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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 (H 3 O + ). Consider the acid-base reaction: HCO 3 - + H 2 O  H 3 O + + CO 3 2-

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THE HYDRONIUM ION

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  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 (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-

  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 • H2CO3 HCO3- + H+

  3. 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

  4. 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!

  5. 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!

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

  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 (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

  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+]=103 = 1000 molal? • But what’s gH+?? 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% 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+ E1 E2 glass H+Gl- H+Gl- Na+Gl- H+Gl- H+Gl- Reference solution Analyte solution H+Gl- H+Gl- Na+Gl- H+Gl- H+Gl-

  11. pH = - log {H+}; glass membrane electrode 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- pH electrode has different H+ activity than the solution E1 E2 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

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

  13. 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

  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 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

  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 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

  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 CO2-H2O system: H2CO3*, HCO3-, CO32-, H+, and OH-. • At each pK value, conjugate acid-base pairs have equal concentrations. • At pH < pK1, H2CO3* is predominant, and accounts for nearly 100% of total carbonate. • At pK1 < pH < pK2, HCO3- is predominant, and accounts for nearly 100% of total carbonate. • At pH > pK2, CO32- 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!

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