sometimes alsocalled methyl orange alkalinity, is the amount ofH+ required to decrease pH to about 4.5 (methylorange end point). In a system dominated bycarbonate, the H+ added is the stoichiometricamount required for the 3 reactions to take place:(i) H+ + OH− = H2O; (ii) H+ + CO3

2− = HCO3−;

and (iii) H+ + HCO3− = H2CO3.

Total acidity is the number of moles/litre of OH−

that must be added to raise the pH of the solutionto approximately 10.8, or to whatever pH is con-

sidered to represent a solution of pure Na2CO3 inwater at the concentration of interest.

CO2 acidity is the amount of OH− required totitrate a solution to a pH of 8.3. This assumes, ofcourse, that the pH of the solution is initially below8.3. Such a solution contains H2CO3 as a majorcomponent, and the titration consists of convertingthe H2CO3 into HCO3

−.If we assume that the alkalinity of water is due

solely to the carbonate species and OH−, thefollowing deductions about the initial compositionof a solution from an alkalinity titration can bemade to a reasonably good approximation. Theseare based on the assumption that the titration is aclosed system, (which is reasonable if the solutionis not shaken and if the titration is conducted).Under these conditions each mole of CO3

2−

present will consume one mole of H+ when thesolution is titrated to pH 8.3 and another mole ofH+ as it is titrated from pH 8.3 to pH ∼4.5.

� If the volume of acid to reach pH 8.3 (Vp) isequal to the volume of acid required to proceedfrom pH 8.3 to 4.5 (Vmo), then the original solu-tion contains only CO3

2− as the major alkalinitycausing species.

� If, on adding a phenolphthalein indicator to thesolution, it immediately becomes colourless (i.e.if the initial pH is below 8.3) and a volume Vmo

is required to reach pH of about 4.5, the originalsolution must have contained only HCO3

− as themajor anion species contributing to alkalinity.

� If the original solution requires Vp ml of acidto reach pH 8.3, but no further acid addition isrequired to reach pH 4.5, the alkalinity is due toOH− alone.

� By similar reasoning, it can be shown that if Vp >

Vmo, then the major alkalinity causning speciesare OH− and CO3

2−.� If Vmo > Vp, then the major species are CO3

2−

and HCO3−.

� The other possible combination of majoralkalinity-causing species, OH− and HCO3

−, can-not exist because there is no pH range overwhich these two species are concurrently themajor alkalinity causing species (Fig. 6.2).

Table 6.7 is a summary of the alkalinity titrationrelationships as explained above.

P1: SFK/UKS P2: SFK



Table 6.7 Approximate relations between the results of an alkalinity titration and the concentrations of predominant speciesin carbonate solutions.

Condition Predominant form of alkalinity Approximate Molar Concentration (M)

Vp = Vmo CO32− [CO3

2−] = VpN/VVp = 0 HCO3

− [HCO3−] = VmoN/V

Vmo = 0 OH− [OH−] = VpN/VVmo > Vp CO3

2− and HCO3− [CO3

2−] = VpN/V; [HCO3−] = (Vmo – Vp) N/V

Vp > Vmo OH− and CO32− [CO3

2−] = VmoN/V; [OH−] = (Vp – Vmo)N/V

Source: http://www.cheml.com/acad/pdf/c3carb.pdf (accessed on 18th April, 2009)Vmo and Vp are the volumes of strong acid of normality, N , required to reach the end points atpH 4.5 and 8.3, respectively. V is the initial volume of the solution.

It may be noted that the stated end points areapproximate values, as the actual end points arenot precisely fixed values and vary with the totalcarbonate concentration in solution.

Some important considerations for understand-ing the behaviour of natural waters in whichcarbonate species are the principal bufferingagents are:

� Alkalinity and acidity are conservative parame-ters, unaffected by temperature, pressure, andactivity coefficients. Note that this is not true forthe pH, which varies with all of these factors.

� Water can simultaneously possess both alkalinityand acidity. Indeed, this is the usual case over thepH range in which HCO3

− predominates.� Addition or removal of CO2 (e.g. by the action of

organisms) has no effect on [Alk]. However, thisaffects both the acidity and total carbonate con-centration; an increase in [H2C O∗

3 ] raises bothof these quantities.

� Addition or removal of solid CaCO3 or other car-bonates has no effect on the acidity. Thus acidityis conserved in solutions that are brought intocontact with calcite and similar sediments.

� In a system that is closed to the atmosphere andis not in contact with solid carbonates, the to-tal carbonate concentration is unchanged by theaddition of a strong acid or a strong base.

� The presence of mineral acidity or caustic alka-linity in natural water is indicative of a sourcerelated to industrial pollution. The limits of pHrepresented by these two conditions correspondroughly to those that are well tolerated by mostliving organisms.

6.1.7.2 Hardness

Another parameter associated with the carbonatechemistry of water is known as the hardness of thewater. Water hardness is the sum of the bivalentcations in water, expressed as equivalent CaCO3.The major bivalent cations are calcium, Ca2+ andmagnesium, Mg2+, though there may also be a mi-nor contribution from iron, Fe2+, and divalent man-ganese, Mn2+. These bivalent cations react withsoap to form a soft precipitate or form solid pre-cipitates that form scale on the inner surface ofboilers.

Hardness = (Ca2+ + Mg2+) as mg CaCO3

per litre of water (6.35)

Sometimes hardness is also expressed as mg Caper litre of water. Common descriptions of waterhardness, in terms of mg CaCO3 per litre of water,are given in Table 6.8.

Table 6.8 Common descriptions of water hardnessbased on WHO (2009).

Hardness(as mg CaCO3.l−1) Description

0–50 Soft50–100 Moderately soft100–150 Slightly hard150–200 Moderately hard200–300 Hard>300 Very hard

P1: SFK/UKS P2: SFK



Example 6.4. What are the alkalinity and hardnessof a groundwater sample from a limestone area withHCO3

− = 270 mg l−1; Ca = 55 mg l−1; and Mg =30 mg l−1?

Solution. Atomic and molecular weights of theatoms and molecules involved are:

Atomic weights:H = 1; C = 12; O = 16; Mg = 24.3; Ca = 40

Molecular weights:HCO3

− = 1 + 12 + 16 + 16 + 16 = 61CaCO3 = 40 + 12 + 16 + 16 + 16 = 100

Alkalinity:

270 mg l−1 HCO−3 = 270

61= 4.43 mmol l−1 HCO−

3

= 4.43 × 100 mg l−1 CaCO3

Since the valency of HCO3− is 1, the alkalinity

of groundwater in various units is 4.43 meq l−1;HCO3

− = 270 mg l−1; HCO3− = 443 mg l−1 CaCO3.

Hardness:

55 mg l−1 Ca = 55

40= 1.375 mmol l−1 Ca

30 mg l−1 mg = 30

24.3= 1.235 mmol l−1 mg

Therefore, Hardness = 1.375 + 1.235 = 2.61 mmoll−1 = 261 mg l−1 CaCO3.

6.1.8 Electro-neutrality

All solutions are required to be electrically neutral,that is, in any given volume of water the sum ofelectrical charges of all cations must equal the sumof charges of all anions. If the number of cationsand anions present in a solution are j and k, respec-tively, the electro-neutrality condition requires:

j∑i=1

ce+i =

k∑i=1

ce−i (6.36)

where cei+ and cei

− are the charge concentrationof the ith cation and ith anion, respectively, ex-pressed in equivalents or milli-equivalents per litreunits. When results of water analysis are pluggedinto Eqn 6.36, they should be approximately equal,otherwise either the analysis is erroneous or oneor more significant ions are being inadvertentlymissed from the analysis.

Example 6.5. Refer to Table 6.9 for calculation.

6.1.9 Oxidation and reduction reactions

Oxidation and reduction (redox) reactions trans-fer energy to many inorganic and life processes.Many common processes such as photosynthesis,

Table 6.9 Calculation of ionic strengths and electro-neutrality – an illustrative example.

Conc. Mol. wt. Valency Molar Conc. Ionic strength %

Cations (mg l−1) (g) mM l−1 (meq l−1)Na+ 8 23 1 0.35 0.35 17K+ 10 39 1 0.26 0.26 13Ca2+ 12 40 2 0.30 0.60 30Mg2+ 10 24.3 2 0.41 0.82 40

Sum 2.03AnionsCl− 10 35.5 1 0.28 0.28 16SO4

2− 10 96 2 0.10 0.21 11HCO3

− 80 61 1 1.31 1.31 73Sum 1.80

Electro-neutrality, EN (%) =∑

Cations (meq.l−1) − ∑Anions (meq.l−1)∑

Cations (meq.l−1) + ∑Anions (meq.l−1)

× 100

= 6.0%

P1: SFK/UKS P2: SFK



respiration, corrosion, combustion, and even bat-teries involve redox reactions. Usually redox reac-tions usually proceed only in one direction, slowlytowards completion, but seldom reach there.

In a chemical reaction when an electron is trans-ferred from one atom to another, the processes ofoxidation and reduction occur simultaneously; theatom gaining the electron is reduced and that los-ing the electron is oxidized. The term ‘oxidation’is used because oxygen has a strong tendency toaccept electrons, which itself gets reduced whileoxidizing other atoms that donate electrons to it.

The oxidation number, or oxidation state,refers to a hypothetical charge that an atom wouldhave if it were to dissociate from the compoundof which it is a part. It provides a useful methodto study redox reactions. Oxidation numbers aredenoted with Roman numerals and are calculatedwith the following set of rules (Stumm and Morgan1996):

� The oxidation state of a monoatomic substanceis given by its electric charge.

� In a compound formed by covalent bonds be-tween various atoms, the oxidation state of eachatom is the charge remaining on the atom wheneach shared pair of electrons is assigned com-pletely to the more electro-negative of the twoatoms sharing a covalent bond.

� The sum of oxidation states of various atoms of amolecule equals zero for neutral molecules, andfor ions it equals the charge on the ion.

Electro-negativity of an element is the measure ofits affinity for electrons; the higher the electro-negativity, the more it tends to attract and gainelectrons. Bonds between different atoms withsimilar electro-negativity tend to be covalent,while bonds between atoms with very differentelectro-negativity tend to be ionic. The followingis a list of some common elements in order ofdecreasing electro-negativity:

F , O, Cl, N, C = S, H, Cu, Si, Fe, Cr, Mn

= Al, Mg, Ca, Na, K

In a compound, the Group 1A elements (H, Li,Na, . . .) are usually with oxidation number (+I);

the Group 2A elements (Be, Mg, Ca, . . .) are usu-ally (+II); and for oxygen it is usually (–II). Theoxidation number is reduced to a lower value onreduction of an element and is increased on itsoxidation. Several common substances and theiroxidation states are listed in Table 6.10.

In the following redox reaction involving oxida-tion of iron, oxygen is reduced from O(0) to O(–II):

O2 + 4F e2+ + 4H+ = 4F e3+ + 2H2O (6.37)

while iron is oxidized from Fe(+II) to Fe(+III); hy-drogen remains H(+I) and is neither oxidized norreduced. This redox reaction can also be thoughtof as the sum of two half reactions as below:

O2 + 4H+ + 4e− = 2H2O (6.38)

4Fe2+ = 4Fe3+ + 4e− (6.39)

It is clear from the above-mentioned reactions thatO2 is an electron acceptor and Fe3+ is an elec-tron donor. Different waters vary in their tendencyto oxidize or reduce, depending on the relativeconcentration of electron acceptors (i.e. O2) anddonors (i.e. Fe(0)), respectively.

Just as pH is a measure of the hydrogen ion con-centration [H+] of a solution, a measure of the ox-idizing or reducing tendency of water is the elec-tron activity [e−]. Similar to pH, the parameter pE is

Table 6.10 Oxidation states of some common substances.After Fitts (2002).

Substance Element and its Oxidation state

H2O H(+I) O(-II)O2 O(0)NO3

− N(+V) O(-II)N2 N(0)NH3, NH4

+ N(-III) H(+I)HCO3

− H(+I) C(+IV) O(-II)CO2, CO3

− C(+IV) O(-II)CH2O C(0) H(+I) O(-II)CH4 C(-IV) H(+I)SO4

2− S(+VI) O(-II)H2S, HS− (H+I) S(-II)Fe2+ Fe(+II)Fe(OH)3 Fe(+III) O(-II) H(+I)Al(OH)3 Al(+III) O(-II) H(+I)Cr(OH)3 Cr(+III) O(-II) H(+I)CrO4

2− Cr(+VI) O(-II)

P1: SFK/UKS P2: SFK



defined as the negative logarithm of electron activ-ity, i.e.:

pE = − log[e−] (6.40)

The electron activity is a measure of the relativeactivity of electron donors and electron acceptorsthat are in solution. The actual electron concen-tration in a solution is likely to be quite low aselectrons are not ‘free’, except very briefly duringredox reactions. High pE (i.e. low [e−]) has fewerreducing species (electron donors) than oxidizingspecies (electron acceptors). Conversely, low pEwaters have an excess of reducing species com-pared to oxidizing species. The pE scale is estab-lished by the convention of assigning K = 1 forthe equilibrium reduction of hydrogen at standardpressure-temperature condition (Morel and Hering1993):

H+ + e− = 1

2H2 (6.41)

For a general half-reaction involving reduction ofOX to RED:

OX + ne− = RED (6.42)

the pE at equilibrium can be estimated as:

K = [RED]

[OX][e−]or pE = − log[e−]

= 1

n

[log K + log

[OX]

[RED]

](6.43)

where K is the equilibrium constant for the halfreaction OX reduced to RED as above. Any redoxpair that is in equilibrium should have the same pEvalue based on Eqn 6.43. Thus there is a uniquevalue of pE for a solution at redox equilibrium.If there is a consistent computed value of pE forseveral redox pairs but a different computed valuefor one of the redox pairs, that particular pair isprobably not in equilibrium.

A parameter closely related to pE is redox poten-tial, Eh, defined as:

Eh = 2.3 RT

FpE = (0.059V ) pE (6.44)

where R is the universal gas constant; T is temper-ature (in Kelvin); and F is Faraday’s Constant. Ehhas units of volt and is essentially proportional to

pE. Eh can be measured through electrochemicalreactions.

Elements that are involved in redox reactions aresensitive to pE in the same way that acids andbases are sensitive to pH. A pE-pH or Eh-pH dia-gram shows which species of a particular elementwould be stable at equilibrium under a given rangeof pE and pH conditions. Such diagrams are use-ful for predicting changes that may be expectedwith change of pE and/or pH. For example, pre-cipitates and scales/stains often form when low-pE anoxic groundwaters are pumped to the sur-face and get exposed to the atmospheric oxygen,which suddenly raises the pE of water. DissolvedFe2+ forms a red-brown precipitate of ferric hy-droxide, Fe(OH)3, on bathroom showers, sinks,laundry, etc. Similarly, manganese-rich water formsa dark brown to black precipitate and copper-richwater forms a bluish precipitate.

6.1.9.1 Biogeochemical redox reactions

Microbially mediated redox processes are regulatedby the ratio of accessible oxidizing agents to theamount of available and degradable organic sub-stances in the water. The resulting biogeochemicalconditions induce further chemically and micro-bially mediated transformations.

During photosynthesis, carbon is reduced, asC(IV ) → C(0), while some of the oxygen is ox-idized as O(–II) → O(0). Photosynthesis occursonly with the availability of additional energy fromthe sunlight. There are other reactions involvingsulphur and nitrogen but they produce far less or-ganic matter than photosynthesis. Hydrocarbonsstore chemical energy, which can be released sub-sequently through a variety of redox reactions.

During aerobic degradation of organic carbon,bacteria present use oxygen directly as an oxidiz-ing agent. Once the available oxygen is consumed,other bacteria take over and use the remaining ox-idants that are present in water during the courseof a characteristic sequence of reactions (Fig. 6.3).The important ones amongst these oxidizing agentsare nitrates, sulphates, and carbon dioxide dis-solved in water as well as solid oxides/hydroxidesof manganese and iron. In chemical parlance,microbial decomposition involves transfer of

P1: SFK/UKS P2: SFK



Manganese oxide reduction

Redox potential

anoxic

oxic

-5 0 +5 +10 P

-0.5 0 +0.5 E [V]h

Oxidisingagent(e-acceptor)

Reducingagent(e-donor)

Aerobic respiration O2 + Corg +CO2H O2

Denitrification NO3– + Corg + CO2+HCO3

–N2

MnO2(S) + Corg +HCO3–

Mn2+

Iron oxide reduction FeOOH(S)+ Corg +HCO3–

Fe2+

Sulphate reduction SO42– + Corg + CO2+HCO3

–HS–

Methane formation CO2 + Corg +CO2CH4

NH4+

Anaerobicmineralisationof Norg

Redox indicators in water

Fig. 6.3 Sequence of the major microbially mediated redox processes in aquatic systems. Redrawn from Zobrist et al. (2000).© EAWAG.

electrons mediated by bacteria from organic car-bon to the various oxidizing agents. These reactionstake place in a typical sequence that is governed bythe chemical energy released during each reaction.Energy output is the highest during the first reac-tion and diminishes steadily thereafter (Fig. 6.3).The energy liberated is used by bacteria for sustain-ing their metabolism and growth. The oxidizingagents and products of the redox processes canserve as redox indicators in water. From their pres-ence or absence, the status of redox processes andconditions in the aquifer can be inferred.

When all the oxygen has been consumed in anaquatic system, its biogeochemical environmentchanges drastically, because products of the redoxprocesses initiate secondary geochemical reactions(von Gunten and Zobrist 1993) as:

� CO2 reacts with the rocks and mineral grainspresent in an aquifer.

� Dissolved iron and manganese, Fe(II) andMn(II), are precipitated by sulphides and car-bonates.

� Any surplus sulphide reductively dissolves ironand manganese (hydr)oxide. It is now knownthat there are certain additional chemicalreactions that do not take place under oxidizing

conditions. For example, Fe(II) is also a reac-tive reducing agent that can react both withinorganic and organic pollutants (Haderlein2000). Anoxic water must be treated before itcan be used as drinking water, since dissolvedmanganese and iron are just as undesirable assulphides, from health as well as taste consid-erations. From a microbiological point of view,anoxic groundwater may contain anaerobicbacteria that decompose organic pollutants thatare not persistent in the presence of oxygen(Van der Meer and Kohler 2000).

Rainwater, due to its contact with the atmosphere,is saturated with dissolved O2, resulting in O2 con-centration close to 10 mg l−1. As water infiltratesinto the ground, oxidation of soil organic mattertends to decrease the amount of dissolved O2

and increase the level of dissolved CO2. Further,oxidation reactions in the saturated zone that isisolated from the atmosphere explain progres-sively lower O2 levels in the infiltrated water as itmoves down into the aquifer. Some groundwatersare anoxic and other oxidation reactions may beimportant (Fig. 6.3). The unpleasant ‘rotten egg’smell of hydrogen sulphide occurs as a by-productof bacterially mediated sulphate reduction in some

P1: SFK/UKS P2: SFK



anoxic groundwaters. The odour of H2S can bedetected at concentrations as low as 0.1 mg l−1

(Tate and Arnold 1990).

6.1.10 Sorption

Sorption is a combination of two processes,namely, adsorption (implying attachment onto theparticle surfaces), and absorption (implying in-corporation into something). Many of the solutes,particularly nonpolar organic molecules and cer-tain metals, get sorbed onto the surfaces of solidsin the aquifer matrix, and onto the rocks, minerals,and sediments. Some of these surfaces are internalto a grain or a rock mass, so technically these casescan be called absorption. Sorption slows downmigration of solutes and is, therefore, a key processin the fate and transport of dissolved contaminants.For instance, certain compounds such as methyltertiary butyl ether (MTBE), a constituent of gaso-line, might sorb negligibly and migrate at aboutthe same rate as the average water molecules,while other compounds such as toluene, also aconstituent of gasoline, might sorb strongly and mi-grate at a much slower rate. As a plume of dissolvedgasoline migrates, the constituents get segregated,depending on the degree of their sorption.

Most sorption reactions are relatively rapid,approaching equilibrium on time scales of minutesor hours (Morel and Hering 1993). Reachingequilibrium can take much longer if the process

is diffusion-limited in either the liquid or the solidphase.

6.1.11 Metal complexes and surfacecomplexation of ions

In water, free metal cations are actually surroundedby a layer of water molecules. This coordinated rindof molecules closest to the metal cation usually con-sists of two, four, or six H2O molecules, six beingthe most common (Morel and Hering 1993). Forexample, dissolved Cr(III) tends to be coordinatedas Cr(H2O)6

3+. Beyond the coordinated rind, thewater molecules are not chemically bonded butare oriented due to the electrostatic charge on thecation, the degree of orientation decreasing withdistance from the cation.

Anions, also called ligands, may displace someor all of the coordinated water molecules and bondwith the central metal cation. Common ligands inwater include the major anions (Table 6.1), otherinorganic anions, as well as a variety of organicmolecules. All ligands have electron pairs availablefor sharing to form bonds with central metal cation.

If the metal and ligands bond directly to eachother, displacing the coordinated water, one hasan inner sphere complex or ion complex. If, onthe other hand, a metal cation along with its co-ordinated water molecule combines with a ligandanion and its coordinated water molecules throughelectrostatic bonding, it leads to the formation of

Table 6.11 Dominant species of trace metals in natural waters. After Morel and Hering (1993).

Metal Aqueous Species Solid Species

Al Al(OH)3, Al(OH)4−, AlF2+, AlF2+ Al(OH)3, Al2O3, Al2Si2O5(OH)4, Al-silicates

Cr Cr(OH)2+, Cr(OH)3, Cr(OH)4−HCrO4−, CrO4

2− Cr(OH)3

Cu Cu2+, CuCO3, CuOH+ CuS, CuFeS2, Cu(OH)2, Cu2CO3(OH)2,Cu2(CO3)2(OH)2, CuO

Fe Fe2+, FeCl+, Fe2SO4, Fe(OH)2+, Fe(OH)4− FeS, FeS2, FeCO3, Fe(OH)3, Fe2O3, Fe3O4, FePO4,Fe3(PO4)2, Fe-silicates

Hg Hg2+, HgCl+, HgCl2, HgCl3−, HgOHCl, Hg(OH)2,HgS2

2−, HgOHS−HgS, Hg(OH)2

Mn Mn2+, MnCl+ MnS, MnCO3, Mn(OH)2, MnO2

Pb Pb2+, PbCl+, PbCl2, Pb(OH)+, PbCO3 PbS, PbCO3, Pb(OH)2, PbO2

P1: SFK/UKS P2: SFK



an outer sphere complex or ion pair and the bond-ing tends to be weak and ephemeral. The num-ber of sites where the central metal atom bondsto ligands is its coordination number. Some lig-ands form bond with metal atom at only one siteand are called unidentate ligands. When the lig-ands bond to metal atoms at multiple sites, theseare called multidentate ligands. Chelates are com-plexes with multidentate ligands and one centralmetal atom. Chelates tend to be more stable thanunidentate complexes. The major inorganic cationsin fresh waters do not typically form complexes toany significant extent. In general, complexation re-actions are rapid and, therefore, assumption of theequilibrium condition is reasonable. Unlike majorcations, many trace metals in water are predom-inantly in the form of complexes as opposed tofree ions. Some of the most common species ofcommon trace metals, based on equilibrium condi-tion in natural waters, are listed in Table 6.10. Thespeciation of metal complexes varies with pH andpE, particularly when there are complexes involv-ing hydroxides, carbonates, or other pH sensitiveligands.

6.1.11.1 Surface complexation of ions

In almost every natural setting, there is a largeamount of mineral surface area in contact withwater, particularly groundwater. Metals and lig-ands in mineral solids are incompletely coordinatedwhen they are at the surface and, therefore, havea tendency to coordinate with ligands and metalions present in the interstitial water. Coordinationof atoms and functional groups that are part of thesemineral surfaces is essentially similar to complexa-tion in the aqueous phase.

Many mineral solids, particularly clay minerals,carry a net charge on their surfaces. This causes thewater in contact with these minerals to becomestructured with respect to ion concentrations.The majority of clays have a negative chargeon their surfaces. In the layer of water closestto the surface, water molecules are orientatedwith their hydrogen ions towards the surface andthe concentration of cations is higher relative tothat of anions. Cations in this zone are known ascounter ions because they counter the charges

on the mineral surfaces. The negative mineralsurface charge and the net positive charge onthe surrounding water together form a chargedistribution called a double layer. Metal cations insolution react with the coordinated hydroxyl ionspresent on mineral surfaces, displacing hydrogenand becoming sorbed onto the mineral surfaces.Similar to complexation in the aqueous phase dis-cussed above, an inner sphere complex is formedwhen the metal ions bond covalently to atoms onthe mineral surface and outer sphere complex isformed when the metal cation is loosely boundto the mineral surface with intervening watermolecules. Since both inner- and outer- surfacecomplexes are removed from the mobile aqueousphase, these are considered as sorbed.

Cations not only compete with H+ ion for sorp-tion sites, but also compete with each other. Ionexchange occurs when one type of ion displacesanother from a coordination site at a mineral sur-face. The ion exchange process differs markedlyfrom mineral to mineral. The cation exchange ca-pacity (CEC) is a parameter that is used to quantifyand compare the ion exchange capacity of differentsoils. CEC is defined as milli-equivalents of cationthat can be exchanged per unit dry mass of thesoil. The procedure for testing of CEC involves firstrinsing the soil sample with ammonium acetatesolution, a process that saturates all the exchange-able sorption sites with NH4

+, after which thesample is flushed with an NaCl solution resultingin Na+ exchanging with NH4

+. The degree of theexchange is quantified by measuring the amountof NH4

+ flushed out from the sample. Becauseion exchange processes vary from ion to ion, andalso with pH as well as the concentration of thesolution, the measured CEC may be of limited usein predicting cation sorption under a variety ofconditions.

6.1.12 Nonpolar organic compounds

Many important contaminants, including hydro-carbon compounds and chlorinated solvents, haverelatively nonpolar distribution of electric chargeon molecules. Because of its polar nature andstrong electrostatic attraction between its differentmolecules, water tends to exclude nonpolar

P1: SFK/UKS P2: SFK



molecules from aqueous solutions. Therefore,nonpolar molecules are also called hydrophobic.Hydrophobic molecules have finite but very lowaqueous solubility, which tends to be lower forlarger and more perfectly nonpolar molecules.

The tendency of water to exclude nonpolarsolute molecules causes them to accumulate onthe surfaces that repel nonpolar molecules to alesser degree than water, giving a false impressionof being sorbed onto these surfaces. Hydrophobicsorption is quite different from the chemicalbonding and electrostatic attraction involved inion sorption. But the net effect is the same: sorbednonpolar molecules get immobilized and removedfrom flowing water.

6.2 Natural hydrochemical conditions in thesubsurface

Changes in the natural groundwater quality starttaking place immediately after its entry into top-soil, where infiltrating rainwater dissolves carbondioxide produced by the biological activity takingplace in the topsoil, which results in higher partialpressure of CO2 in it, producing weak carbonicacid. This process assists removal of soluble min-erals from the underlying such as calcite cements.Simultaneously, soil organisms consume some ofthe dissolved oxygen in the rainfall. In temperateand humid climates with significant recharge,groundwater moves relatively rapidly through theaquifer. Due to its short contact time with therock matrix, groundwater in the outcrop areasof aquifers is likely to be low in overall chemicalcontent, i.e. it has a low content of major ions andlow TDS. Igneous rocks usually have less solubleconstituents than sedimentary rocks (Hem 1985).

In the recharge area of the aquifer, oxidizingconditions occur and dissolution of calcium andbicarbonate dominates. As water continues tomove down gradient, a sharp redox barrier iscreated beyond the edge of the confining layer,corresponding to complete consumption of thedissolved oxygen. The bicarbonate concentrationincreases and the pH rises until buffering occurs atabout pH 8.3. Sulphate concentration remains sta-ble in the oxidizing water but decreases abruptly

just beyond the redox limit due to sulphatereduction. Groundwater becomes steadily morereducing while it flows in the aquifer down gradi-ent, as demonstrated by the presence of sulphide,increase in the solubility of iron and manganese,and denitrification of nitrate present in water. Aftermoving through the aquifer for a few kilometres,sodium concentration begins to increase at the ex-pense of calcium, due to the ion exchange process,causing a natural softening of the water. Eventually,the available calcium in the water is exhaustedbut sodium continues to increase to a level higherthan what could be accounted for purely by cationexchange. The point at which chloride begins toincrease is indicative of recharging water movingslowly through the aquifer and mixing with themuch older saline water present in the sediments(Fig. 6.4). The hydrochemical characteristics ofgroundwaters can thus be interpreted in terms ofoxidation/reduction, ion exchange, and mixingprocesses.

In arid and semi-arid environments, evapo-transpiration rates are much higher, recharge islower, flow paths longer, and residence timesmuch greater and hence much higher levels ofnatural mineralization, often dominated by sodiumand chloride, can be encountered. Thus in aridand semi-arid regions, the levels of major ionconcentration and TDS are often high. In somedesert regions, even if groundwater can be found,it may be too salty (very high TDS) for it to bepotable and the difficulty of meeting even thebasic domestic requirements can have seriousimplications on health and livelihood.

In many tropical regions, weathered basementaquifers and alluvial sequences have low pH,and the reducing conditions that can promotemobilization of metals such as arsenic, and otherconstituents of concern to health. Thus the pre-vailing hydrochemical conditions of groundwaterthat are naturally present at a given point andget modified as it flows in the aquifer, need tobe taken into account for: (i) developing schemesfor groundwater abstraction for various uses andin protecting groundwater reservoir; and (ii)considering transport and attenuation of additionalchemicals getting into groundwaters due to humanactivity.

P1: SFK/UKS P2: SFK



Aquitard

Increasing reducing conditions and salinity down gradient

Solution reactionsOxidisingconditions

Calcium bicarbonatewater

Sodium bicarbonatewater

Ion exchange

Sulphate reduction

Methanogenesis

Sodium chloridewater

Surface waterConfined aquifer

Con nir

fi ng st ata

Upflow

Fig. 6.4 Schematic representation of down-gradient change of hydrochemical facies in a typical regional groundwatersystems. Redrawn from Rivett et al. (2006).

6.3 Presenting inorganic chemical data

Typically analytical chemistry data are reportedin tables and numbers. However, when there arelarge numbers of analyses, these may be difficult tocomprehend and interpret. Graphical methods ofdata presentation are, therefore, helpful for quickinspection and identification of general trends inthe data.

Some of these graphical procedures are usefulmainly for visual display purposes, for example,in audio-visual presentations or written reportson water quality to provide a basis for com-parison of different analyses or to emphasizedifferences/similarities amongst them. Graphicalprocedures are more effective in this regard thanthe raw data presented in the form of tables. Inaddition, graphical procedures have been devisedto help detect and identify mixing of waters of dif-ferent origins/compositions and to identify someof the chemical processes that may take place asnatural waters circulate through different environ-ments. Some of the commonly employed graphicalmethods are presented here. Hem (1985) providesa more detailed account of these. Also one shouldnot lose sight of the fact that the graphical methodof presenting water quality data is a tool only toidentify broad trends and not an end in itself.

Methods of graphical analyses are generally de-signed to simultaneously present the total soluteconcentration and the proportions assigned to each

ionic species for a single analysis or a group ofanalyses. The units in which solute concentrationsare generally expressed in these diagrams are milli-equivalents per litre (meq l−1).

The chemical data of a limited number ofsamples can be depicted using bar charts or piediagrams, which can be easily made by usingdatabase management and spreadsheet software.In bar charts (e.g. in Collins’ ion-concentrationgraphical procedure; Collins 1923), each analysisis represented by drawing a vertical bar graphwhose total height is proportional to the measuredconcentration of anions or cations, expressedin meq l−1. The bar is divided into two partsby a vertical line with the left half representingthe cations and the right half the anions. Thesesegments are then divided by horizontal lines toshow concentrations of major ions. These ionsare identified by a single distinctive colours orpatterns. Usually six divisions are made but morecan be made if necessary. The concentrations ofclosely related ions are often added together andrepresented by a single pattern. A perfect chargebalance results in columns of equal height.

The Collins’ system, as described above, doesnot consider non-ionic constituents but they maybe represented, if desired, by adding an extra baror some other marker along with a supplementaryscale. In Fig. 6.5a, the hardness of two waters isshown. On a pie chart, anions and cations are plot-ted in opposite hemispheres and subdivisions of

P1: SFK/UKS P2: SFK



LEGEND

Na+K Cl+NO3

Mg SO4

Ca HCO3

Hardness

0

2

4

6

8

10

0

100

200

300

Mill

ieq

uiv

qle

nts

per

litr

e (m

eq/l)

Har

dn

ess

as C

aCO

, (m

eq/l)

3

(a)

Ca

HCO3

Mg

SO4

Na+K

ClCl

SO4

HCO3

Ca

MgNa+K

MgCa

HCO3

Na+K

ClSO4

Ca

HCO3

Na+KMg

SO4

Cl

Scale of Radius(meq/l)

(b)

Fig. 6.5 (a) Bar charts and (b) Pie charts representing arbitrary water analyses. Units are milli-equivalents per litre (meq l−1).The Collins’ system of ion-concentration diagrams for bar charts does not consider non-ionic constituents, but if desired, theymay be represented by adding extra bars with a supplementary scale as for hardness, as shown in (a). In pie diagrams, thescale of radii for the plotted graphs of the analysed data is so chosen as to make the area of the circle represent the total ionicconcentration, as shown in (b). Sample codes are indicated above the bars or circles. Redrawn from Hem (1985). © U.S.Geological Survey.

the area represent proportions of different ionsso that with perfect charge balance, anions andcations occupy half the circle each. A pie diagramcan be plotted with the radii drawn to scale, pro-portional to the measured total concentration ofdifferent ions (Fig. 6.5b).

Another graphical method is the Stiff diagram(Stiff 1951), as shown in Fig. 6.6a. Polygons aredrawn by plotting vertices at scaled distances to theleft (cation) and the right (anions) of a central axis.Waters of different origins form polygons of differ-ent shapes in such plots. Normally only three par-

allel horizontal axes are used for plotting concen-trations of three major cations and anions. A fourthaxis for Fe2+ and CO3

2− ions is often added at thebottom when these ions are present in significantconcentrations. The ions are plotted in the samesequence. The width of the polygonal pattern is anapproximate indicator of the total ionic content.

Bar charts, Pie charts, and Stiff diagrams arepractical procedures for visual comparison of asmall number of samples. A nomographic proce-dure proposed by Schoeller (1935), using verticallyscaled axes for individual cations and anions, is a

Cations Anionsmeq/l

0 5510 10

Ca

Mg

Na+K

Fe

Cl

HCO3

SO4

CO3

Stiff Diagram(a)

ClHCO3SO4 Na+KCaMg1

2

5

10

Schoeller Diagram(b)

Fig. 6.6 (a) Stiff diagram for presenting dataof major ion chemistry. Three or four parallelhorizontal axes may be used. Each sample isrepresented by a polygon and samplenumbers may be added below each polygon.Waters of differing origins revealdifferent-shaped polygons in such plots. (b)Schoeller diagram, which is a means ofdepicting groups of analyses. In this plot,waters of similar composition plot as nearparallel lines. Different symbols and types oflines can be used to label and/or identifydifferent samples.

P1: SFK/UKS P2: SFK



Cations Anions% Total meq/l

Ca

80

60

40

2080

60

40

20

40 206080

Mg

Na+K

A-1

A-6

A-3

A-10

80

60

40

20

CO

+HC

O3

3

4020

80

60

40

20

60 80Cl

A-1A-6A-3

A-10

80 80

60 60

40 40

20 20

SO+C

l4

Ca+M

g

A-1

A-6

A-3

A-10

0 100

500

1000

5000

1000

0

Scale of diameters

Total dissolved solids(meq/l)

SO4

Fig. 6.7 Trilinear diagram to depict major ion chemistry of four arbitrary samples with sample numbers marked on thefigure. The dotted lines show how the data for sample A-1 are projected from the two triangular plots to the diamond-shapedplot. The circles plotted in the diamond field have areas proportional to total concentration in meq l−1 for each sample andthe scale of diameters is shown in the upper left corner. Redrawn from Hem (1985). © U.S. Geological Survey.

means of depicting a group of analyses (Fig. 6.6b).As with the Stiff diagram, different Schoeller plotsshould have the same sequence of cations andanions for comparison. In this diagram, waters ofsimilar composition plot as nearly parallel lines.This diagram, however, uses logarithmic scalesand this may complicate interpretation of watersthat differ considerably in composition.

A trilinear or Piper diagram (Fig. 6.7) is aconvenient method of visualizing the results of alarge number of analyses in a single plot (Piper1944). The percentage values of total cations (inmeq l−1) is plotted in the lower left triangle, usingCa2+, Mg2+, and (Na+ + K+) on the three axes ofthe triangle. For example, a sample in which Ca2+

is the only cation present would plot at the lowerleft vertex of the triangle. In a similar way, anions

plot in the lower right triangle using percentagemeq l−1 of (HCO3

− + CO32−), SO4

2−, and Cl−. Thediamond-shaped part between the two trianglesshows projections from the anion and cationtriangles onto a field that shows percentage valuesof major anions and cations simultaneously. Fig. 6.7shows four arbitrary samples plotted on a trilineardiagram. Waters are often classified on the basisof where they plot on a trilinear diagram; such aclassification is called the hydrochemical facies.For example, Sample No. A-1 would be classifiedas a calcium–bicarbonate facies water. Otherfacies include sodium–chloride, sodium–sulphate,calcium–sulphate, etc. To enhance the visualimpact of a Piper diagram, different ranges of TDSmay be represented by symbols of different sizesor colours.

P1: SFK/UKS P2: SFK



A simple application of the Piper diagram may beto show whether particular water may be a mixtureof different waters for which analyses are availableor whether it is affected by dissolution or precip-itation of a single salt in the same water. Analy-sis of mixture of waters A and B should plot onthe straight line joining A and B in the plottingfield, provided the ions in A and B do not reactchemically on mixing. If solutions A and C define astraight line pointing towards the NaCl vertex, themore concentrated solution could form from a di-lute solution spiked by addition of sodium chloride.Plotting of samples from wells successively locatedin the direction of the hydraulic gradient may showlinear trends and other relationships that may havesome geochemical implications.

6.4 Impact of human activities

All life forms interact with their environments invarious ways. As a result of this, complex ecologicstructures have evolved in which diverse life formsinteract with one another in supportive as well asin predatory ways. One might view the present-dayEarth-surface environment as having been shapedin many ways through the biological processes in-teracting with their surroundings over the long ge-ological time span. Lovelock and Margulis (1974)suggested that metabolic processes of various or-ganisms have resulted in maintaining the compo-sition of the atmosphere at an optimum level forsurvival of various life forms on the Earth.

Many human activities have, however, adverseimpacts on the environment. When these activitiescause deterioration in the quality of natural waters,the result is ‘water pollution’. Water pollution maybe defined as the human-induced deterioration ofwater quality that is sufficiently severe to decreasethe usefulness of the resource substantially, eitherfor human beings or other life forms.

A serious consequence of human impact on wa-ter supply is the contamination and resulting pol-lution of the water source from diverse activitiesgenerating waste products such as: (i) industrialwastes, including mining, petrochemicals, etc.; (ii)agricultural non-point sources of pollution; (iii)wastewater treatment outfalls; (iv) accidental spills

and oil slicks; and (v) loading by sediments andnutrients. A point source is an identifiable source,such as a leaking septic tank, which may resultin a well-defined plume of the pollutant emanat-ing from it. On the contrary, non-point sources aremore difficult to control and pose a greater riskto water quality. Non-point sources are larger inscale compared to point sources and produce rela-tively diffuse pollution plumes originating from ei-ther widespread application of polluting material ina given area or caused by a large number of smallersources. The aggregate of point sources in a leak-ing sewerage system may be taken as an exampleof a non-point source of contamination to ground-water (usually described as multi-point pollutionsource).

Pollution of water sources ranges from smallstreams, rivers, lakes, and reservoirs to coastal wa-ters and even groundwater. Pollution of streams,rivers, and groundwater can form plumes of a pointsource downstream. Common water pollutants andtheir sources are listed in Table 6.12.

Contamination of groundwater can occur froma number of potential sources such as: (i) leakageof liquid fuels and chemicals from undergroundstorage tanks; (ii) septic tanks or cesspools usedfor disposal of domestic sewage or other wastew-ater; (iii) leachates from landfill refuse dumps orsewer pipes; (iv) injection wells disposing off liq-uid wastes in the subsurface; (v) pesticides, herbi-cides, and fertilizers sprayed on agricultural fieldsas part of farming operations; (vi) leachates frommine workings and tailings; (vii) spraying of saltfor de-icing the snow/ice-covered roads in winter;and (viii) sea water intrusion in coastal areas due toexcessive pumping of groundwater in the coastalaquifers. Groundwater contamination can also af-fect surface water and soil vapours.

Much of the subject of water-pollution controland problems relating to it are beyond the scope ofthis book. However, the basic principles of naturalwater chemistry can be applied to understand, pre-dict, and remedy/mitigate pollution problems. Itshould be appreciated that environmental changeattributable to anthropogenic factors is to a largeextent unavoidable and that some deterioration inwater quality may have to be accepted if the avail-able alternatives entail unacceptable social costs.

P1: SFK/UKS P2: SFK



Table 6.12 Common pollutants in water and their origin (adapted from http://www.epa.vic.gov.au/students/water/pollutants.asp)

Pollutant Origin

Sediments � Land surface erosion� Pavement and vehicle wear and tear (tyres, brakes, etc.)� Atmospheric particulate matter� Spillage/illegal wastewater discharge� Organic matter (e.g. leaf litter, grass, bird and animal excreta)� Waste water generated from washing of cars and other vehicles� Weathering of buildings/structures

Nutrients � Organic matter� Fertilizers� Sewer overflows/leaks from septic tanks� Animal/bird droppings� Detergents (household laundry/car washing)� Atmospheric particular matter� Accidental spillage/illegal wastewater discharges

Oxygen demanding substances � Decaying organic matter� Atmospheric emissions of pollutants, e.g. CO2, SO2� Sewer overflows/leaks from septic tanks� Animal/bird droppings� Spillage/illegal discharges

pH (acidity) � Atmosphere� Spillage/illegal discharges� Decaying organic matter� Erosion of roofing material

Micro-organisms � Animal/bird droppings� Sewer overflows/ improperly designed septic tanks/ cesspools� Decaying organic matter

Toxic organics � Pesticides� Herbicides� Spillage/illegal discharges� Sewer overflows/leaks from septic tanks

Heavy metals � Atmospheric particulate matter� Vehicle wear and tear� Sewer overflows/leaks from septic tanks� Weathering of buildings/structures� Spillage/illegal discharges

Oils and surfactants � Asphalt pavements� Spillage/illegal discharges from automobiles/tankers� Leaks from vehicles� Washing of vehicles� Organic matter from various sources

High water temperatures � Runoff from impervious surfaces� Removal of riparian vegetation� Industrial discharges (e.g. power production)

P1: SFK/UKS P2: SFK



In recent years, as a result of improved technol-ogy and application of analytical chemistry, detec-tion of progressively lower levels of organic or in-organic pollutants has become relatively easy. Butthe mere presence of a particular substance doesnot establish the existence of pollution. Variousother aspects need to be considered. Major prob-lems that need in-depth scientific study include de-termination of form, stability, transport rates, andmechanisms for generation of pollutant species;predicting probable effects of current or foresee-able uture practices in waste disposal or productconsumption causing contamination; assessing rel-ative impacts of artificial sources vis-a-vis naturalsources; and providing methods for identifying themost significant existing as well as potential pollu-tion sources.

Some common ionic species may be dispersedinto the environment with no serious repercus-sions. Chloride, for example, may not be harmful ifmaintained at a low level of concentration. Further-more, chloride gets easily conveyed to the oceanwhere it causes no significant ill effects. Some othersolutes, however, may tend to accumulate and be-come concentrated in such places as sediments instreams or in biota and can be released from suchsources in unforeseen ways to cause troublesomelocal concentrations.

Although some polluted surface waters can berejuvenated to a reasonable quality fairly rapidly bydecreasing the waste loads or concentrations, theprocess can be expensive. A polluted groundwater,on the other hand, may be so slow to recover thatone may think of the pollution of aquifers as almostirreversible. For this reason, great care is needed toprotect groundwater, particularly deeper aquifers.Incidents of contamination of groundwater fromseptic tanks, sewage, and industrial waste-disposalsystems, solid-waste disposal practices, and naturalgas and petroleum storage leaks have been a grow-ing problem both in urban as well as rural areas.

In an industrial society with developed agricul-ture, a large number of organic and inorganicproducts are produced and used, which in thenatural course do not enter the environment. In-evitably, some of these products, their residues, orby-products enter the water cycle. During the pastfew decades, attention has been drawn, for exam-

ple, to the presence of lead, mercury, and variouscommon and exotic organic substances in the aque-ous environment. A matter of more recent concernworldwide is the occurrence of rainfall having alow pH – the so-called ‘acid rain’, which is harmfulboth for the plant as well as the animal kingdom.

6.4.1 Pathogenic micro-organisms

Waterborne diseases remain one of the ma-jor health concerns globally and waterbornepathogens from various sources, in particular agri-cultural and urban sources, are the main diseasecausing agents. Typically, the health risk fromchemicals is lower than that from pathogens. This isbecause, unlike pathogens, health effects many ofthe hazardous chemicals manifest after a prolongedexposure and tend to be limited to specific geo-graphical areas or particular types of water source.Following is a brief summary of the current knowl-edge on the distribution of pathogens in water ingeneral and groundwater in particular, and the fac-tors that control their transport and attenuation.

The ability of a pathogen – a group of disease-causing micro-organisms – to inflict damage uponthe host is controlled by a combination of factors,in particular the nature of the organism (e.g. the de-gree of its virulence) and the susceptibility of thehost. Water offers an easy carrier for transmission ofmost pathogenic micro-organisms, some being nat-ural aquatic organisms and others introduced intowater from an infected host. Overall, pathogens inwater that are the main concern to public healthoriginate in the faeces of humans and animals andcause infection when water contaminated with fae-cal matter is consumed by a susceptible host. Wa-terborne pathogenic micro-organisms of concernand the health effects caused by these are summa-rized in Table 6.13.

Furthermore, currently unknown and, therefore,undetectable pathogens may also be present inwater. To overcome this difficulty, a separategroup of micro-organisms is used as an indicatorfor the presence of potential pathogens. Thecommonly used term for this group of organismsis faecal indicator organisms. Gleeson and Gray(1997) published an exhaustive review of theapplication of faecal indicator organisms in water

P1: SFK/UKS P2: SFK



Table 6.13 Waterborne pathogenic micro-organisms and their associated health effects. Adapted from Macler and Merkle(2000).

Organism Associated health effects

VirusesCoxsackie virus Fever, pharyngitis, rashes, respiratory diseases, diarrhoea, haemorrhagic

conjunctivitis, myocarditis, pericarditis, aseptic meningitis, encephalitis, reactiveinsulin-dependent diabetes, diseases of hand, foot, and mouth

Echovirus Respiratory diseases, aseptic meningitis, rash, feverNorovirus (formerly Norwalk virus) GastroenteritisHepatitis A Fever, nausea, jaundice, liver failureHepatitis E Fever, nausea, jaundice, deathRotavirus A and C GastroenteritisEnteric adenovirus Respiratory diseases, haemorrhagic conjunctivitis, gastroenteritisCalicivirus GastroenteritisAstrovirus Gastroenteritis

BacteriaEscherichia coli Gastroenteritis, Haemolytic Uraemic Syndrome (enterotoxic E. coli)Salmonella spp. Enterocolitis, endocarditis, meningitis, pericarditis, reactive arthritis, pneumoniaShigella spp. Gastroenteritis, dysentery, reactive arthritisCampylobacter jejuni Gastroenteritis, Guillain-Barre syndromeYersinia spp. Diarrhoea, reactive arthritisLegionella spp. Legionnaire’s disease, Pontiac feverVibrio cholerae Cholera

ProtozoaCryptosporidium parvum DiarrhoeaGiardia lamblia Chronic diarrhoea

Note: The coliform group of bacteria comprises several genera belonging to the family Enterobacteriaceae (Gleeson and Gray1997). The relatively limited number of biochemical and physiological attributes used to define the group means that itsmembers include a heterogeneous mix of bacteria. Although the total coliform group includes bacteria of faecal origin, it alsoincludes species that are found in unpolluted environments. Thermo-tolerant coliforms are those bacteria from within thecoliform group that grow at 44◦C. E. coli is a thermo-tolerant coliform bacteria.

quality monitoring. This group comprises: (i) totalcoliform bacteria; (ii) thermo-tolerant coliformbacteria; (iii) E. coli; (iv) faecal streptococci; and(v) bacteriophages.

6.4.1.1 Transport and attenuation ofmicro-organisms in the subsurface

Many instances of groundwater contaminationhave occurred, possibly by rapid transportpathways accidentally introduced by humanintervention and connecting the contaminationsource to the groundwater abstraction point.Such pathways could be provided by impropercompletion of springs, wells, and boreholes, andconduits left carelessly connecting the source of

contamination to the groundwater abstractionpoint/system. Existence of rapid transport path-ways cannot, however, explain all occurrences ofgroundwater source contamination and it is nowwidely accepted that the transport of microbialpathogens through groundwater systems is an im-portant mechanism for transmission of waterbornediseases.

6.4.1.1.1 Unsaturated zone

Hydrogeological processes in the unsaturated zoneare complex and the behaviour of micro-organismsis often difficult to predict. Nevertheless, theunsaturated zone can play an important role in re-tarding (and in some cases eliminating) pathogens.

P1: SFK/UKS P2: SFK



Attenuation of pathogens is generally moreeffective in the top soil layers where significant mi-crobial activity occurs. Within this zone, presenceof protozoa and other predatory organisms, rapidchanges in moisture content and temperatureof the soil, competition from the establishedmicrobial community, and the effect of sunlight atthe surface combine together to reduce the levelof pathogens.

Movement of pathogens from the surface intothe subsurface requires presence of moisture. Evenduring relatively dry periods, soil particles on thesurface retain sufficient moisture for pathogens tomigrate. Within the thin film of moisture the organ-isms present are brought into close contact withthe surface of particulate matter, thus increasingthe opportunity for adsorption onto the surfacesof particulate matter, which further retards theirmovement. If the soil moisture content decreases,the strength of association between organisms andparticle surfaces will increase to a point where or-ganisms are attached irreversibly to particle sur-faces. In laboratory experiments, soil moisture con-tent between 10 and 15% has been shown to be op-timal for survival of several strains of enteric viruses(Bagdasaryan 1964; Hurst et al. 1980a, b).

Increase in the moisture content of the unsatu-rated zone, on the other hand, may increase thevulnerability of the aquifer to pathogen contami-

nation in two ways by: (i) providing rapid trans-port pathways; and (ii) remobilizing the adsorbedmicro-organisms. During periods of high ground-water recharge, for example during prolongedheavy spells of rain, the inter-granular spaces inthe unsaturated zone become waterlogged, pro-viding hydraulic pathways for rapid transport ofpathogens.

In the interval between individual rechargeevents, the chemistry of water in the unsaturatedzone changes as it equilibrates with the soil matrix.In some soil types, these changes may favour ad-sorption of micro-organisms onto surfaces in thesoil matrix. A lowering of the ionic strength or saltcontent of the surrounding medium, which can oc-cur during a rainfall event may, however, be suffi-cient to cause desorption of the organisms allowingtheir further migration into the soil matrix. Vari-ability in the size of different micro-organisms can,to some extent, control their mobility in the sub-surface. Pore sizes of soil and rock particles arealso variable and the size ranges of the two areknown to overlap (Fig. 6.8). Thus in soils that arecomposed of fine-grained particles, typically clayeysilts, the pore space is sufficiently small (<4 µm)to physically prevent the passage of pathogens,such as bacteria and protozoa, into the subsurface.This removal process is called physical filtration orstraining.

Fig. 6.8 Comparison of pathogen size diameter with pore size diameter in an aquifer matrix. Adapted from Pedley et al.(2006).

P1: SFK/UKS P2: SFK



6.4.1.1.2 Saturated zone

Shallow groundwater (<5 m deep) has the highestprobability of contamination, irrespective of thelithology of the unsaturated zone. As depth tothe water table increases so does the capacity ofthe unsaturated zone to attenuate micro-organisms,although this also depends upon composition andstructure of the unsaturated zone. On reachingthe saturated zone, microbial contaminants aresubject to the same processes of attenuation asin the unsaturated zone but under conditions ofnatural or artificially induced groundwater flow.Thus death, adsorption, filtration, predation, anddilution contribute to attenuation of pathogens inthe saturated zone. However, there may be largevariations in hydraulic conductivity of an aquiferdepending on the nature of aquifer material, whichcan significantly influence mobility of micro-organisms in the aquifer. Micro-organisms aretransported through groundwater by advection,dispersion, and diffusion, as defined in Section7.6. It may be recalled that advection refers totransport of dissolved solute mass due to bulk flowof groundwater and to a non-reactive solute forwhich the advection velocity would equal the bulkgroundwater velocity through pores. Diffusionand dispersion, arising from the tortuosity of flowwithin an aquifer, cause mixing of a solute andmodify its concentration as the flow proceeds.The result of both advection and dispersion ismigration and spreading of the contaminant in theaquifer, resulting in a decrease in its concentrationboth in space and time. This may result in contam-ination of increasingly large volumes of the aquiferas the pollutant moves down gradient with thegroundwater. Although transport of pathogens insome aquifer types can be both rapid and exten-sive, there are several factors that may attenuatepathogens in groundwater (West et al. 1998).

Inactivation rates of bacteria and viruses ingroundwater vary considerably, not only betweendifferent groups of bacteria and viruses but alsobetween different strains within each group. Thereis also variation between the results of differentmeasurements. Usually inactivation proceeds fasterat higher temperatures, although highly dependenton type of micro-organism. Often, inactivation

of micro-organisms can be described reasonablywell by a first-order rate process, especially underrelatively mild conditions, such as temperaturesbetween 5 and 20◦C and pH values in the range 6to 8. Thus:

Ct = C0 e−µt or ln

(Ct

C0

)= −µt or

(6.45)

log10

(Ct

C0

)= − µ

2.3t

where Ct is the concentration of micro-organismssurviving after time t; C0 is the initial concentrationat t = 0; and µ is the inactivation rate coefficient[T−1]. For ease of interpretation, µ is often dividedby a factor of 2.3 (equal to the natural logarithm of10). The inactivation rate coefficient then reflectsthe number of log10 units per unit time; forexample, a decrease in virus count every 10 daysby 2 units of log10 is equal to reduction by a factorof 100.

Under more extreme conditions, the rate of in-activation of a virus, e.g., is often found to proceedinitially at a higher rate followed by a lower rate,as if two or more sub-populations differing in theirstability exist simultaneously (Hurst et al. 1992).

6.4.1.2 Subsurface transport and attenuationof chemicals

In the classical conceptual model of contaminantfrom a point source to a receptor, the chemicalcontaminant is leached from a near-surface leach-able source zone through the process of dissolutionin the infiltrating water (Fig. 6.9a). Subsequently,a dissolved phase chemical solute plume emergesin water draining from the base of the contami-nant source zone and moves vertically downwardsthrough the unsaturated zone. This solute plume ul-timately reaches the water table and subsequentlymigrates laterally in the aquifer along with ground-water flow. If the source (e.g. landfill, chemicalwaste lagoon, contaminated soils from industrialsites, pesticide residues in field soils, etc.) hassufficient pollutant mass, the contaminant plumecan last for decades. In this manner, continuousdissolved phase plumes extending from the sourcethrough the groundwater pathway can grow withtime and may eventually reach a distant location.

P1: SFK/UKS P2: SFK



Fig. 6.9 (a) A conceptual model for transportation and attenuation of chemical and hydrophilic organic pollutants ingroundwater starting from a point source up to the location of the observer. (b) Polarity-volatility diagram for select organiccontaminants. Redrawn from Rivett et al. (2006).

The near surface leachable source–dissolvedplume conceptual model as described above isfrequently invoked for vulnerability and protectionof groundwater and also for groundwater riskassessment. It is important to note, however, thatthe above conceptualization may be too simplisticand alternative conceptual models may be neededin some cases, notably for non-aqueous phaseliquids (NAPLs).

The processes involved in retardation of a plumemovement are: (i) sorption by which chemicals ororganisms become attached to the aquifer solids;(ii) cation exchange by which interchanging ofcations occurs in solution with those on the sur-faces of clay particles or organic colloids; and (iii)filtration that affects relatively large-size particulatecontaminants by preventing their movement byadvection.

In addition, abiotic (i.e. not mediated by bacteria)reactions such as precipitation, hydrolysis, com-plexation, elimination, substitution, etc. that trans-form the chemicals present into some other chem-icals and potentially alter their phase/state (solid,liquid, gas, dissolved) are also involved in trans-

portation and attenuation of contaminant plumesin the underground.

Biodegradation (biotic reactions) is a reactionprocess that is facilitated by microbial activity, forexample, by bacteria present in the subsurface.Typically contaminant molecules are degraded(broken down) to molecules of simpler structurethat often have lower toxicity.

6.4.2 Organic compounds

Natural sources invariably contribute some organiccompounds to water, though at low levels. Nat-ural organic matter comprises water-soluble com-pounds of a rather complex nature having a broadrange of chemical and physical properties.

Typically, natural organic matter in surface- andgroundwater is composed of humic substances(mostly fulvic acids) and non-humic materials(e.g. proteins, carbohydrates, and hydrocarbons)(Stevenson 1994; Thurman 1985). Natural organicmatter is a complex, heterogeneous mixturevarying in size, structure, functionality, and re-activity. It can originate from terrestrial sources

P1: SFK/UKS P2: SFK



(allochthonous natural organic matter) and/oralgal and bacterial sources within the water body(autochthonous natural organic matter). Dissolvedorganic carbon (DOC) is considered to be asuitable parameter for quantifying the organicmatter present in groundwater. However, DOCis a bulk organic quality parameter and does notprovide specific identification data and may alsoincorporate organic compounds resulting fromhuman activity. Natural organic matter, althoughconsidered innocuous, may still indirectly influ-ence water quality. For example, contaminantsmay bind to organic-matter colloids, facilitatingtheir transport with water flow.

Also, routine chlorination of water supplies con-taining natural organic matter may form disinfec-tion by-products such as trihalomethanes that areknown to be carcinogenic. However, natural or-ganic substances are of little direct health concernand, therefore, are not addressed any further here.

Human activity has contributed to the release ofa vast range of anthropogenic organic chemicalsto the environment; some of these may adverselyimpact groundwater quality. In the following, thefocus is specifically on the subsurface transport ofnon-aqueous phase liquids as part of the anthro-pogenic suite of organic chemicals.

6.4.2.1 Subsurface transport models fornon-aqueous phase liquids

As already mentioned (see Section 6.1.12), due tothe polar nature of water molecules, it causes non-polar molecules to be excluded from aqueous so-lutions – called hydrophobic as opposed to po-lar molecules that are hydrophilic. Hydrophobicmolecules do have finite but very low aqueous sol-ubility, which tends to be lower for larger and moreperfectly nonpolar molecules.

Hydrophobic non-ionic organic contaminantspreferentially sorb onto the low polarity com-ponents of geo-solids, for example, any organicmaterial present. Sorption is inversely related tothe solubility of an organic compound; the morehydrophobic and less soluble an organic solute is,the greater its intrinsic potential for sorption ontoany organic material present in the aquifer solids.Hence hydrophilic organics have negligible sorp-

tion and mild to moderately hydrophobic organics,such as the volatile organic compounds (VOCs),which show limited sorption. In contrast, hy-drophobic, high molecular weight, large organicssuch as polynuclear aromatic hydrocarbons (PAHs)and polychlorinated biphenols (PCBs) of low sol-ubility exhibit high sorption. Hydrophobicity andsorption retardation (indicative of solubility) trendsof select organic contaminants against vapour pres-sure (a measure of volatilization tendency) areshown in Fig. 6.9b in the form of a polarity-volatility diagram. It is seen that PAHs are unlikelyto volatilize and undergo a high degree of sorptionand this also implies that in soil/unsaturated zonesolids, concentrations of PAHs could often be high,which is frequently the case. There is relatively lim-ited development of PAH plumes in groundwater;often in real-life situations relatively small plumesare encountered. The chlorinated hydrocarbons,in contrast, are volatile but with low sorptionpotential. It is likely that they get vaporized, andpotentially become a vapour hazard to living organ-isms at the soil surface. Chlorinated hydrocarbonsare also leached into the groundwater, resulting intheir low concentrations in soils and unsaturatedsamples, which is quite often the case.

The classical near-surface leachable sourcezone dissolved plume model (presented in the Sec-tion 6.4.1.2; Fig. 6.9a) is not applicable to all or-ganic substances. Of key importance is the realiza-tion that different organic chemicals have differentaffinities for water, ranging from hydrophilic to hy-drophobic.

Most organic liquids, however, are so hydropho-bic that they form a separate organic phase withinthe water (aqueous) phase. They are immisciblewith water and a phase boundary exists betweenthe organic phase and the aqueous phase. The or-ganic phase is generally referred to as the non-aqueous phase liquid (NAPL). When a separate or-ganic NAPL exists, it is important to consider thedensity of the NAPL relative to water, as this con-trols whether the NAPL will be above or below rel-ative to the water phase. Most hydrocarbon-basedorganic liquids have a density <1 g cm−3, for ex-ample, the density of benzene is 0.88 g cm−3 andthat of pentane is 0.63 g cm−3. When in contactwith water, these liquids form the upper phase and

P1: SFK/UKS P2: SFK



‘float’ on the water phase, which has density of 1g cm−3. Such ‘light’ organic compounds are gener-ally referred to as LNAPLs.

In contrast, other hydrophobic organics havea relatively high density due to incorporation ofdense chlorine (or other halogen) atoms in theirstructure. For example, chlorinated solvents suchas trichloroethene (TCE), 1,1,1-trichloroethane(1,1,1-TCA), and polychlorinated biphenyl (PCB)mixtures have densities in the range 1.1–1.7 gcm−3. Due to their higher density, such organicphases will form the lower phase and ‘sink’ belowthe water phase. Such ‘dense’ organic compoundsare generally referred to as DNAPLs.

Although hydrophobic, LNAPL and DNAPL or-ganics still have the potential for some of their or-ganic molecules to dissolve into the adjacent aque-ous phase. These organics are ‘sparingly soluble’and have a finite solubility in water, leading to theirfinite dissolved concentrations in the water phase.

In contrast to hydrophilic miscible organics, hy-drophobic immiscible organics, that is NAPLs, ex-hibit quite different behaviour. Conceptual modelsfor LNAPL and DNAPL releases (Mackay and Cherry1989) are shown in Fig. 6.10.

If sufficient pressure head exists to overcomethe entry pressure to the pores or fractures, NAPLsmay migrate as a separate phase and displace airand water from the pores into which they pene-trate. NAPL migration is also controlled by its den-sity and viscosity. For example, petrol and chlori-nated solvents have viscosities lower than waterand migrate more easily in the subsurface. In con-trast, PCB oils or coal tar (PAH-based) hydrocar-bons may be very viscous and perhaps take yearsfor the NAPLs to come to a resting stable posi-tion in the subsurface. Chlorinated solvents, suchas PCE, have high densities and may penetrate tosignificant depths through aquifer systems in veryshort time periods. Whereas dissolved pesticidesmay take a couple of years to decades to migratethrough a 30 m unsaturated zone, DNAPLs maymigrate through such a zone in the matter of afew hours or days (Pankow and Cherry 1996). Atthe water table, LNAPLs, being lighter than water,form a floating layer on the water table, often form-ing a slightly elongated plume in the direction ofthe hydraulic gradient of the water table. DNAPLs,in contrast, may penetrate as a separate immisci-ble layer below the water table. Its predominant

LNAPL

ResidualLNAPL Vapours

LNAPL ‘Pancake’Vapours

Vapours

Dissolved-phasesolute plume

Groundwater

Flow

Aquifer base

(a)

Groundwater

Flow

Aquifer base

Dissolved-phase

solute plume

Vapours

ResidualDNAPL

DNAPLpool

Clay layer

DNAPL

(b)

Fig. 6.10 Conceptual models of: (a) light non-aqueous phase liquid (LNAPL) release; and (b) dense non-aqueous phase liquid(DNAPL) release. Redrawn from Rivett et al. (2006).

P1: SFK/UKS P2: SFK



movement in a vertical (downward) direction isdue to its density, but some lateral spreading doesoccur as it encounters a lower permeability strata.If sufficient volume and pressure head exist, theDNAPLs may penetrate the entire aquifer depthup to the underlying aquitard/bedrock (Kueperet al. 1993). Migrating NAPLs leave behind a trailof immobile residual NAPL droplets along their mi-gration pathways held by capillary forces causingNAPLs to spread across the entire aquifer thick-ness. DNAPLs accumulating on low permeabilityfeatures, often referred to as pools, are potentiallymobile. They may ultimately penetrate the forma-tion due to changes in hydraulic pressure arisingfrom their continued spillage, pumping, or reme-diation attempts or during drilling (for boreholes,piling, etc.) through the layer. Often NAPLs remainrelatively localized at the site.

Risks posed to groundwater resources and sup-plies are generally concerned with migration ofthe dissolved-phase plume formed through contactof flowing groundwater with spilled NAPLs. Oftenthe mass of NAPLs is so large and the dissolutionof NAPLs into water so slow that the entire NAPLbody subsequent to spilling should be regarded as alargely immobile source zone that can continuouslygenerate a dissolved-phase solute plume of organ-ics moving down-gradient for years to decades, oreven centuries for low-solubility NAPLs. In general,DNAPLs tend to pose the greatest threat to ground-water as they reside deep in groundwater systemsand many, being chlorinated, are less susceptibleto attenuation. In contrast, LNAPLs are restricted toshallower groundwater table depths and are moresusceptible to attenuation via biodegradation.

6.5 Geochemical modelling

An important issue that one is often confrontedwith relates to understanding the reactions thatmay occur along a flow path – given the parametersthat are measured by aqueous geochemistry. Thisis essentially the process of inverse geochemicalmodelling vis-a-vis forward modelling, wherein oneattempts to figure out the geochemical evolution ofsurface water when flowing over a given terrain orgroundwater flowing through an aquifer, and how

the model-computed results compare with obser-vations. Essentially, one has to look for tools tounderstand the evolution of chemical species inwater at specific sites, for example: (i) to see if cer-tain minerals would dissolve or precipitate in thewater; (ii) to model water–rock interaction (dis-solution, precipitation, ion exchange); and (iii) toestimate changes in chemical properties of waterduring mixing of different water masses.

6.5.1 Computer models

Many of the earlier computer models were de-signed for specific problems related to aqueousspeciation (Allison et al. 1991; Nordstrom et al.1979). Some important recent applications includemodelling the disposal of high-activity radioac-tive waste, environmental issues associated withmining operations, landfill leachates, injection ofhazardous chemical/biological wastes into deepwells, water resources-related issues, and artificialrecharging of aquifers, particularly deep aquifers(Zhu and Anderson 2002).

Existing models use the same basic approach,calculating thermodynamic equilibrium state of aspecified system that includes water, solutes, sur-faces, and solid and gas phases. These models com-prise four components:

1) Input: specific information that defines the sys-tem under consideration, such as concentra-tion of solutes, temperature, partial pressure ofgases, and composition of solid phases;

2) Equations that are to be solved by the model;3) Equilibrium and kinetic formulations between

solutes of interest;4) Output: in tabular or graphic form.

The computer codes require initial input con-straints that generally consist of water chemistryanalyses along with, units used for chemical mea-surements, temperature, dissolved gas content,pH, and redox potential (Eh). The models convertthe chemical concentrations, usually reported inwt./wt. or wt./volume units such as mg kg−1 ormg l−1 to moles, and then solving a seriesof simultaneous non-linear algebraic equations(involving chemical reactions, charge balanceand mass balance equations) to determine the

P1: SFK/UKS P2: SFK



activity–concentration relationship for all thechemical species in a specified system. The modelsusually require electro-neutrality conditions andimpose the charge balance condition with one ofthe designated components, as they solve a set ofnon-linear equations formulated to describe theproblem. The capabilities of modern codes includecalculation of pH and Eh, speciation of aqueousspecies, equilibration with gases and minerals,oxidation and reduction reactions (redox), kineticreactions, and reactions with surfaces.

The non-linear algebraic equations are solved us-ing an iterative approach by the Newton-Raphsonmethod (Bethke 1996). The equations to be solvedare drawn from a database that contains equationsin the standard notations of chemical mass action.In principle, any reaction such as sorption ofsolutes onto surfaces that can be represented inthis form can be incorporated into the model.Reactions are assumed to reach equilibrium state(the point of lowest free energy in the system)when there is no change in the concentration ofreactants and products.

Kinetic reactions, that involve time, are includedby assuming that chemical reactions proceed toreach the equilibrium at a specified rate. Examplesof kinetic reactions include mineral dissolution andprecipitation, redox reactions, microbial growth,and metabolism of solutes. The rate laws used inthe codes vary but all codes with kinetic capabil-ities include simple first-order rate laws and mayinclude more complex rate formulations such ascross-affinity, Michaelis-Menten, and Monod formu-lations (Bethke 1996).

Computer models are divided into two basictypes: speciation models and reaction-path mod-els. In both cases, the models are fundamentallystatic, that is, there is no explicit transport func-tion. However, some forms of transport can besimulated by manipulation of the models. Morecomplex reaction-transport models that explicitlyincorporate transport are briefly described below.

All equilibrium models are speciation models inthat they can calculate the speciation (distribution)of aqueous species for any element or compoundincluded in the database.

Speciation models calculate activities (chemi-cally reactive concentration), species distribution

for elements in the database, saturation indicesand ion ratios for specified conditions of pH, andredox potential (ORP or Eh). Most of the modelsallow choice of the method employed for activitycalculation (Davies, Debye-Huckle, extendedDebye-Huckle, Pitzer). Some models incorporatesurface reactions such as adsorption and multiplekinetic formulations. Only one model, PHREEQC,has provision for inverse modelling option. Thisfeature uses mass balance constraints to calculatethe mass transfer of minerals and gases that wouldproduce the final water composition, given aspecified starting water composition (Garrelsand Mackenzie 1967). This method does notmodel mass transport but calculates and providesonly statistical measures of fit to obtain possiblesolutions to the mass balance between starting andfinal water compositions.

The next step in the complexity is the reactionpath (mass transfer) models. These models usespeciation calculation as the starting point andthen make forward predictions of changes alongthe specified reaction path (specified changesin T, P, pH, and addition of new reactantssuch as another fluid or solid). The programmakes small incremental steps with step-wiseaddition or removal of mass (dissolution orprecipitation) and can also include changesin temperature or pressure along the reactionpath.

There are, however, limitations associated withany of the above models. The input field data maybecome corrupted due to bad analysis, missingparameters, or violation of electro-neutrality.Speciation models assume equilibrium conditions,which may not be the case in real-life situations.The databases are also a source of uncertainty. Theydo not always contain all the elements or species ofinterest, the data invariably has some uncertainty,and some data may be inaccurate (Drever 1997).Some of the available codes attempt to minimizethis problem by including popular databases suchas the MINTEQ database (EPA-approved databaseon metals), WATEQ (USGS database specifically onminerals), and the LLNL database (the most com-plete database available, which is compiled andmaintained by the Lawrence Livermore NationalLaboratory). For environmental applications, lack

P1: SFK/UKS P2: SFK



Table 6.14 A comparison of capabilities and features of select computer codes commonly used in geochemical modelling.Adapted from Thyne (2005).

Program Source SpeciationReactionPath

Tabularoutput

Graphicoutput

SurfaceRxns. Kinetics Inverse Transport

Multipledatabases

EQ3/6 LLNL yes yes yes no no yes no no noGWB Rock-

ware∗yes yes yes yes yes yes no yes yes

HYDRO-GEOCHEM2

SSG∗ yes yes yes no yes yes no yes no

MINTEQ EPA yes no yes no yes no no no noMINEQL+ ERS∗ yes no yes some no no no no noPHREEQC USGS yes yes yes no yes yes yes yes yes

∗ - Commercial programs, others are freeware.EQ3/6 – http://www.llnl.gov/IPandC/technology/software/softwaretitles/eq36.phpGWB – Rockware – http://www.rockware.comHYDROGEOCHEM – http://www.scisoftware.com/environmental software/software.phpMINEQL+ – http://www.mineql.com/MINTEQ – http://soils.stanford.edu/classes/GES166 266items%5Cminteq.htmPHREEQC – http://wwwbrr.cr.usgs.gov/projects/GWC coupled/phreeqc/

of adequate data for organic compounds remainsa matter of concern.

Other limitations include the redox reactionsthat are of particular importance in metal transport.These reactions are difficult to model correctlysince redox reactions may have different rates, pro-ducing natural systems that are not in equilibrium(Lindberg and Runnels 1984). This problem canbe addressed by modelling redox reactions as rate-limited (kinetic) formulations if appropriate dataare available.

Some of the commonly used programs in geo-chemical modelling, their sources, and some oftheir useful capabilities are listed in Table 6.14.

6.6 Chemical tracers

In many cases, water pollution is an indirect conse-quence of human activities in the sense that the pol-luting material is not being directly let out into thereceiving water as a waste effluent. One such formof indirect pollution of water is the intrusion of sea-water into the coastal aquifers. This may be consid-ered as an example in which both hydrologic andchemical knowledge are essential in controlling thepollution, and chemical species often act as tracers.

6.6.1 Seawater intrusion

Along the sea coast there is a saltwater–freshwatercontact zone, both in streams and in aquifers thatextend under the seabed. The relation between seawater and fresh water in aquifers along the sea coastis generally described by the hydraulic relationshipknown as the Ghyben-Herzberg equation (see Sec-tion 3.6.4, in Chapter 3; Eqn. 3.51). The effect of ex-cessive withdrawal of groundwater from the land-ward parts of these aquifers may have far-reachingeffects on the position of the saltwater–freshwaterinterface.

Hydrologists generally consider that the bound-ary between fresh water and salt water in coastalaquifers depends on the balance of forces in adynamic situation. Normally, there is a continuousseaward movement of fresh water at a rate thatis related to the head above mean sea level in thefreshwater aquifer. Cooper (1959) described themovement of fresh- and salt water along a contactzone, which tends to produce a diffuse zone ofmixing rather than a sharp interface as predictedby the Ghyben-Herzberg relation. So long as a highhead of fresh water is maintained inland in theaquifer, freshwater discharge maintains the zoneof contact with saline water in the aquifer up to a

P1: SFK/UKS P2: SFK



considerable distance offshore. Pumping the inlandaquifers reduces the head of the fresh water and ashead changes are transmitted rapidly to maintainthe hydrostatic equilibrium, the seaward flowof fresh water decreases. The head may declineaccordingly to completely stop the seaward flowof fresh water past the interface. With reducedfreshwater flow, the system becomes unstable andsalt water invades the aquifer. The saltwater frontmoves inland to the point where the reduced fresh-water head is sufficient to produce a balancing sea-ward movement of fresh water past the interface.

Overdevelopment of coastal aquifers can greatlydecrease the freshwater head and can bring aboutconditions favourable for the intrusion of salt wa-ter inland. The migration of the saltwater front,however, is rather slow, as it represents actualmovement of water in the system under low hy-draulic gradients against high resistance offered bythe aquifer material. The appearance of salty waterin a well may, however, not occur until some yearsafter the head decline in the coastal aquifers hasreached serious proportions.

The rate of movement of some of the ions in thesaltwater front is also influenced by ion exchange,and diffusion and head fluctuation will cause theinterface to become a broad transition zone ratherthan a sharp front as given by the Ghyben-Herzbergrelation. Saltwater intrusion into highly developedaquifers is a serious problem in many places alongcontinental margins. Hydrologists are frequentlyconfronted with the need to recognize incipientstages of saltwater intrusion so that steps can betaken to remedy the situation in time.

Chloride is the major anion in sea water, whichmoves through an aquifer at nearly the same rateas the intruding sea water. Increase in chlorideconcentrations with time may well be the firstindication of the onset of a seawater intrusionfront. In an area where no other known source ofsaline water contamination exists, high chlorideconcentrations in groundwater can be considereddefinite evidence for seawater contamination.However, if significant amounts of chloride arecontributed by other sources, establishment ofunambiguous proof of the seawater source may bedifficult. Components of sea water other than chlo-ride may also be used to identify contamination but

with some caution. Magnesium is present in seawater in much greater concentration than calcium.A low calcium/magnesium ratio may sometimes beindicative of seawater contamination. Presence ofsulphate in anionic proportions similar to that ofsea water may also be indicative. Because of possi-ble cation exchange reactions and sulphate reduc-tion in the aquifers that can be expected to occurwhen sea water is introduced, the proportions ofanions and cations in the first contaminated waterto reach the sampling point cannot be expected tobe exactly the same as those in a simple mixtureof sea water and fresh water. It is indeed likelythat, even after moving only for a short distancethrough an aquifer, the water in the advancingsaltwater front will have superficial resemblanceto a simple mixture of sea water and groundwater.An example of cation exchange altering the Ca/Nabalance in sea water intruded in the coastal karsticlimestone aquifer from part of Saurashtra coast ofGujarat, India, was reported by Desai et al. (1979).

In some instances, minor constituents of sea wa-ter may aid in determining whether a particularaquifer has been contaminated by sea water or bysome other saline source. Incipient stages of con-tamination are, however, difficult to detect by theseconstituents. Piper et al. (1953) reported some suc-cess in differentiating seawater contamination ofan aquifer from contamination by connate brineby comparing iodide, boron, and barium concen-trations in the suspected source of contamination.The constituents that might be useful in identify-ing the contributing sources should be selected byusing knowledge of the composition of contami-nating solution and by considering chemical- andexchange behaviour of the solutes.

6.6.2 Injected chemical tracers

An introduction to hydrologic tracers is given inSection 7.2. Various kinds of ‘chemical’ measuringtechniques suitable for surface streams have beendeveloped (Corbett et al. 1945; Rantz et al. 1982).A slug injection technique has been applied exten-sively to determine the rate of solute movementor time of travel of water through a given reachof a river. This estimate cannot be made accu-rately from stream-discharge data. The procedure

P1: SFK/UKS P2: SFK



involves adding a readily detectable solute to thestream in the form of a concentrated slug and thelength of time required for the material to appearat a downstream point is measured. An index oftime-of-travel studies by US Geological Survey wascompiled by Boning (1973).

The salt-dilution method of measuring the flowrate of a stream consists of adding a known quan-tity of tracer, usually sodium chloride, at a con-stant rate and measuring its concentration in thewater upstream from the point of addition, as wellas far enough downstream so that mixing of thetracer in the stream is complete. The flow must re-main reasonably constant during the interval whilethe measurements are made and enough time mustbe allowed so that the tracer concentration at thedownstream point becomes stable (i.e. time invari-ant). This provides enough data for calculating dis-charge rate of the stream. The method can be usedin streams in which other procedures are not pos-sible because of inaccessibility or extreme condi-tions due to the presence of turbulence or duringa flood. Amounts of salt added should not be ex-cessively large, as it adversely affects water quality(see also Section 2.6.1.4).

In areas where interconnection between surfacewater and groundwater systems is of interest,detailed studies often include seepage estimation.These consist of a series of measurements of riverflow and tributary inflow taken in a downstreamdirection, with gains/losses between measuringpoints being ascribed to groundwater inflowsinto the stream/losses of water from the streamto the groundwater reservoir. However, if watersamples are taken at appropriate measuring sitesand analysed, regions of inflow and outflow maybe better identified to understand the hydrologicsystem. Data of this kind were used, for example,to help evaluate stream-aquifer interconnectionsalong Sabarmati River in Gujarat, India (Bhandariet al. 1986).

Several investigations of direction and ratesof groundwater movement have been made byinjecting slugs of salt, dye, or radioactive material(Kaufman and Orlob 1956). Any tracer materialthat may be added must be similar in density andtemperature to that of the groundwater. Largeamounts of tracer might constitute unacceptable

levels of pollution. In aquifers where groundwatermovement is through large fissures or cavernousopenings, a tracer technique becomes simple andeasy to interpret. Organic dyes have been used totrace water movement through limestone aquifersand to identify pollution sources in such systems.Kaufman and Orlob (1956) observed that chlorideions seemed to move at effectively the same rateas water through porous material.

Materials that are naturally present in ground-water are also potentially useful for tracinggroundwater flow and estimating rates of itsmovement. Use of environmental radioisotopesfor studying groundwater movement is discussedin the Section 7.3.

6.7 Groundwater – numerical modellingof solute transport

As discussed in Chapter 3, the process of ground-water flow is governed by Darcy’s law (Eqn. 3.24)and the law of conservation of mass (Eqn. 3.25).The purpose of a model that simulates solute trans-port in groundwater is to compute the concentra-tion of a dissolved chemical species in an aquifer atany given time and place. Theoretical basis for theequation describing solute transport is well doc-umented in the literature (Bear 1979; Domenicoand Schwartz 1998). A conceptual framework foranalysing and modelling physical solute-transportprocesses in groundwater has been provided byReilly et al. (1987). Changes in chemical concen-tration occur in a dynamic groundwater system,primarily due to four distinct processes:

1) Advective transport, in which dissolved chem-icals move with the flowing groundwater;

2) Hydrodynamic dispersion, in which molecu-lar and ionic diffusion and small-scale variationsin the flow velocity through the porous mediacause the paths of dissolved molecules and ionsto diverge or spread from the average directionof groundwater flow;

3) Fluid sources, where water of one compositionis introduced and mixed with water of a differ-ent composition.

P1: SFK/UKS P2: SFK



4) Reactions, in which some amount of a particu-lar dissolved chemical species may be added toor removed from the groundwater as a result ofchemical, biological, and physical reactions tak-ing place in water or between the water and thesolid aquifer materials or other separate liquidphases (e.g. NAPL).

6.7.1 Governing equations

A general form of the equation (Eqn. 3.36) describ-ing the transient flow of a compressible fluid ina non-homogeneous anisotropic aquifer can be de-rived by combining Darcy’s law with the continuityequation as:

Ss∂h

∂t+ W = ∂

∂x

(Kx

∂h

∂x

)+ ∂

∂y

(Ky

∂h

∂y

)

+ ∂

∂z

(Kz

∂h

∂z

)(6.46)

An equation describing the transport and disper-sion of a dissolved chemical in flowing groundwa-ter may also be derived from the principle of con-servation of mass that requires that the net massof solute entering or leaving a specified volume ofaquifer during a given time interval must equal theaccumulation or loss of mass stored in the volumeduring the interval. This relationship may then beexpressed mathematically by considering all fluxesinto and out of a representative elementary volume(REV). A generalized form of the solute-transportequation was presented by Grove (1976), in whichappropriate terms are incorporated to representchemical reactions and solute concentration bothin the pore fluid and on the solid surfaces, as:

∂ (εC )

∂t= ∂

∂xi

(εDij

∂C

∂xj

)− ∂

∂xi(εC Vi)

− C ′W∗ + CHEM (6.47)

where CHEM equals:

−ρb∂C

∂t

for linear equilibrium controlled

sorption or ion-exchange reactions,

s∑k=1

Rk

for s chemical rate-controlled

reactions, and (or)

−λ(εC + ρbC

)for decay

where ε is the effective porosity; Dij is the coef-ficient of hydrodynamic dispersion (L2T−1); C ′ isthe concentration of the solute in the source orsink fluid; C is the concentration of the species ad-sorbed on the solid (mass of solute/mass of solid);ρb is the bulk density of the sediment [ML−3]; Rk

is the rate of production of the solute in reactionk [ML−3T−1]; and λ is the decay constant (equal toln 2/T 1/2) [T−1] (Grove 1976).

The first term on the right-hand side of Eqn6.47 represents change in the concentration dueto hydrodynamic dispersion. This expression isanalogous to Fick’s Law describing the diffusiveflux. This Fickian model assumes that the drivingforce is the concentration gradient and that thedispersive flux occurs in a direction from higherconcentration towards lower concentration. Thesecond term represents advective transport anddescribes movement of solutes at the averageseepage velocity of groundwater. The third termrepresents effect of mixing with a source fluidthat has a concentration different from that inthe groundwater at the location of recharge orinjection. The fourth term lumps together all of thechemical, geochemical, and biological reactionsthat cause transfer of mass between the liquid andsolid phase or conversion of dissolved chemicalspecies from one form to another. Chemicalattenuation of inorganic chemicals can occur bysorption/desorption, precipitation/dissolution, oroxidation/reduction. Organic chemicals can adsorbor degrade by microbiological processes. Therehas been considerable progress in modelling thesereactions. However, a comprehensive review ofthe reaction processes and their representation intransport models is beyond the scope of this book.

If reactions are limited to equilibrium-controlledsorption or exchange and first-order irreversiblerate (decay) reactions, then the general governingequation (Eqn 6.47) can be written as:

P1: SFK/UKS P2: SFK



∂C

∂t+ ρb

ε

∂C

∂t= ∂

∂xi

(Dij

∂C

∂xj

)− ∂

∂xi(C Vi)

+ C ′W∗

ε− λC − ρb

ελC (6.48)

Any temporal change in sorbed concentration inEqn 6.48 can be represented in terms of the soluteconcentration using the chain rule of calculus, asfollows:

dC

dt= dC

dC

∂C

∂t(6.49)

The quantities dC /dC as well as C are functionsof C alone for equilibrium sorption and exchangereactions. Therefore, the equilibrium relation forC and dC /dC can be substituted into the gov-erning equation to reduce the partial differentialequation in terms of only C. The resulting singletransport equation is solved for solute concentra-tion. Sorbed concentration can then be calculatedusing the equilibrium relation. The linear-sorptionreaction considers that the concentration of solutesorbed onto the porous medium is directly propor-tional to the concentration of the solute in the porefluid, according to the relation:

C = KdC (6.50)

where Kd is the distribution coefficient [L3M−1].This reaction is assumed to be instantaneous andreversible. The curve relating sorbed concentra-tion to dissolved concentration is known as anisotherm. If this relation is linear, the slope (givenby the derivative) of the isotherm, dC /dC , isknown as the equilibrium distribution coefficient,Kd. Thus, in the case of a linear isotherm:

dC

dt= dC

dC

∂C

∂t= Kd

∂C

∂t(6.51)

After substituting and rewriting Eqn 6.48:

∂C

∂t+ ρbKd

ε

∂C

∂t= ∂

∂xi

(Dij

∂C

∂xj

)− ∂

∂xi(C Vi)

+ C ′W∗

ε− λC − ρbKd

ελC (6.52)

Defining a dimensionless retardation factor, Rf, as:

R f = 1 + ρbKd

ε(6.53)

and substituting this relation into Eqn 6.52:

R f∂C

∂t= ∂

∂xi

(Dij

∂C

∂xj

)− ∂

∂xi(C Vi)

+ C ′W∗

ε− R f λC (6.54)

As Rf is constant under these assumptions, so-lution to this governing equation is identical tothe solution of the governing equation with nosorption, except that the velocity, dispersive flux,and source strength are reduced by a factor, Rf.The transport process thus appears to be ‘retarded’because of instantaneous equilibrium sorptiononto the particle surfaces in the porous medium.

In the conventional formulation of the solute-transport equation (Eqn 6.47), the coefficient of hy-drodynamic dispersion is defined as the sum of me-chanical dispersion and molecular diffusion. Themechanical dispersion is a function of both theintrinsic properties of the porous medium (suchas heterogeneities in hydraulic conductivity andporosity) as well as of the fluid flow. Molecular dif-fusion in a porous medium differs from that in freewater because of the effects of porosity and tortu-osity. These relations are commonly expressed as:

Dij = αi jmnVmVn

|V| + Dm i, j, m, n = 1, 2, 3

(6.55)

where αijmn is the dispersivity of the porousmedium (a fourth-order tensor) [L]; Vm and Vn

are components of the flow velocity of the fluidin the m and n directions, respectively [LT−1]; Dm

is the effective coefficient of molecular diffusion[L2T−1]; and |V| is the magnitude of the velocityvector [LT−1], defined as (Bear 1979; Domenicoand Schwartz 1998; Scheidegger 1961):

V =√

V 2x + V 2

y + V 2z

The dispersivity of an isotropic porous medium canbe defined by two constants – the longitudinal dis-persivity of the medium, αL, and the transverse dis-persivity of the medium, αT. These are related to thelongitudinal and transverse dispersion coefficientsby DL = αL|V| and DT = αT|V|. Most applicationsof transport models to groundwater problems arebased on this conventional formulation.

P1: SFK/UKS P2: SFK



RELIABILITY

LowIntermediateHigh

104

103

102

101

100

10-2

10-3

10-1

10510-1 104103102101100 106

Lo

ng

itu

din

al D

isp

ersi

vity

(m

)

Scale of Transport (m)

Fig. 6.11 Model-fitted longitudinal dispersivity in saturatedmedia versus scale of modelled plume. Macro-dispersivitieswere determined by calibrating solute transport modelswith the observed solute plumes. Reliability classificationfollows original diagram of Gelhar et al. (1992). ©American Geophysical Union.

Although conventional theory holds that αL isgenerally an intrinsic property of the aquifer, it isfound in practice to be proportional to the scale ofthe measurement (Fig. 6.11). But this trend is muchless evident when reliability of data (Fig. 6.11) isconsidered (Gelhar et al. 1992). Most reported val-ues of αL fall in the range from 0.01 to 1.0 times thescale of the measurement, although the ratio of αL

to scale of measurement tends to decrease at largerspatial scales (Anderson 1984; Gelhar et al. 1992).Field-scale dispersion (commonly called macro-dispersion) results from large-scale spatial varia-tions in hydraulic properties. Consequently, use ofrelatively large values of dispersivity together withuniform hydraulic properties (Kij and ε) is inappro-priate for describing transport in geological systems(Smith and Schwartz 1980). If a model applied to asystem having variable hydraulic conductivity usesmean values and thereby does not explicitly repre-sent the variability, it is likely that the model cali-bration will yield values for the dispersivity coeffi-cients that are larger than what would be measuredlocally in the field. Similarly, representing the tran-sient flow field by a mean steady-state flow field,as is commonly done, inherently ignores some ofthe variability in velocity and must be compensated

for by using higher values of dispersivity (primar-ily transverse dispersivity) (Goode and Konikow1990). Overall, the higher the accuracy with whicha model can represent or simulate the true veloc-ity distribution in space and time, the uncertaintyconcerning representation of dispersion processeswill be correspondingly a smaller problem.

A special form of the solute-transport equationcan be used for direct simulation of groundwaterages (Goode, 1996, 1999). This is accomplishedby adding a zero-order growth term, which repre-sents production of the solute [ML−3T−1] withinthe system itself. In developing an age transportequation, concentrations are replaced with cor-responding ages representing a volume-averagedgroundwater age in the aquifer and the zero-ordergrowth rate having a value equal to unity. Decayand sorption reactions are assumed to be absentand, in general, the age of incoming water (analo-gous to C′) is specified as zero. This type of analysisallows a direct comparison of groundwater mod-elling results with measured environmental tracerdata, while accounting for effects of dispersion andother transport processes (see Section 7.6).

6.8 Relation between use and qualityof water

An obvious purpose of water quality investigationis to determine if a given water supply is satisfac-tory for the intended use(s). Some discussion onwater quality, in terms of physical and chemicalparameters and pathogens, is given in Section1.2. Standards for water meant to be used fordrinking and other domestic purposes have beenestablished in many countries. Published literaturecontains tolerance levels and related data forconstituents of water to be used for agriculture,in industry, for development of fisheries, and fora number of other specific purposes. Water thatis meant to be used for domestic supply may beemployed for many purposes. Therefore, the stan-dards used to evaluate the suitability of water forpublic supplies are generally more stringent thanthose applied to water for a small domestic or farmsupply. Water from zones of mineralization and hotsprings is used medicinally in many places and the

P1: SFK/UKS P2: SFK



Table 6.15 Web links with details of drinking water quality standards and guidelines.

WHO http://www.who.int/water sanitation health/dwq/gdwq3rev/en/index.htmlUSA http://www.epa.gov/safewater/standards.htmlIndia http://www.chennaimetrowater.tn.nic.in/qualitymainpage.htmWHO/EU http://www.lenntech.com/WHO-EU-water-standards.htm

mystic qualities of natural hot springs have been ofgreat interest to man since prehistoric times.

6.8.1 Domestic uses and public supplies

Besides being chemically safe for human con-sumption, water for domestic use should be freeof undesirable physical properties such as colouror turbidity and should not have an unpleasanttaste or odour. Harmful micro-organisms should bevirtually absent, even though these are not usuallyconsidered in routine chemical analyses. Presenceof harmful micro-organisms is considerably moredifficult to ascertain than other properties of water,but they are of utmost concern. Additional risksarise from toxic chemicals and radiological hazards.

Mandatory standards for dissolved constituentsbelieved to be harmful to humans were first estab-lished in the United States in 1914 by the US PublicHealth Service. In 1974, the federal Safe DrinkingWater Act was legislated and standards for concen-tration of inorganic constituents in public watersupplies became effective in 1977. Presently theUnited States has one of the safest water suppliesin the world, but drinking water quality is still anissue of concern for human health in developingas well as many of the developed countries world-wide. Drinking water quality varies from place toplace, depending on condition of the source fromwhich it is drawn and the treatment it receives priorto supply. WHO has stipulated international normson water quality and human health in the form ofguidelines that are used worldwide as a basis forsetting up regulatory guidelines in developing aswell as in developed countries. For example, WHOguidelines for drinking water are used as a basis forthe standards in the Drinking Water Directive in Eu-ropean Union, but with some differences. In India,the Bureau of Indian Standards (BIS) has notifiedstandard drinking water specifications through BIS

10500: 1990. The web links in Table 6.15 can beaccessed to obtain details of some of the drinkingwater quality standards and guidelines.

The limiting concentrations of radioactive sub-stances in drinking water are viewed somewhatdifferently from those of non-radioactive solutes. Itis generally agreed that the effects of radioactivityare harmful, and unnecessary exposure should beavoided. Strontium-90 is a fission product, but ra-dium occurs naturally. Both nuclides are preferen-tially absorbed in bone structure and are, therefore,especially undesirable in drinking water.

6.8.2 Agricultural use

Water required for non-domestic purposes onfarms includes that consumed by livestock and forirrigation. Water for livestock is subject to simi-lar quality considerations as those related to drink-ing water for human consumption. Most animals,however, can tolerate water that has a considerablyhigher concentration of dissolved solids than thatwhich is considered safe for humans.

The chemical quality of water is an importantfactor to be considered in evaluating its usefulnessfor irrigation. Features of the chemical composi-tion that need to be considered include concen-tration of total dissolved matter, concentrations ofcertain potentially toxic constituents, and relativeproportions of some specific constituents. Suitabil-ity of particular water for irrigation also depends onmany factors not directly associated with composi-tion of water. A brief discussion of some of thesefactors is given here to highlight the complexityof the problem for deciding whether or not givenwater is suitable for irrigation.

Part of the irrigation water that is actually con-sumed by plants or evaporated is virtually free ofdissolved material. Growing plants selectively re-tain some nutrients and a part of the mineral matter

P1: SFK/UKS P2: SFK



originally dissolved in the water, but the amount ofmajor cations and anions thus retained is only asmall part of their total content in the irrigation wa-ter. Eaton (1954) showed that this consists mostlyof calcium and magnesium salts. The bulk of thesoluble material originally present in irrigation wa-ter stays behind in solution in residual water. Con-centration of solutes in soil moisture cannot be al-lowed to rise too high, in order to avoid interfer-ence with the osmotic process by which plant rootmembranes assimilate water together with nutri-ents. Some compounds of low solubility, especiallycalcium carbonate, which is virtually harmless, mayprecipitate in the soil as solute concentrations in-crease, but the bulk of the residual solutes mustbe managed effectively to maintain productivity ofirrigated soils.

The extent and severity of salt accumulationproblems in irrigated areas depend on severalfactors. These include: (i) chemical compositionof the water supply; (ii) nature and composition ofthe topsoil and subsoil; (iii) topography of the land;(iv) amount of water used; (v) method of irrigationemployed (flooding the fields/sprinkler/drip); (vi)types of crops grown; (vii) climate of the region,especially the amount and distribution of rainfall;and (viii) groundwater conditions (depth to watertable, quality) and nature of surface-water drainagesystem.

In most areas, excess of the soluble material leftin the soil from previous irrigation is removed byleaching of the topsoil and percolation below theroot zone of a part of the resulting solution into thegroundwater reservoir. In areas where the watertable beneath the irrigated land can be kept suffi-ciently below the surface, this process of drainageis reasonably effective. The leaching may be ac-complished by rainfall in areas where precipita-tion is sufficient to saturate a large depth of soil.Leaching of soluble salts also occurs during irriga-tion when an excess amount of water is added,with the aim to store the extra supply of waterin the soil or to use up the surplus amount ofwater that happens to be available at a particu-lar time. The need for leaching of the soil with aview to remove excess salts is generally recognizedby farmers who use highly mineralized irrigationwater.

For long-term successful operation of an irriga-tion project, all the ions present in the irrigationwater and those extracted by plants must be dis-posed off either by carrying them away from thearea or by storing them safely within the area. Thenet ion load in an irrigated area can be expressedin terms of the salt balance, i.e. the difference be-tween ion inflow and outflow. Besides the generalincrease in major solute ion concentrations thatirrigation drainage may cause in the groundwaterunderlying the irrigated land, there may be addi-tions of specific solutes that are undesirable. A ma-jor problem in some irrigated regions has been theincreasing concentration of nitrate in groundwa-ter received from the drainage of irrigated fields onwhich nitrogenous chemical fertilizers (particularlyurea) has been applied. Some types of pesticidesmay also persist in the drainage water.

In addition to problems of excessive concentra-tion of dissolved solids, certain constituents in irri-gation water are especially undesirable, even whenpresent in trace concentrations. Boron, for exam-ple, is an essential plant nutrient and is sometimesadded to fertilizers in small amounts because somesoils in humid regions are deficient in boron. Buteven a small excess of boron over the plant tol-erance level is toxic to some types of plants, par-ticularly citrus fruits and walnut trees. Lithium inwater in small concentrations (0.06–0.10 mg l−1)has been shown to cause damage to citrus plants(Bradford 1963). Soils of high salinity interfere withcrop growth and a high pH may decrease the solu-bility of some essential elements.

Some minor constituents of irrigation water,notably molybdenum, selenium, and cadmium,may accumulate in plant tissues and cause toxicitywhen the plants or their seeds are consumed byhumans/animals.

The process of cation exchange also occurs inirrigated soils and may influence soil properties, es-pecially when concentrations of solutes are high.Irrigation water with a high ratio of sodium tototal cations tends to put sodium ions in the ex-change positions on the soil-mineral particles. Inwater having mostly divalent cations, this processis reversed. In soils, clay minerals have the highestexchange capacity per unit weight. Physical prop-erties of soils are optimal for plant cultivation and

P1: SFK/UKS P2: SFK



growth when their exchange sites are occupied bydivalent ions of calcium and magnesium. However,when exchange positions become saturated withsodium, soils tends to become deflocculated and,therefore, impermeable to water. A soil of this typeis difficult to cultivate and may not support plantgrowth.

The cation-exchange process is reversible andcan be controlled either by adjusting the compo-sition of the water or by using soil amendments.The condition of a sodium-saturated soil can be im-proved by liberal application of gypsum, which re-leases calcium to occupy exchange positions. Thesoil also may be treated with sulphur, sulphuricacid, ferrous sulphate, or other chemicals that tendto lower the pH of the soil solution. The lowerpH brings calcium into solution by dissolving car-bonates or other calcium minerals. The tendencyof water to replace adsorbed calcium and magne-sium with sodium can be expressed by the sodium-adsorption ratio (SAR):

SAR = (Na+)√1/2

[(Ca2+) + (Mg2+)

] (6.56)

where ion concentrations (in parentheses) are ex-pressed in meq l−1.

Two other related indices are: (i) Soluble SodiumProportion (SSP) indicating proportion of sodiumions in solution in relation to the total cation con-centration in water, defined as:

SSP = Soluble Sodium Concentration (meq.l−1)

Total Cation Concentration (meq.l−1)

×100 (6.57)

and (ii) Exchangeable Sodium Percentage (ESP)in soil defined as:

ESP =Exchangeable Sodium

(meq.(100g)−1 soil)

Cation Exchange Capacity(meq.(100g)−1 soil)

× 100 (6.58)

Whereas SSP is an indicator of the sodium hazardfrom the irrigation water, ESP, on the other hand,indicates the extent to which the adsorption com-plex of a soil is occupied by sodium. It has beenobserved that where irrigation water and drainageconditions are good, the ESP value of the soil varies

only slightly from season to season or year to year.This implies that the cation exchange material ofthe soil has reached a steady state relative to thecations in the soil solution, which are derived fromthe irrigation water. Under such conditions, an em-pirical relation (Eqn 6.59) has been observed be-tween ESP and SAR:

ESP = 100 (−0.0126 + 0.01475 × S AR)

1 + (−0.0126 + 0.0145 × S AR)(6.59)

On the basis of this relationship (Eqn 6.59), SARappears to be a useful index for designating sodiumhazard of waters used for irrigation.

Eaton (1950) suggested that if much of the cal-cium and magnesium originally present were pre-cipitated, the residual water would be considerablyenriched in sodium relative to the other cations.Some waters, in which the bicarbonate contentis higher than the amount equivalent to the to-tal amount of calcium and magnesium, could thusevolve into solutions containing mostly sodium andbicarbonate and would have a high pH and poten-tial for deposition of sodium carbonate (commonlyknown as black alkali). Residual sodium carbonate(RSC) is defined as an excess of carbonate or bi-carbonate that water contains after subtracting anamount equivalent to the calcium plus the magne-sium, i.e.:

RSC = (CO2+

3 + HCO−3

)

− (Ca2+ + Mg2+)

in meq.l−1 (6.60)

RSC is another alternative measure of the sodiumcontent in relation to Mg and Ca. This value mayappear in some water quality reports, although itis not frequently used. Water is considered safe forirrigation if RSC <1.25 and not appropriate if theRSC >2.5.

From this brief discussion it should be evidentthat the relationship between water quality andsuitability of water for irrigation is not simple.Further complications can arise as salinity of waterincreases. A diagram widely used for evaluatingsuitability of waters for irrigation, published bythe US Salinity Laboratory (US-SL 1954), is givenin Fig. 6.12. In this diagram specific conductanceor electrical conductivity (EC), as an index of dis-solved solids concentration, is plotted on one axisand the sodium-adsorption ratio on the other. The

P1: SFK/UKS P2: SFK



Class

Low Medium High V. HighC1 C2 C3 C4

100 250 750 2250

Salinity Hazard

So

diu

m (

Alk

ali)

Haz

ard

Lo

wM

ediu

mH

igh

V. High

S1

S2

S3

S4

So

diu

m-A

dso

rpti

on

-Rat

io (

SA

R)

Specific Conductance (EC)in micro-Siemens per cm

at 25° C1000 5000100

0

2

4

6

8

10

12

14

16

18

20

22

24

26

28

30

10

20

30

0

Fig. 6.12 Diagram used in interpreting the analysis ofirrigation water. Redrawn from US Salinity Laboratory(US-SL 1954). © U.S. Geological Survey.

diagram is divided into 16 areas that are used tocategorize/specify the degree to which a particularwater source may contribute to salinity problemsand undesirable ion-exchange effects in soil. Waterhaving EC >5000 µS cm−1 is also being used withsome success in certain areas where proper soilconditions exist for growing suitable crops, andappropriate irrigation techniques are employed. Ahydrologist needs to consider the local experiencerather than arbitrarily deciding whether givenwater is suitable for irrigation at a given place.Salinity problems, however, may be slow to de-velop and may be observable only through indirectmeans, such as reduced crop yields or otherindicators that are not easy to evaluate. A water ofhigh salinity must always be viewed with cautionuntil proof of it being safe for a specific use isestablished.

6.9 Industrial use

Quality requirements for industrial water suppliesvary widely as almost every industrial applicationhas its own norms. For some uses, such as single-pass condensing of steam or for cooling or con-centrating ores, chemical quality is not particularlycrucial and almost any water may be used. At theother extreme, water of nearly distilled water qual-ity is required for processes such as manufacturingof high-grade paper or pharmaceuticals, as impuri-ties in water can seriously impact the product qual-ity. In nuclear reactors, water of very high purity isdesirable to minimize the radioactivity induced byneutron activation of the dissolved constituents.

Technically it is possible to treat any water tomake it suitable for a specific use. However, if ex-tensive treatment is required and large volumes ofwater are involved, it may not be economically vi-able to use some of the supply sources. Industrialplants requiring large quantities of water, therefore,need to be suitably located by considering availabil-ity of water of the desired quality.

Although not a chemical property, the tempera-ture of a water supply source and its seasonal fluc-tuations are major considerations in its use in indus-try for cooling purposes. In some areas, groundwa-ter is used extensively for this purpose because itstemperature is uniform and is below ambient airtemperatures during warm weather and above am-bient air temperatures during cold weather. Someindustries recharge groundwater aquifers with coldwater from surface streams during winter and with-draw it in the summer when the surface wateris too warm for efficient cooling. In recent yearsthe practice of discharging excessively warm wa-ter into streams is regulated to prevent ecologicalstress due to depletion of dissolved oxygen in waterat elevated temperatures.

Much of industrial use, unlike agricultural use,is non-consumptive in the sense that the water isnot evaporated or incorporated into the finishedproduct but is discarded after a single use withoutsignificant change in its quantity, but generally withan increased load of dissolved or suspended mate-rial. As water supplies generally become more fullycommitted in the course of time, many industriesfind it necessary to conserve and reuse water which

P1: SFK/UKS P2: SFK



in the past would have been allowed to flow downa sewer or released into a surface stream. In severalinstances, reclaimed sewage is being used for cer-tain non-critical industrial applications in terms ofwater quality.

6.9.1 Recreational and aesthetic uses

Considerable attention is now being paid to uses ofrivers and lakes for such purposes as religious cere-monies, swimming, fishing, boating, and aestheticand recreational purposes. Many of the surface wa-ter bodies in India are intimately linked to religiouspractices and periodic congregation of large num-bers of devotees on the banks of streams and lakesand holy dips in rivers are a common practice.Restoring such water bodies entails enormous costsbecause of the high level of pollution but there isstrong public support with the aim of creating orprotecting them for religious purposes.

Water for swimming and other sports in whichwater is in contact with human skin must obviouslyconform to sanitary standards. Fish require cleanwater with a good supply of dissolved oxygen.Certain metal ions may be lethal to fish and otheraquatic life forms when present at levels close tothe limits specified for public water supplies. Cop-per, zinc, and aluminium, which are not amongstthe metals for which limits are prescribed forpublic water supplies, are toxic to fish and manyother aquatic life forms. Assimilation of dissolvedmetal ions by aquatic biota has a tendency ofincreasing concentrations in species higher up inthe food chain. One of the more insidious effectsof mercury-containing wastes that enter rivers andlakes is an increase in mercury content of fish, tothe extent that they become dangerous for humanconsumption.

6.10 Tutorial

Ex 6.1 How many moles of hydrogen ions are con-tained in 1 litre solution of pH 10? How many gramsand how many numbers of hydrogen ions are con-tained in the same solution?

[Ans. 10−10, 10−10, 6.022 × 10−13]

Ex 6.2 What is the pH of pure water in equilibriumwith the atmosphere (pCO2 =10−3.5 atm) at 25◦C?

Solution Using Eqn 6.30, [H2CO3∗] = 10−1.47 ×

10−3.5 = 10−4.97. Now using Eqn 6.31, [H+][HCO3

−] = 10−6.35 × 10−4.97 = 10−11.32. Becausethe solution must have neutral charge, i.e. concen-trations of cations and anions must balance so that[H+] = [HCO3

−]; [H+]2 = 10−11.32. Or [H+] =10−5.66; so that pH = 5.66. One thus expects thepH of rain to be around 5.66. In reality the pH ofrain is quite variable, influenced by other solutesderived from the atmosphere.

Ex 6.3 Calculate the pH of a 0.0250 M solution ofCO2 in water.

[Hint. Use Eqn 1.31. Ans. pH = 3:97]

Ex 6.4 A 100-ml sample of natural water whosepH is 6.6 requires 12.2 ml of 0.10 M HCl for titra-tion to the methyl orange end point and 5.85 ml of0.10 M NaOH for titration to the phenolphthaleinend point. Assuming that only carbonate speciesare present in significant quantities, find the totalalkalinity, carbon dioxide acidity, and the total acid-ity of the water.

Solution Since the pH is below 8.3, bicarbonate isthe major alkalinity species.

Total alkalinity: 12.2 mM l−1.

CO2 acidity (conversion of CO2 to HCO3−):

5.85 mM l−1

Total acidity (conversion of all carbonate speciesto CO3

2−): Because it is impractical to carry outthis titration, one can make use of the data alreadyavailable. Conversion of the initial CO2 to HCO3

−

would require 5.85 mM l−1 of NaOH, and thenan additional equal amount of HCO3

−. Similarly,the 12.2 mM l−1 of HCO3

− initially present willrequire the same quantity of NaOH for conversionto CO3

2−. The total acidity is thus (5.85 + 5.85 +12.2) = 23.9 mM l−1.

Ex 6.5 Calculate the amount of helium dissolved inair-saturated water under normal atmospheric con-ditions at 25◦C, given Henry’s constant for oxygenat 25 ◦C, i.e. Hi = 2.865 ×103 atm/(mol l−1) can becalculated as follows. Under normal atmosphericconditions there is 5.24 × 10−4 mole per cent he-lium, which makes the partial pressure of helium5.24 ×10−6 atm. Using Henry’s law, Eqn 6.25, the

P1: SFK/UKS P2: SFK



Table 6.16

Ca+2 +Mg+2 Na+ K+ CO3

−2 HCO3−1 Cl− F− SO4

−2 SAR RSC SSP

S. No. pHEC µmhos.cm−1

TDSmg.l−1 meq.l−1 ppm meq.l−1 %

1 8.62 1023 560 5.8 5.08 0.03 1.3 6.8 2.8 0.30 0.02 9.20 3442 1910 3.1 32.87 0.01 1.1 14.2 17.5 0.34 2.63 9.40 1945 1040 4.8 15.56 0.01 1.4 6.3 10.5 0.44 0.84 9.00 2739 1456 4.2 26.30 0.00 1.6 10.7 14.0 0.24 1.05 8.41 2006 1078 7.9 12.00 0.02 1.2 6.6 11.0 0.24 3.06 8.60 3304 1715 6.7 24.87 0.02 1.4 11.9 19.5 0.26 0.07 8.80 231 145 3.6 0.50 0.00 0.6 2.9 3.6 0.30 0.08 8.65 1189 660 3.8 9.78 0.01 1.0 9.6 2.0 1.10 0.09 8.68 614 350 4.1 4.50 0.03 1.6 4.9 1.0 0.22 0.0

10 8.80 1620 940 12.3 7.50 0.03 0.9 4.8 8.5 0.20 0.511 8.00 3281 1880 4.6 29.78 0.02 1.2 15.0 16.0 0.52 1.312 8.45 1831 1070 3.5 15.30 0.02 1.3 6.3 11.0 0.39 0.813 8.75 640 410 3.2 3.39 0.01 1.1 4.1 2.2 0.56 0.0

concentration of helium is [5.24 × 10−6 atm /Hi],which is 1.83 × 10−9 mol l−1 or 1.83 × 10−6 mmoll−1.

Ex 6.6 Calculate the amount of carbon dioxide dis-solved in 1 litre of soda pop if the manufactureruses a pressure of 2.4 atm of CO2 to carbonate thesoda pop. Given Hi for CO2 = 2.976 ×101 atm/(mol l−1).

Ex 6.7 Estimate the amount of nitrogen that a divermust lose from his bloodstream (∼5 l) in risingfrom a depth of 100 m to the surface, in order toavoid formation of nitrogen bubbles in his blood-stream. Given Hi for nitrogen = 1.55 × 108 Pa/(moll−1).

Solution The density of water is about 1 kg l−1

or 1000 kg m−3. A column of water 100 m thickwould have a mass of 100,000 kg m−2 at its base,which would exert an additional force of 980,665Nm2 or 980,665 Pa. (The total pressure would be980,665 kPa plus 101.325 kPa or 1081.990 kPa.)

The pressure change of 980,665 Pa would pro-ducea concentration change of 980,665/(1.55 ×108) = 6.33 mmol l−1. Therefore, the amount ofnitrogen that must be lost is 5 × 6.33 = 31.65mmol or >750 ml of nitrogen at room tempera-ture and pressure. This is enough nitrogen gas tocreate massive bubbles in the bloodstream.

Ex 6.8 Is a diver is safer using a helium/oxygenmixture than a nitrogen/oxygen mixture when div-ing to longer depths or remaining submerged forlonger periods of time? Given Hi for helium =2.83 × 108 Pa/(mol l−1).

Ex 6.9 Table 6.16 gives results of some chemicalanalyses of groundwater from Bhiloda Taluka (N.Gujarat), India. Source: Acharya et al. (2008).

Plot the data of these samples as bar charts, piecharts, Stiff diagram, Schoeller diagram, and trilin-ear diagram. Complete the table by computing val-ues of SAR, RSC, and SSP. What conclusions canbe drawn about irrigation water quality and otherregional aspects from these analyses?

modern hydrology and sustainable water development (gupta/modern hydrology and sustainable water...

Documents