WATER QUALITY IN STREAMS AND
RIVERS IS THE END PRODUCT OF ALL
PROCESSES IN THE BASIN
WATERSHEDS ARE THE
KIDNEYS OF AN ECOSYSTEM
KIDNEY ANALOGOUS
TO A WATERSHED
NITRATE EXAMPLE
Fingerprint water
Isotopes
Geochemical content
Nutrients
Rock
Weathering
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LITHOSPHERE
• Linkage between the atmosphere and
the crust. Weathering results in:
• Igneous rocks + acid volatiles =
sedimentary rocks + salty oceans
IMPORTANCE OF ROCK
WEATHERING
[1] Bioavailability of nutrients that have no
gaseous form:
– P, Ca, K, Fe
• Forms the basis of biological diversity,
soil fertility, and agricultural productivity
• The quality and quantity of lifeforms and
food is dependent on these nutrients
IMPORTANCE OF ROCK
WEATHERING
[2] Buffering of aquatic systems
-Maintains pH levels
-regulates availability of Al, Fe, PO4
Example: human blood.
-pH highly buffered
-similar to oceans
IMPORTANCE OF ROCK
WEATHERING
[3] Forms soil
[4] Regulates Earths climate
[5] Makes beach sand!
NATURAL ACIDS that
WEATHER ROCK
• Produced from C, N, and S gases in
the atmosphere
• H2CO3 Carbonic Acid
• HNO3 Nitric Acid
• H2SO4 Sulfuric Acid
• HCl
Hydrochloric Acid
Stoichiometry
Stoichiometry is the accounting, or math, behind chemistry.
Given enough information, one can use stoichiometry
to calculate masses, moles, and percents within a chemical equation.
Keep track of atoms, molecules, and charge
Calcite
dissolution
CaCO3 + CO2 + H2O  Ca2+ + 2 HCO3reactants
products
TYPES OF CHEMICAL
WEATHERING
•
•
•
•
Carbonate weathering
Dehydration
Oxidation
Hydrolysis
CARBONIC ACID
Carbonic acid is produced in rainwater by
Reaction of the water with carbon dioxide
Gas in the atmosphere.
CARBONATE
(DISSOLUTION)
All of the mineral is completely
Dissolved by the water.
Congruent weathering.
DEHYDRATION
Removal of water from a mineral.
OXIDATION
Reaction of minerals with oxidation.
An ion in the mineral is oxidized.
OXIDATION: REDOX
REACTIONS
Loss of electrons = Oxidation
Gain of electrons = Reduction
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For example, rusting
* Oxidation of elemental iron to iron(III) oxide by oxygen
4 Fe + 3 O2 → 2 Fe2O3
Occurs in nature as the mineral hematite,
and is the principal ore for iron
HYDROLYSIS
H+ replaces an ion in the mineral.
Generally incongruent weathering.
HYDROLYSIS
• Silicate rock + acid + water = base
cations + alkalinity + clay + reactive
silicate (SiO2)
HYDROLYSIS
• Base cations are
– Ca2+, Mg2+, Na+, K+
• Alkalinity = HCO3• Clay = kaolinite (Al2Si2O5(OH)4)
• Si = H4SiO4; no charge, dimer, trimer
Mineral Solubility
• Solubility - relative capability of being
dissolved
• Salt dissolution - solids break down in
solution to yield ions
• Example: Barium chloride BaCl2
BaCl2 (s) = Ba2+ + 2 Cl–
– Define K using the Law of Mass Action
(“activity” in brackets):
2
 2
[Ba ][Cl ]
K
BaCl 2 (s)
•Inside the [] are the measured concentrations
•Multiply [] by number of atoms
Solubility constant Ksp
– Because the activity of the solid is 1, the
equation becomes
Ksp = [Ba2+] · [Cl–]2
– The equilibrium constant for the dissolution
reaction is called the solubility product
constant or Ksp.
Measurements of
Disequilibrium
• It can be important to know whether a
solution is saturated or undersaturated with
respect to a mineral
• Consider:
AaBb = aA + bB
• At equilibrium: Ksp = [A]a [B]b
• How do we know the solution is in equilibrium
with the mineral? Measure [A] and [B] in
solution (activity product or ion activity
product) and compare to Ksp
Degree of saturation W
a
b
[A] [B]
W
Ksp
»where [A] and [B] are for the solution,
»which may or may not be in equilibrium with the mineral
» W > 1 Supersaturated
» W = 1 Saturated
» W < 1 Undersaturated
Problem: What is the degree of saturation
of anhydrite in College Station tap water?
• (Ca2+) = 3 mg/L = 0.003 g·L-1/40 g Ca·mol-1 =
0.000075 M
• (SO42-) = 10 mg/L = 0.010 g·L-1/96 g SO42-·mol-1 =
0.00010 M
• T = 25°C
• Assume ideal behavior (g = 1)
• Write the reaction in terms of dissolution and make
use of Ksp values
CaSO4 = Ca2+ + SO42-
• We calculate the ion activity product in
solution:
IAP = [Ca2+][SO42-] = 0.000075 · 0.00010 = 7.5
x 10–9 = 10–8.1
• Degree of saturation
8.1
IAP
10
3.74
W
 4.36  10
K sp, anh 10
Water is undersaturated with respect to annhydrite
Calcite dissolution:
CaCO3 = Ca2+ + CO32–
Ksp,calcite  [Ca2  ][CO23 – ]
Is water undersaturated or oversaturated with
respect to calcite?
Get stalagmites/stalagtites?
Or dissolve them?
Tea pots: where does mineral deposits come from?
But ions don’t behave ideally . . .
• Concentration related to activity using the activity
coefficient g, where
[z] = gz (z)
• The value of g depends on:
– Concentration of ions and charge in the solution
– Charge of the ion
– Diameter of the ion
• Ionic strength I = concentration of ions and
charge in solution
I = 1/2 Smizi2
– where mi = concentration of each ion in moles per
kg, zi = charge of ion
Activity and Concentration
• Activity – “effective concentration”
• Ion-ion and ion-H2O interactions (hydration
shell) cause number of ions available to react
chemically ("free" ions) to be less than the
number present
• Concentration can be related to activity using
the activity coefficient g, where [z] = gz (z)
• Activity coefficient gz  1 as concentrations  0 and
tend to be <1 except for brines
Carbonate
Chemistry
The Carbonate System
• pH of most natural waters controlled by
reactions involving the carbonate system
• Groundwater and seawater chemistries are
often poised near calcite equilibrium, with pH
buffered by calcite dissolution and
precipitation
• Applications
– Fate of CO2 from fossil fuels and other CO2
sources on the atmosphere
– Effect of acid rain on lakes
– Effect of acid mine drainage on rivers
Carbonate System
• Carbonate species are necessary for
all biological systems
• Aquatic photosynthesis is affected by
the presence of dissolved carbonate
species.
• Neutralization of strong acids and
bases
• Effects chemistry of many reactions
• Effects global carbon dioxide content
PCO2
= 10–3.5 yields pH = 5.66
»What is 10–3.5? 316 ppm CO2
What
is today’s PCO2? ~368 ppm = 10-3.43
»pH = 5.63
pH of Global Precipitation
DIPROTIC ACID SYSTEM
• Carbonic Acid (H2CO3)
– Can donate two protons (a weak acid)
• Bicarbonate (HCO3-)
– Can donate or accept one proton (can be
either an acid or a base
• Carbonate (CO32-)
– Can accept two protons (a base)
TOTAL CARBONATE SPECIES (CT)
OPEN SYSTEM
• Water is in equilibrium with the partial
pressure of CO2 in the atmosphere
• Useful for chemistry of lakes, etc
• Carbonate equilibrium reactions are
thus appropriate
First the CO2 dissolves according to:
(1) CO2 (g) ⇔ CO2 (l)
According to Henry’s Law, solubility increases
as water temperature decreases
Equilibrium is established between the
dissolved CO2 and H2CO3, carbonic acid.
(2) CO2 (l) + H2O (l) ⇔ H2CO3 (l)
Carbonic acid is a weak acid that dissociates in two steps.
(3) H2CO3 + H2O  H3O+ + HCO3
-
pKa1 (25 C) = 6.37
(4) HCO3- + H2O  H3O+ + CO32pKa2 (25 C) = 10.25
The above presen ted more schematically:
CO 2(g)

CO 2 (l)
+ H2O

H2CO 3
+ H2O

-
HCO 3
+
+ H3O
+ H2O

2+
2-
CO 3
+
+ H3O
+ Ca

Note that the revers e is also true and that the scheme represent s the solubility of CaCO3 in
an acidic solution resulting in the liberation of CO 2 in the atmosphere.
CaCO 3 
Activity of Carbonate Species versus pH
CARBONATE SPECIES AND pH
Carbonate Buffering: Humans
We
can describe the formation
and dissociation of carbonic acid
through the following chemical and
equilibrium equations
Carbonic
acid forms when CO2 dissolves
in and reacts with water:
CO2(g) + H2O = H2CO3
»Most dissolved CO2 occurs as “aqueous CO2”
rather than H2CO3, but we write it as carbonic
acid for convenience
»The equilibrium constant for the reaction is:
KCO2
[H2CO3 ]
[H2CO3 ]


PCO2 [H2O]
PCO2
»Note we have a gas in the reaction and use partial
» pressure rather than activity
»First dissociation:
H2CO3 = HCO3– + H+

–
3
[H ][HCO ]
K1 
[H2CO3 ]
FIRST REACTION
»Second dissociation:
HCO3– = CO32– + H+

2–
3
–
3
[H ][CO ]
K2 
[HCO ]
SECOND REACTION
Variables and Reactions Involved in
Understanding the Carbonate System
Gas
Dissociation of Dissociation Cations
equilibria carbonic acid
of water
PCO2
[H2CO3]
[H+]
[Ca2+]
[HCO3–]
[CO32–]
[OH–]
Measurements
DIC
Alkalinity
ALKALINITY refers to
water's ability, or inability, to
neutralize acids.
The terms alkalinity and total
alkalinity are often used to
define the same thing.
Total
alkalinity - sum of the bases in equivalents that
are titratable with strong acid (the ability of a solution
to neutralize strong acids)
Bases
which can neutralize acids in natural waters:
HCO3–, CO32–, B(OH)4–, H3SiO4–, HS–, organic acids
(e.g., acetate CH3COO–, formate HCOO–)
Carbonate alkalinity
• Alkalinity ≈ (HCO3–) + 2(CO32–)
• Reason is that in most natural waters,
ionized silicic acid and organic acids are
present in only small concentrations
• If pH around 7, then
– Alkalinity ≈ HCO3–
Bicarbonate
dominates
alkalinity of
sea water
Alk = OH– + HCO3– + 2CO3–2 + B(OH)4- - H+
Gran Titration for Acid
Neutralizing Capacity (total
alkalinity)
• This method determines ANC by
titration with 0.1 N Hydrochloric Acid
between the pH range of 4.5 and 3.5 at
which the contributions of organic acids,
carbonate and bicarbonate are
neutralized.
• Explicitly accounts for most organic
acids
Charge Balance
• Fundamental principle of solution
chemistry is that solutions are
electronically neutral
• Sum of positive charges must equal the
sum of negative charges in any sample
+ = -
• + > -, there is an unmeasured anion
• Equivalents: moles/L x valence
ION Percent Difference
Quality Control
•
•
•
•
•
+ - - / + + Normalizes the charge balance
NADP guidelines for allowable error
Error should be random and equal zero
If error always positive, means there is
an unmeasured anion (negative charge)
• Alkalinity is routinely measured in
natural water samples. By measuring
only two parameters, such as alkalinity
and pH, the remaining parameters that
define the carbonate chemistry of the
solution (PCO2, [HCO3–], [CO32–],
[H2CO3]) can be determined.
Alkalinity and hardness.
Calcium (Ca++) and Magnesium (Mg++)
are primarily responsible for hardness.
However, in most waters, alkalinity and
hardness have similar values because the
carbonates and bicarbonates responsible for
total alkalinity are usually in the form of
Calcium carbonate or Magnesium carbonate.
However, waters with high total alkalinity
are not always hard, since the carbonates
can be in the form of Sodium or Potassium
carbonate.
CLOSED CARBONATE
SYSTEM
• Carbon dioxide is not lost or gained to
the atmosphere
• Total carbonate species (CT) is constant
regardless of the pH of the system
• Occurs when acid-base reactions much
faster than gas dissolution reactions
• Equilibrium with atmosphere ignored
How does [CO3–2] respond to changes in Alk or DIC?
CT = [H2CO3*] + [ HCO3–] + [CO3–2]
~ [ HCO3–] + [CO3–2] (an approximation)
Alk = [OH–] + [HCO3–] + 2[CO3–2] + [B(OH)4-] – [H+]
~ [HCO3–] + 2[CO3–2] (a.k.a. “carbonate alkalinity”)
So (roughly):
[CO3–2] ~ Alk – CT
CT ↑ , [CO3–2] ↓
Alk ↑ , [CO3–2] ↑
Diurnal changes in DO and pH
What’s up?
Photosynthesis is the biochemical process in which plants and algae
harness the energy of sunlight to produce food. Photosynthesis of
aquatic plants and algae in the water occurs when sunlight acts on the
chlorophyll in the plants. Here is the general equation:
6 H20 + 6 CO2 + light energy —> C6H12O6 + 6 O2
Note that photosynthesis consumes dissolved CO2 and produces
dissolved oxygen (DO). we can see that a decrease in
dissolved CO2 results in a lower concentration of carbonic acid
(H2CO3), according to:
CO2 + H20 <=> H2CO3 (carbonic acid)
As the concentration of H2CO3 decreases so does the concentration
of H+, and thus the pH increases.
Cellular Respiration
Cellular respiration is the process in which organisms,
including plants, convert the chemical bonds of energy-rich molecules
such as glucose into energy usable for life processes.
The equation for the oxidation of glucose is:
C6H12O6 + 6 O2 —> 6 H20 + 6 CO2 + energy
As CO2 increases, so does H+, and pH decreases.
Cellular respiration occurs in plants and algae during the day and night,
whereas photosynthesis occurs only during daylight.
Rock Cycle
Hydrolysis
Carbonate weathering
Classic reference on geochemical weathering
%A R. M. Garrels
%A F. T. MacKenzie
%T Origin of the chemical composition of some springs
and lakes
%B Equilibrium Concepts in Natural Water Systems
%E R. G. Gould
%S Am. Chem. Soc. Adv. Chem. Ser.
%V 67
%P 222-242
%D 1967