Acid-Base Geochemistry
description
Transcript of Acid-Base Geochemistry
Acid-Base Geochemistry• Arrhenius’ definition:
– Acid any compound that releases a H+ when dissolved in water
– Base any compound that releases an OH- when dissolved in water
• Bronstead-Lowry’s definition:– Acid donates a proton– Base receive/accept a proton
• Lewis’ definition:– Acid electron pair donor acceptor– Base electron pair donor
Conjugate Acid-Base pairs
• Generalized acid-base reaction:HA + B A + HB
• A is the conjugate base of HA, and HB is the conjugate acid of B.
• More simply, HA A- + H+
HA is the conjugate acid, A- is the conjugate base
• H2CO3 HCO3- + H+
Hydrolysis
• Mz + H2O M(OH)z-1 + H+
• Reaction of a cation, which generates a H+ from water is a hydrolysis reaction
• Described by the equilibrium constant Ka
• Hydrolysis also describes an organic reaction in which the molecule is cleaved by reaction with water…
zz
a MHMOHK ]][[ 1
AMPHOTERIC SUBSTANCE• Now consider the acid-base reaction:
NH3 + H2O NH4+ + OH-
In this case, water acts as an acid, with OH- its conjugate base. Substances that can act as either acids or bases are called amphoteric.
• Bicarbonate (HCO3-) is also an amphoteric
substance:Acid: HCO3
- + H2O H3O+ + CO32-
Base: HCO3- + H3O+ H2O + H2CO3
0
Strong Acids/ Bases
• Strong Acids more readily release H+ into water, they more fully dissociate– H2SO4 2 H+ + SO4
2-
• Strong Bases more readily release OH- into water, they more fully dissociate– NaOH Na+ + OH-
Strength DOES NOT EQUAL Concentration!
Acid-base Dissociation• For any acid, describe it’s reaction in water:
– HxA + H2O x H+ + A- + H2O– Describe this as an equilibrium expression, K (often
denotes KA or KB for acids or bases…)
• Strength of an acid or base is then related to the dissociation constant Big K, strong acid/base!
• pK = -log K as before, lower pK=stronger acid/base!
][]][[
AHHAKx
x
• LOTS of reactions are acid-base rxns in the environment!!
• HUGE effect on solubility due to this, most other processes
Geochemical Relevance?
Dissociation of H2O
• H2O H+ + OH-
• Keq = [H+][OH-]• log Keq = -14 = log Kw
• pH = - log [H+]• pOH = - log [OH-]• pK = pOH + pH = 14
• If pH =3, pOH = 11 [H+]=10-3, [OH-]=10-11
][]][[
2OHOHHKeq
Definition of pH
pH• Commonly represented as a range between
0 and 14, and most natural waters are between pH 4 and 9
• Remember that pH = - log [H+]– Can pH be negative?– Of course! pH -3 [H+]=103 = 1000 molal?– But what’s ?? Turns out to be quite small
0.002 or so…
pKx?
• Why were there more than one pK for those acids and bases??
• H3PO4 H+ + H2PO4- pK1
• H2PO4- H+ + HPO4
2- pK2
• HPO41- H+ + PO4
3- pK3
BUFFERING
• When the pH is held ‘steady’ because of the presence of a conjugate acid/base pair, the system is said to be buffered
• In the environment, we must think about more than just one conjugate acid/base pairings in solution
• Many different acid/base pairs in solution, minerals, gases, can act as buffers…
Henderson-Hasselbach Equation:
• When acid or base added to buffered system with a pH near pK (remember that when pH=pK HA and A- are equal), the pH will not change much
• When the pH is further from the pK, additions of acid or base will change the pH a lot
][][log
HAApKpH
Buffering example
• Let’s convince ourselves of what buffering can do…
• Take a base-generating reaction:– Albite + 2 H2O = 4 OH- + Na+ + Al3+ + 3 SiO2(aq)
– What happens to the pH of a solution containing 100 mM HCO3- which starts at pH 5??
– pK1 for H2CO3 = 6.35
• Think of albite dissolution as titrating OH- into solution – dissolve 0.05 mol albite = 0.2 mol OH-
• 0.2 mol OH- pOH = 0.7, pH = 13.3 ??
• What about the buffer??– Write the pH changes via the Henderson-Hasselbach
equation
• 0.1 mol H2CO3(aq), as the pH increases, some of this starts turning into HCO3-
• After 12.5 mmoles albite react (50 mmoles OH-):– pH=6.35+log (HCO3-/H2CO3) = 6.35+log(50/50)
• After 20 mmoles albite react (80 mmoles OH-):– pH=6.35+log(80/20) = 6.35 + 0.6 = 6.95
][][log
HAApKpH
Greg Mon Oct 11 2004
0 10 20 30 40 50 60 70 80 90 1005
5.5
6
6.5
7
7.5
8
8.5
Albite reacted (mmoles)
pH
Bjerrum Plots
• 2 D plots of species activity (y axis) and pH (x axis)
• Useful to look at how conjugate acid-base pairs for many different species behave as pH changes
• At pH=pK the activity of the conjugate acid and base are equal
pH0 2 4 6 8 10 12 14
log
a i
-12
-10
-8
-6
-4
-2H2S
0HS-
S2-
H+OH-
7.0 13.0
Bjerrum plot showing the activities of reduced sulfur species as a function of pH for a value of total reduced sulfur of 10-3 mol L-1.
pH0 2 4 6 8 10 12 14
log
a i
-8
-7
-6
-5
-4
-3
-2
6.35 10.33H2CO3* HCO3- CO3
2-
H+
OH-
Common pHrange in nature
Bjerrum plot showing the activities of inorganic carbon species as a function of pH for a value of total inorganic carbon of 10-3 mol L-1.
In most natural waters, bicarbonate is the dominant carbonate species!