Module 8 (incomplete)

Question 1

Explain how adding CO2 or CO32- changes pH of aqueous systems and list common sources of these
compounds to natural waters

Answer 1

The carbonate system controls the pH of oceans and surface water.
An open system (ocean/surface water) will interact with atmospheric CO2

Question 2

What equations are involved in the open carbonate system?

Answer 2

CO2 dissolution
CO2 + H2O -> H2CO3

carbonic acid dissociation
1. H2CO3 -> H+ + HCO3 -
2. HCO3- -> H+ + CO3 2-
pka1=6.35
pka2=10.33

other:
H2O -> H+ + OH-

Question 3

Assumptions for Carbonate problems

Answer 3

pH < 9, ignore CO32- in charge balance
-basic means CO3-2 is low
If pH < 7, ignore OH- in charge balance
-acidic means OH- is low

system in equilibrium with atmosphere
-> atmospheric CO2

-CO2(aq) converted to H2CO3(aq)
-can use in henrys law
[H2CO3] = Pco2 + KH
-ideal gas so Pco2=[co2]

Question 4

Calculate the pH of water open to the atmosphere, in equilibrium with dissolution of a carbonate
mineral

Answer 4

Ksp = [Ca 2+][co3 2-] = 10^-8.48

Charge balance & assumptions give:
[Ca 2+]=1/2[HCO3 -]

rearrange dissociation 1 to get:
[HCO3 -]=Ka1[H2CO3] /[H+]

rearrange dissociation 2 to get:

[CO3 2-] = Ka2 [HCO3 -] /[H+]

use henrys law to get:
[H2CO3]=[CO2]KH

Resulting Equation:

Ksp = 1/2Ka1^2[CO2]^2KH^2Ka2/[H+]^3

solve for H+:

[H+] = (1/2Ka1^2[CO2]^2KH^2Ka2) ^1/3

Question 5

Explain how alkalinity helps buffer natural waters. Calculate alkalinity given appropriate data (either acid and base concentrations or titration data).

Answer 5

Alkalinity
-Ability of a water system to neutralize acid
-Higher alkalinity = higher tolerance to acid rain, acid mine drainage, etc.
-Measures bases which can accept a H+
-controlled by carbonates in natural systems

Carbonate alkalinity, meq/L
= [HCO3−] + 2[CO32−]

Total Alkalinity includes OH and H+ but they are low so we ignore
also includes other weak bases minus weak acids

-[HCO3−] and [CO32−] are measured in mmol/L

Alkalinity titrations
-measure alkalinity without measuring acids and bases separately
alkalinity of a water:
-titrate with a strong acid (e.g., HCl)

-measure pH and look for inflection
points rather than use indicators

At equivalence point:
meq/L titrant × V(mL) titrant
= meq/L sample × V(mL) sample

Concentration (meq/L) sample = alkalinity of sample

Question 6

Give the range of pH in unpolluted natural waters and what controls the extremes of the range.
Explain why some natural waters experience a 24-cycle in pH

Answer 6

pH increase
free CO2

pH decrease
CO2 being used

pH decrease during the day when photosynthesis is happening
pH increase at night when photosynthesis stops

Question 7

pH in Freshwaters

Answer 7

controls many other things in freshwater:
 Solubility and weathering of minerals
-acid dissolution and hydrolysis
 Speciation of weak acids and conjugate bases
 Shifts equilibria of many other reactions
 Many aquatic species have a narrow pH they can tolerate

pH=-log[H+]
0-7 acidic
8-14 basic

higher H+ = lower pH = more acidic

Question 8

Acid Dissociation Constants (Ka)

strong and weak acids

Major anions in fresh and marine water

Answer 8

An acid (HA) can donate H+ to aqueous solutions

Strong acid: Undergoes significant dissociation, Ka is very high, pKa is low (< 2)
 E.g., HCl, H2SO4, HNO3

Weak acid: Undergoes little dissociation, Ka is low, pKa is high
 E.g., acetic acid (CH3COOH)
 H2CO3 (carbonic acid)
 H4SiO4 (silicic acid)

pKa = -logKa

major anions:
Cl-, SO42-, HCO3-

At pH 6.35, half H2CO3 has dissociated into H+ and HCO3-. Acts as buffer in
natural waters

pH range of natural freshwaters: 4 – 10