Transcript Lecture 7 - Acid-base chemistry

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+ + CO32Here 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
• 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+ + OHIn 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+ + CO32Base: HCO3- + H3O+  H2O + H2CO30
Strong Acids/ Bases
• Strong Acids more readily release H+ into
water, they more fully dissociate
– H2SO4  2 H+ + SO42-
• Strong Bases more readily release OHinto 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…)
 x
[ A][ H ]
K
[ H x A]
• 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
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 gH?? 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 ionexchange 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+
E1
H+GlAnalyte solution
E2
H+Gl-
Na+Gl-
H+Gl-
H+Gl-
H+Gl-
glass H+Gl-
+ H+Gl- Na Gl H+Gl-
Reference solution
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
pH = - log {H+}
K = reference and
junction potentials
Values of NIST primary-standard
pH solutions from 0 to 60 oC
pKx?
• Why were there more than one pK for
those acids and bases??
• H3PO4  H+ + H2PO4pK1
• H2PO4-  H+ + HPO42pK2
• HPO41-  H+ + PO43pK3
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:

[A ]
pH  pK  log
[ HA]
• 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
• 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??
[A ]
pH  pK  log
[ HA]
– Write the pH changes via the Henderson-Hasselbach
equation
8
8.5
• 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-):
7.5
pH
7
6.5
–
pH=6.35+log (HCO3-/H2CO3) = 6.35+log(50/50)
6
• After 20 mmoles albite react (80 mmoles OH-):
5.5
–
pH=6.35+log(80/20)
=40 6.35 50+ 0.6 60= 6.9570
5
0
10
20
30
80
90
100
Albite reacted (mmoles)
Greg Mon Oct 11 2004
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
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.
-2
H2S0
7.0
13.0
-
HS
-4
S2-
log ai
-6
-8
OH-
-10
H+
-12
0
2
4
6
pH
8
10
12
14
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.
-2
Common pH
range in nature
6.35
H2CO3*
-3
-
HCO3
10.33
2-
CO3
log ai
-4
OH-
-5
+
H
-6
-7
-8
0
2
4
6
8
10
12
14
pH
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?
-2
Common pH
range in nature
6.35
H2CO3*
-3
-
HCO3
10.33
2-
CO3
log ai
-4
OH-
-5
+
H
-6
-7
-8
0
2
4
6
8
pH
10
12
14
Carbonate System Titration
11
10
9
pH
8
7
6
5
4
3
2
0
5
10
15
20
25
30
35
40
45
50
NaOH reacted (mmoles)
Greg Wed Oct 06 2004
-2
-
CO2(aq)
--
HCO3
CO3
-4
-6
-
Some species w/ HCO3 (log activity)
• From low
pH to high
pH
12
-8
-10
-12
-14
-16
0
5
10
15
20
25
30
35
40
45
50
NaOH reacted (mmoles)
Greg Wed Oct 06 2004
Titrations  precipitate
Some minerals (log moles)
-3.5
-4
-4.5
Boehmite
-5
-5.5
Fe(OH) 3(ppd)
-6
2
2.5
3
3.5
4
4.5
5
5.5
6
6.5
7
pH
Greg Wed Oct 06 2004
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.
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.
-2
Common pH
range in nature
6.35
H2CO3*
-3
-
HCO3
10.33
2-
CO3
log ai
-4
OH-
-5
+
H
-6
-7
-8
0
2
4
6
8
10
12
14
pH
In most natural waters, bicarbonate is the dominant carbonate species!