and alkalinity is, in this example, approximately linear, although for every mmol

of Ca^2 +, 2 mmol of HCO 3 - are released, as predicted in equation 5.12.

5.4 Aluminium solubility and acidity

Aluminium is largely insoluble during weathering processes (Table 5.1), but

becomes soluble when pH is both low and high. At the simplest level, three

aluminium species are identified; soluble Al^3 +, dominant under acid conditions,

insoluble aluminium hydroxide (Al(OH) 3 ), dominant under neutral conditions,

and Al(OH) 4 - , dominant under alkaline conditions.

eqn. 5.13

eqn. 5.14

Aluminium solubility is therefore pH-dependent and aluminium is insoluble in

the pH range 5–9, which includes most natural waters. The details of aluminium

solubility are also complicated by the formation of partially dissociated Al(OH) 3

species and complexing between aluminium and organic matter (see Box 6.4).

Although aluminium is soluble at high pH, alkaline waters are uncommon

because they absorb acid gases, for example CO 2 and SO 2 , from the atmosphere.

However, alkaline rivers with aluminium mobility are known. For example, the

industrial process for abstracting aluminium from bauxite involves leaching the

ore with strong sodium hydroxide (NaOH) solutions. In Jamaica, discharge of

wastes from bauxite processing produces freshwater streams with high pH (>8)

in addition to high sodium and aluminium concentrations. As these streamwaters

evolve, the pH falls to about 8 and the dissolved aluminium : sodium ratio declines

as the aluminium precipitates.

It is, however, acidificationof freshwaters that commonly results in aluminium

mobility resulting in ecological damage. This acidification is typically caused by

two anthropogenic processes, acid rain and acid mine drainage.

Al OH() 3 ()s ªAl^3 ()aq+ + 3 OH()-aq

Al OH() 34 ()s+OH()aq ªAl OH()-

The Chemistry of Continental Waters 155

eqn. 16

(see also Box 3.7), and so:

eqn. 17

Now equation 13 can be rewritten:

eqn. 18

Consider the case of the Mackenzie River,
where HCO 3 - =1.8 mmol l-^1 (Table 5.2) and
atmospheric pCO 2 =3.6¥ 10 -^4 atm:

aH+- aHCOCOmol l

• 
  
  
  
  • =¥ 1064 1014 ¥^2 -
    3
    .. p 1
  
  
  
  

aHCO 23 =¥ 10 -^14. pCO 2

KCO COHCOH O pHCOCO

mol l atm

(^22223232)
14 1 1
004
10

= ---
a
aa
a

..
. eqn. 19

eqn. 20

Although this treatment is simplified,

it serves to illustrate the way in which

pH can be calculated. In practice the pH

of most natural water containing HCO 3 -

and CO 32 - is buffered between pH 7

and 9.

pH=-log 10 aH+-=-log 10 32 10.¥^9 = 849.

aH

mol l

+- ---

---

=¥◊¥¥

=¥ =¥

(^1010) 18 103 6 10
10 0 008 3 2 10
64 14 3 4
64 9 1

..
.

.

.

..