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-
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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.
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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+ + CO3
2-
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+ + CO3
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+
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+ + CO3
2-
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!
][
]][[
AH
HAK
x
x
• 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• 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• 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…– How would you determine this??
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• 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
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