dissociation and ph dissociation of weak acids/bases controlled by ph knowing the total amount of s...
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1
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
2
Hydrogen Sulfide Activity Diagram
3
Hydrogen Sulfide Activity Diagram
4
Solubility of Quartz
• The oxides of many metals react with H2O to form bases
• SiO2(s) + 2H2O H4SiO4°
5
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
6
Quartz Activity Diagram
7
2 4 6 8 10 12 14–6
–5
–4
–3
–2
–1
0
pH
log
a S
iO2(
aq
)
SiO2(aq)
H2SiO4--
H3SiO4-
Quartz
25°C
Walt Tue Feb 14 2006
Dia
gram
SiO
2(aq)
, T
=
25
°C
, P
=
1.
013
bars
, a
[H2O
] =
1;
Sup
pres
sed:
H4(H
2SiO
4) 4----
H4SiO4
8
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
9
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
10
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
11
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
12
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
13
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
14
DIC and pH in Open System
• What is the pH of natural rainwater?– Controlled by DIC equilibrium
•
– At 25°C, KCO2 = 10-1.47
15
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+ ↔ CO3
2- + H+
– At pH = 6.35, [H2CO3] = [HCO3-]
– At pH = 10.33, [HCO3-] = [CO3
2-]
• We can do same calculations we did for H2S
16Walt Tue Feb 21 2006
2 3 4 5 6 7 8 9 10 11 12–16
–14
–12
–10
–8
–6
–4
–2
0
pH
Sp
eci
es
with
HC
O3- (
log
mo
lal)
CO2(aq) CO
3--
HCO3-
Total DIC = 10-1 M
pH = 6.35 pH = 10.33
Common pH rangein natural waters
17
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
18
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-]
19
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+ + H2SiO4
2-, pKa = 13.28
20
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
21
Alkalinity
• Total alkalinity = [HCO3-] + 2[CO3
2-] + [B(OH) 4-] +
[H3SiO4-] + [HS-] + [OH-] – [H+]
– Typically in groundwater, [HCO3-] >> [CO3
2-], [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[CO3
2-] + [OH-] – [H+]– Directly convertible to [HCO3
-] when it is >> than others
• Measured by titration of solution with strong acid
22Walt Tue Feb 21 2006
2 3 4 5 6 7 8 9 10 11 12–16
–14
–12
–10
–8
–6
–4
–2
0
pH
Sp
eci
es
with
HC
O3- (
log
mo
lal)
CO2(aq) CO
3--
HCO3-
Total DIC = 10-1 M
23
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
24
Determining Alkalinity by Titration
InitialpH = 8.26
RapidpH change
RapidpH change
Slow pH change:Buffered
Determine maximum pH change by: ΔpH ÷ mL acid added
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