1. THE HYDRONIUM ION The proton does not actually exist in aqueous solution as a bare H + ion. The proton exists as the hydronium ion (H 3 O + ). Consider the acid-base reaction: HCO 3 - + H 2 O  H 3 O + + CO 3 2- Here water acts as a base, producing the hydronium ion as its conjugate acid. For simplicity, we often just write this reaction as: HCO 3 -  H + + CO 3 2-

2. 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 H 2 CO 3  HCO 3 - + H +

3. AMPHOTERIC SUBSTANCE Now consider the acid-base reaction: NH 3 + H 2 O  NH 4 + + 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 (HCO 3 - ) is also an amphoteric substance: Acid: HCO 3 - + H 2 O  H 3 O + + CO 3 2- Base: HCO 3 - + H 3 O +  H 2 O + H 2 CO 3 0

4. Strong Acids/ Bases Strong Acids more readily release H+ into water, they more fully dissociate –H 2 SO 4  2 H + + SO 4 2- Strong Bases more readily release OH- into water, they more fully dissociate –NaOH  Na + + OH - Strength DOES NOT EQUAL Concentration!

5. Acid-base Dissociation For any acid, describe it’s reaction in water: –H x A + H 2 O  x H + + A - + H 2 O –Describe this as an equilibrium expression, K (often denotes K A or K B 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!

6. LOTS of reactions are acid-base rxns in the environment!! HUGE effect on solubility due to this, most other processes Geochemical Relevance?

7. Organic acids in natural waters Humic/nonhumic – designations for organic fractions, –Humics= refractory, acidic, dark, aromatic, large – generally meaning an unspecified mix of organics –Nonhumics – Carbohydrates, proteins, peptides, amino acids, etc. Aquatic humics include humic and fulvic acids (pK a >3.6) and humin which is more insoluble Soil fulvic acids also strongly complex metals and can be an important control on metal mobility

8. 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 + ]=10 3 = 1000 molal? –But what’s   ?? Turns out to be quite small  0.002 or so… –How would you determine this??

9. pH pH electrodes are membrane ion-specific electrodes Membrane is a silicate or chalcogenide glass Monovalant cations in the glass lattice interact with H + in solution via an ion- exchange reaction: H + + Na + Gl - = Na + + H + Gl -

10. The glass Corning 015 is 22% Na 2 O, 6% CaO, 72% SiO 2 Glass must be hygroscopic – hydration of the glass is critical for pH function The glass surface is predominantly H + Gl - (H + on the glass) and the internal charge is carried by Na + glass H + Gl - Na + Gl - E1E1 E2E2 Analyte solution Reference solution

11. pH = - log {H + }; glass membrane electrode pH electrode has different H + activity than the solution SCE // {H + }= a 1 / glass membrane/ {H + }= a 2, [Cl - ] = 0.1 M, AgCl (sat’d) / Ag ref#1 // external analyte solution / E b =E 1 -E 2 / ref#2 E1E1 E2E2 H + gradient across the glass; Na + is the charge carrier at the internal dry part of the membrane soln glass H + + Na + Gl -  Na + + H + Gl -

12. Values of NIST primary-standard pH solutions from 0 to 60 o C pH = - log {H + } K = reference and junction potentials

13. pK x ? Why were there more than one pK for those acids and bases?? H 3 PO 4  H + + H 2 PO 4 - pK 1 H 2 PO 4 -  H + + HPO 4 2- pK 2 HPO 4 1-  H+ + PO 4 3- pK 3

14. 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…

15. 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

16. Buffering example Let’s convince ourselves of what buffering can do… Take a base-generating reaction: – Albite + 2 H 2 O = 4 OH- + Na + + Al 3+ + 3 SiO 2(aq) –What happens to the pH of a solution containing 100 mM HCO3- which starts at pH 5?? –pK 1 for H 2 CO 3 = 6.35

17. 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 H 2 CO 3(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 - /H 2 CO 3 ) = 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

18. 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

19. 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.

20. 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!

21. Titrations When we add acid or base to a solution containing an ion which can by protonated/deprotonated (i.e. it can accept a H + or OH - ), how does that affect the pH?

22. Carbonate System Titration From low pH to high pH

23. Titrations  precipitate

24. BJERRUM PLOT - CARBONATE closed systems with a specified total carbonate concentration. They plot the log of the concentrations of various species in the system as a function of pH. The species in the CO 2 -H 2 O system: H 2 CO 3 *, HCO 3 -, CO 3 2-, H +, and OH -. At each pK value, conjugate acid-base pairs have equal concentrations. At pH < pK 1, H 2 CO 3 * is predominant, and accounts for nearly 100% of total carbonate. At pK 1 < pH < pK 2, HCO 3 - is predominant, and accounts for nearly 100% of total carbonate. At pH > pK 2, CO 3 2- is predominant.

25. 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!