Mica Biotite KMg$3/2$Fe$3/2[$AlSi$3]$O$10($OH$)2($s$)+7$H$++1/2$H$2$O $->$ K$++3/2$Mg$2++3/2$Fe$2++$

$2$H$4$SiO$4($aq$)+1/2$A$12$Si$2$O$5($OH$)4($s$)(2\mathrm{.}16)$

Muscovite KAl$2[$AlSi$3]$O$10($OH$)2($s$)+$ H$++3/2$H$2$O $->$ K$++3/2$A$12$Si$2$O$5($OH$)4($s$)(2\mathrm{.}17)$

Figure $2\mathrm{.}2$: The pH dependence of solubility for aluminium$-$ and iron$-$hydroxide minerals $($from Stumm and Morgan, $1996).$ Note that below neutral pH$,$ solubility of these phases increases with decreasing pH$.$ This emphasises the role of pH as a master variable to control the solubility of metal ions and thus their mobility. Metal ions such as Fe$3+$ react with hydroxide ions to form hydrolysis complexes such as FeOH$2+.$The straight lines that are plotted show how concentrations of individual metal species such as Fe$3+,$ FeOH$2+,$ Fe$($OH$)4-$ etc. depend on pH$.$ The dashed line in each plot is the sum of concentrations for all species and thus corresponds to the total dissolved concentration of Fe or AI for solutions in solubility equilibrium with minerals. The metal ions released by sulphide weathering can also precipitate as sulphate$-,$carbonate$-$ and in some: cases silicate$-$ mineral phases: Some examples includeminerals that immobilize: Ferric lron Jarosite: KFe$3($SO$4)2($OH$)6($s$),$Ferrous Iron Melanterite: FeSO$4\mathrm{.}7$H$2$O $($s$),$ Aluminum Alunite: KAl$3($SO$4)2($OH$)6($s$),$Lead Anglesite: PbSO$4($s$)$; Cerrusite PbCO$3($s$)$ and Copper Malachite: Cu$2($OH$)2($CO$3)($s$).$

ACIDITY, ALKALINITY AND pH

Because acidity is a key contaminant in mine water, and because metal ion mobility is

critically dependent on pH$,$ it is necessary to clearly understand how pH depends on

acidity and alkalinity. The relationship is not trivial. A rigorous definition requires review

of some basic concepts from aqueous chemistry.

oxidation of the ferrous iron and precipitation of ochre resulting in no net loss or

gain of acidity.

FeCO$3($s$)+$ H$+->$ Fe$2++$ HCO$3-(2\mathrm{.}18)$

Fe$2++1/4$O$2+5/2$H$2$O $->$ Fe$($OH$)3($s$)+$

Before the acidity and alkalinity of a mine water can be related to the pH$,$ some

additional concepts are needed. In particular it is necessary to understand what

differentiates "strong" acids and bases from "weak" ones.

Strong acids and bases are compounds that dissociate completely when dissolved in

water, releasing protons or hydroxide ions respectively, in the process. For natural

waters, strong acids include H$2$SO$4,$ HNO$3$ and HCI, while strong bases include NaOH,

KOH, Mg$($OH$)2$ and Ca$($OH$)2$ Because these compounds dissociate completely, the

amount added can be determined by the amount of "acid anion" $($SO$42-,$ NO$3$

$-,$ CI$-)$ or "base cation" $($Na$+,$ K$+,$ Mg$2+,$ Ca$2+)$ in solution. Equations $(2\mathrm{.}20)$ and $(2\mathrm{.}21)$

mathematically define acidity $($Acy$)$ and alkalinity $($Alk$)$ according to the excess of strong

acid or base in solution.

Each mole of H$2$SO$4$ dissociates to produce two charge equivalents of protons, and

each mole of Ca$($OH$)2$ and Mg$($OH$)2$ dissociate to provide two equivalents of hydroxide

ion. The concentrations of the respective acid anion and base cations must therefore be

multiplied by two. The square brackets refer to the molar concentration scale with units

of mole litre$-1$

$[$Acy$]=2[$SO$42-]+[$NO$3-]+[$CI$-]-[$Na$+]-[$K$+]-2[$Mg$2+]-2[$Ca$2+](2\mathrm{.}20)$

$[$Alk$]=[$Na$+]+[$K$+]+2[$Mg$2+]+2[$Ca$2+]-2[$SO$42-]-[$NO$3-]-[$CI$-](2\mathrm{.}21)$

$[$Acy$]=-[$Alk$]$

Weak acids: the CO$2-$H$2$O buffer system

Unlike strong acids, weak acids do not dissociate completely. Equation $(2\mathrm{.}23)$ shows

the dissociation of carbonic acid.

H$2$CO$3$ HCO$3-+$ H$+(2\mathrm{.}23)$

A common shorthand for this reaction represents the left hand side as "carbonic acid"

with an asterisk. This is because the equilibrium for the reaction CO$2($aq$)+$ H$2$O $$

H$2$CO$3$ lies far to the left and true carbonic acid $($H$2$CO$3)$ is therefore present in

negligible concentrations compared to dissolved CO$2($ aq$)).$

The dissociation constant for the deprotonation of H$2$CO$3*$ reflects both the hydration of

CO$2($aq$)$ and the deprotonation of true H$2$CO$3$ to form HCO$3$

$-($Equation $2\mathrm{.}24).$ H$2$CO$3*$ HCO$3-+$ H$+-$IogK$1=6\mathrm{.}35(25\backslash$deg C$)(2\mathrm{.}24)$

The principle of thermodynamic mass action can be applied to a general stoichiometric

reaction: aA $+$ bB $+$ cC $+$ nN $+$ pP $+$ rR $+...,$

with an equilibrium constant Keq that has been corrected for temperature and the ionic

composition of the solution. The brackets indicate units of molar concentration $($moles L$-$

l$)$ for reactants $($A$,$B$,$C$,$ etc$)$ and products $($N$,$P$,$R$,$etc.$).$

K$=(2\mathrm{.}26)$

Although there are a large number of solutes which act as weak acids, these can often

be neglected in natural waters except for carbonic acid which forms when carbon

dioxide dissolves in water.

CO$2($g$)+$ H$2$O $$ H$2$CO$3*-$logK $=1\mathrm{.}47(25\backslash$deg C$)2\mathrm{.}27)$

For the dissolution of CO$2($g$)$ in water, and the dissociation of the resulting carbonic

acid, the following thermodynamic mass action laws follow from equations $(2\mathrm{.}27),(2\mathrm{.}24)$

and $(2\mathrm{.}25),$ respectively. Note that gases have concentrations expressed as a partial pressure. With units of atmospheres $($atm$).$ kindly present this detailed