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.

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