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
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-
Values of NIST primary-standard pH solutions from 0 to 60 oC
pH = - log {H+}
K = reference and junction potentials
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
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:
• 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 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!
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?
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
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
-
Titrations precipitate
Greg Wed Oct 06 2004
2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7-6
-5.5
-5
-4.5
-4
-3.5
pH
Som
e m
iner
als
(log
mol
es)
Fe(OH)3(ppd)
Boehmite
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.
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!