THE HYDRONIUM ION: THE HYDRONIUM ION The proton does not actually exist in aqueous solution as a bare H+ ion. The proton exists as the hydronium ion (H3O+). Consider the acid-base reaction: HCO3- + H2O  H3O+ + CO32- Here water acts as a base, producing the hydronium ion as its conjugate acid. For simplicity, we often just write this reaction as: HCO3-  H+ + CO32-


Conjugate Acid-Base pairs: 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+


AMPHOTERIC SUBSTANCE: 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 + H2CO30


Strong Acids/ Bases: Strong Acids/ Bases Strong Acids more readily release H+ into water, they more fully dissociate H2SO4  2 H+ + SO42- Strong Bases more readily release OH- into water, they more fully dissociate NaOH  Na+ + OH- Strength DOES NOT EQUAL Concentration!


Acid-base Dissociation: 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!


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


Organic acids in natural waters: 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 (pKa>3.6) and humin which is more insoluble Soil fulvic acids also strongly complex metals and can be an important control on metal mobility


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 gH+?? Turns out to be quite small  0.002 or so… How would you determine this??


pH: 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-


The glass: The glass Corning 015 is 22% Na2O, 6% CaO, 72% SiO2 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- H+Gl- H+Gl- H+Gl- H+Gl- H+Gl- H+Gl- H+Gl- Na+Gl- Na+Gl- E1 E2 Analyte solution Reference solution


Slide11: pH = - log {H+}; glass membrane electrode pH electrode has different H+ activity than the solution SCE // {H+}= a1 / glass membrane/ {H+}= a2, [Cl-] = 0.1 M, AgCl (sat’d) / Ag ref#1 // external analyte solution / Eb=E1-E2 / ref#2 E1 E2 H+ gradient across the glass; Na+ is the charge carrier at the internal dry part of the membrane soln glass soln glass H+ + Na+Gl-  Na+ + H+Gl-


Slide12: Values of NIST primary-standard pH solutions from 0 to 60 oC pH = - log {H+} K = reference and junction potentials


pKx?: pKx? Why were there more than one pK for those acids and bases?? H3PO4  H+ + H2PO4- pK1 H2PO4-  H+ + HPO42- pK2 HPO41-  H+ + PO43- pK3


BUFFERING: 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:: 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


Buffering example: 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


Slide17: 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


Bjerrum Plots: 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


Slide19: 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.


Slide20: 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!


Titrations: 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?


Carbonate System Titration: Carbonate System Titration From low pH to high pH


Titrations  precipitate: Titrations  precipitate


BJERRUM PLOT - CARBONATE: 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 CO2-H2O system: H2CO3*, HCO3-, CO32-, H+, and OH-. At each pK value, conjugate acid-base pairs have equal concentrations. At pH pK2, CO32- is predominant.


Slide25: 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!