Dissociation and pH
• Dissociation of weak acids/bases controlled by
pH
• Knowing the total amount of S and pH, we can
calculate activities of all species and generate
curves
• Example: H2S
1
Hydrogen Sulfide Activity Diagram
2
Hydrogen Sulfide Activity Diagram
3
Solubility of Quartz
• The oxides of many metals react with H2O to
form bases
• SiO2(s) + 2H2O  H4SiO4°
4
Quartz Activity Diagram
• When including a solid, the activity diagram
looks a little different
– Showing fields of stability for each species
• Note: we don’t need to define initial log[SiO2]
concentration
– Activity of solid = 1
5
Quartz Activity Diagram
6
log a SiO2(aq)
–3
–6
H2SiO4
--
–4
H3SiO4
-
–5
SiO2(aq)
H4SiO4
25°C
2
4
6
8
pH
10
12
Walt T ue Feb 14 2006
----
–2
Diagram SiO 2(aq), T = 25 °C , P = 1.013 bars, a [H 2O] = 1; Suppressed: H 4(H 2SiO 4) 4
0
–1
Quartz
14
7
Buffering of pH
• Weak acids and bases can buffer pH of a
solution
– pH changes very little as acid (or base) is added
– Need both a protonated and unprotonated
species present in significant concentrations
• e.g., H2CO3(aq) and HCO3• Carbonic acid-bicarbonate is the major buffer in most
natural waters
• Organic acids and sometimes silicic acid can be
important buffers
8
pH Buffering capacity of an aquifer:
Minerals as well as aqueous species
• Reactions with minerals: carbonate most
important, fastest
– CaCO3 + H+ ↔ Ca2+ + HCO3-
• Silicates, slower, less important
– 2KAlSi3O8 + 2H2CO3 + 9H2O  Al2Si2O5(OH)4 + 2K+
+ 4H4SiO4 + 2HCO3• H2CO3 consumes acid, HCO3- creates alkalinity
• Ion exchange of charge surfaces
– Negatively charged S- + H+ ↔ SH
9
Dissolved Inorganic Carbon (DIC)
• Initially, DIC in groundwater comes from CO2
– CO2 (g) + H2O ↔ H2CO3°
• Equilibrium expression with a gas is known as
Henry’s Law
– PCO2: partial pressure (in atm or bar); pressure in
atmosphere exerted by CO2
– Assuming atmospheric pressure of 1 atm, PCO2 = 10-3.5;
concentration of CO2 = 350 ppm
• At atm = 1, N2 is 78%, PN2 = 0.78, O2 21%, PO2 = 0.21
10
Dissolved Inorganic Carbon (DIC)
• PCO2 of soil gas can be 10-100 times the PCO2 of
atmosphere
• PCO2 for surface water controlled by atmosphere
and biological processes
– Photosynthesis (day): drives PCO2 down, less H2CO3,
pH increases
• 6CO2 + 6H2O + Energy ↔ C6H12O6 + 6O2
– Respiration: increases PCO2, more H2CO3, pH drops
11
Dissolved Inorganic Carbon (DIC)
• In groundwater, no photosynthesis, no diurnal
variations
– CO2 usually increases along a flow path due to
biodegradation in a closed system
– CH2O + O2  CO2 + H2O
• CH2O = generic organic matter
12
DIC and pH in Open System
• CO2 can be dissolved into or volatilize out of
water freely
– Surface waters
• PCO2 is constant = 10-3.5 atm at Earth’s surface
13
DIC and pH in Open System
• What is the pH of natural rainwater?
– Controlled by DIC equilibrium
•
– At 25°C, KCO2 = 10-1.47
14
DIC and pH in Closed System
• In a closed system (no CO2 exchange), for a
given amount of TIC, speciation is a function
of pH
• CO2 + H2O ↔ H2CO3 ↔ HCO3- + H+ ↔ CO32+ H+
– At pH = 6.35, [H2CO3] = [HCO3-]
– At pH = 10.33, [HCO3-] = [CO32-]
• We can do same calculations we did for H2S
15
Total DIC = 10-1 M
0
-
CO2 (aq)
--
HCO3
CO3
pH = 10.33
pH = 6.35
–4
–6
-
Species with HCO3 (log molal)
–2
–8
–10
Common pH range
in natural waters
–12
–14
–16
2
3
4
5
6
7
8
9
10
11
12
pH
Walt Tue Feb 21 2006
16
Rainwater pH and PCO2
• What if we double PCO2 (10-1.75 atm)
– [H2CO3] = [10-1.47] [10-1.75] = 10-3.22
–
• Doubling the PCO2 does not have a large effect on pH
• Acid rain can have pH < 4
– Due to other acids (nitric and sulfuric) that are injected
into the atmosphere by vehicles and smokestacks
17
Special points about DIC, pH, and
other weak acids
• At pH 6.35, Ka1 = [H+], therefore [H2CO3] =
[HCO3-]
–
–
• Likewise, at pH 10.33, Ka2 = [H+], therefore
[HCO3-] = [CO32-]
18
Special points about DIC, pH, and
other weak acids
• When pH = pKa, concentration of protonated in
reactant = deprotonated in product
– pKa = -log Ka
– for H2CO3 ↔ HCO3- + H+, Ka = 10-6.35, pKa = 6.35
– so for H4SiO4 ↔ H3SiO4- + H+, pKa = 9.71
– And for H3SiO4-  H+ + H2SiO42-, pKa = 13.28
19
Alkalinity
• Alkalinity = acid neutralizing capability (ANC)
of water
– Total effect of all bases in solution
– Typically assumed to be directly correlated to
HCO3- concentration in groundwater
20
Alkalinity
• Total alkalinity = [HCO3-] + 2[CO32-] + [B(OH) 4-] +
[H3SiO4-] + [HS-] + [OH-] – [H+]
– Typically in groundwater, [HCO3-] >> [CO32-], [B(OH) 4-],
[H3SiO4-], [HS-], [OH-], [H+]
– Whenever there are significant amounts of any of these
other species, they must be considered
• Carbonate alkalinity = [HCO3-] + 2[CO32-] + [OH-] – [H+]
– Directly convertible to [HCO3-] when it is >> than others
• Measured by titration of solution with strong acid
21
Total DIC = 10-1 M
0
-
CO2 (aq)
--
HCO3
CO3
–4
–6
-
Species with HCO3 (log molal)
–2
–8
–10
–12
–14
–16
2
3
4
5
6
7
8
9
10
11
12
pH
Walt Tue Feb 21 2006
22
Alkalinity Titration
• Determine end-point pH:
– The pH at which the rate of change of pH per
added volume of acid is at a maximum
– Typically in the range 4.3-4.9
– Function of ionic strength
– Reported as mg/L CaCO3
–
– HCO3- = alkalinity
0.82
23
Determining Alkalinity by Titration
Initial
pH = 8.26
Rapid
pH change
Rapid
pH change
Slow pH change:
Buffered
Determine maximum pH change by: ΔpH ÷ mL acid added
24