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!

][

]][[

AH

HAK

x

x

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

• 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

][

]][[

2OH

OHHKeq

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…

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

HA

ApKpH

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…

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

HA

ApKpH

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

ai

-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

ai

-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!

Carbonate System Titration

• From low pH to high pH

Greg Wed Oct 06 2004

0 5 10 15 20 25 30 35 40 45 502

3

4

5

6

7

8

9

10

11

12

NaOH reacted (mmoles)

pH

Greg Wed Oct 06 2004

0 5 10 15 20 25 30 35 40 45 50-16

-14

-12

-10

-8

-6

-4

-2

NaOH reacted (mmoles)

So

me

sp

eci

es

w/

HC

O3- (

log

act

ivit

y) CO2(aq) CO3--HCO3

-