Page 109 - Hydrogeology Principles and Practice

HYDC03 12/5/05 5:36 PM Page 92

92 Chapter Three

BO X
Acid–base reactions
3.4
+
Acids and bases are important chemical compounds that exert par- H O j H + OH − eq. 5
2
ticular control over reactions in water. Acids are commonly consid-
+
+
ered as compounds that dissociate to yield hydrogen ions (protons) In reality, H cannot exist, and H O (hydronium) is formed by the
3
+
+
in water: interaction of water and H . However, it is convenient to use H in
chemical equations. From the law of mass action, the equilibrium
HCl (aq) j H + (aq) + Cl − (aq) eq. 1 constant for this dissociation is:
+
HOH ]
Bases (or alkalis) can be considered as those substances which yield K = [][ − = 10 − 14 eq. 6
−
hydroxide (OH ) ions in aqueous solutions: HO [ HO]
2
2
−
NaOH (aq) j Na + (aq) + OH (aq) eq. 2 For water that is neutral, there are exactly the same concentrations
+
−
−7
−7
(10 ) of H and OH ions such that pH =−log [10 ] = 7. If pH
10
+
−
< 7, there are more H ions than OH ions and the solution is acidic.
Acids and bases react to neutralize each other, producing a dis- − +
solved salt plus water: If pH > 7, there are more OH ions than H ions and the solution is
basic. It is important to notice that pH is a logarithmic scale and so
HCl (aq) + NaOH (aq) j Cl − (aq) + Na + (aq) + H O eq. 3 it is not appropriate to average pH values of solutions. Instead, it is
+
2 (l)
better to average H concentrations.
Acid–base pairs commonly present in groundwater are those
Hydrochloric acid (HCl) and sodium hydroxide (NaOH) are recog- associated with carbonic acid and water itself. Boric, orthophospho-
nized, respectively, as strong acids and bases that dissociate com- ric and humic acids are minor constituents of groundwater but are
pletely in solution to form ions. Weak acids and bases dissociate relatively unimportant in controlling acid–base chemistry. Many
only partly. aquifers of sedimentary origin contain significant amounts of solid
The acidity of aqueous solutions is often described in terms of the carbonate such as calcite (CaCO ), a fairly strong base that con-
3
pH scale. The pH of a solution is defined as: tributes a carbonate ion, thus rendering the solution more alkaline,
and dolomite (CaMg(CO ) ) which participate in equilibrium reac-
3 2
+
pH =−log [H ] eq. 4 tions involving carbonic acid. All acid–base reactions encountered
10
in natural aqueous chemistry are fast such that acid–base systems
Water undergoes dissociation into two ionic species as follows: are always in equilibrium in solution.
[ HCO ] Using a mass balance expression for the dissolved
K = 2 3 eq. 3.12
CO 2 Pco inorganic carbon (DIC) in the acid and its dissociated
2
anionic species, expressed in terms of molality, then:
Carbonic acid is polyprotic (i.e. it has more than one − 2−
+
H ion) and dissociates in two steps: DIC = (H CO ) + (HCO ) + (CO ) eq. 3.17
2
3
3
3
+
H CO j H + HCO − eq. 3.13 Rearranging equations 3.13–3.17, and taking an arbit-
2 3 3
rary value of unity (1) for DIC, equations for the relat-
−
+
−
HCO j H + CO 2− eq. 3.14 ive concentrations of H CO , HCO and CO 3 2− as
2
3
3
3 3
a function of pH are obtained as shown graphically
in Fig. 3.13. It can be seen from Fig. 3.13 that over
From the law of mass action, dissociation constants
most of the normal pH range of groundwater (6–9),
can be expressed as follows: −
HCO is the dominant carbonate species and this
3
−
+ − explains why HCO is one of the major dissolved
[ H ][ HCO ] 3
K = 3 eq. 3.15 inorganic species in groundwater.
HCO 3 [ HCO ]
2
2 3 To calculate actual concentrations of inorganic
carbon species in groundwater, first consider the
+ 2−
[ H ][ CO ] dissolution of calcite by carbonic acid (eq. 3.5). With
K − = 3 eq. 3.16
HCO 3 HCO − reference to equation 3.10, if the partial pressure of
3