Let's consider water:

Water

"the universal solvent"

Why?

Bond angle
105° between hydrogens and oxygen creates a permanent dipole

Dielectric Constant: of 78.5 (twice that of all other common solvents)
Measure of the ability of a solvent to cause charges of opposite charge to separate and behave independently.

Will salt dissolve in oil?

Hydrogen bonding

1. causes solvation of neutral molecules
2. assists with solvation of molecules with permanent dipoles
3. stabilizes colloids
4. determines the tertiary shape of proteins and enzymes
5. stabilizes ionic double layers
6. assists in solvation of charged ions
7. responsible for surface tension

Sets the pH scale for aqueous systems from 0-14.
What would the pH scale be in say, glacial acetic acid?

Solubilizes all phases
Liquids, solids, and gases dissolve in water.

Equilibrium Constant:
Along with temperature it sets the equilibrium constant.
Temperature is always stated
Aqueous system is always assumed

Aqueous Chemistry

The complexity of the chemistry cannot be over emphasized.

Models are simplified systems to provide insight into the chemical reactions that occur.

Gasses in water (O₂, CO₂, CH₄ etc.)

All gasses in water are governed by Henry's Law - "the solubility of a gas in a liquid is proportional to the partial pressure of that gas in contact with the liquid" (calculations in Ch. 5)

Oxygen

Oxygen (O₂)is very important for higher aquatic life.

7-8 mg of organic matter will stoichiometrically consume all the O₂ in 1L of water.

Oxygen is the final electron acceptor in biological energy systems for most higher animal and fish life. Higher life cannot exist without it!

If O₂ concentration goes to Zero or become very small, fish and other marine organisms die.

Carbon Dioxide

CO₂ is an increasing atmospheric gas

A naturally occurring and biologically produced and sinks in natural waters

Organic matter decomposition is also a significant source of CO ₂

Natural Water Sinks for CO₂

In Oceans 1/3 of excess CO₂ produced by anthropogenic activity is absorbed by the Oceans

Also

Another important CO₂ sink is photosynthesis by plants, grasses, trees, weeds etc. in the lithosphere

Ref. In the Introduction to Env. Sci. text.
Living in the Environment, by Miller, 8th ed. Wadsworth Publishing Co. CA, 1994. Text pgs 95-96. - Carbon Cycle Fig. 4-28

Carbon Dioxide & Carbonate, pH & Alkalinity

Carbon dioxide is a primary reagent in natural water buffer systems.

In many surface water systems the primary source of CO₂ is the atmosphere at 0.035%. However, unlike O₂, in some systems the primary source of CO₂ may be the by product of biological reactions. Depending on the biosphere and the biological productivity.

The solubility of CO₂ is primarily due to the formation of HCO₃^- and CO₃^2- .

Greater than 25 mg/L of CO₂ in water will cause death to many aquatic life forms.

Different species (CO₂, HCO₃^- , CO₃^2-) predominate depending on the pH.

The typical concentration of [HCO₃^- ] is 1.0 x 10^-3 mol/liter.

The term Alkalinity is used to describe the particular water systems ability to accept H₃O⁺ or H⁺ ions (buffer capacity).
(sometimes called Total Alkalinity)

Alkalinity is usually due to natural buffers such as HCO₃ ^- (the predominate CO₂ species in natural water) and is usually 1.0 x 10^-3 mol/liter of water. (assume 1eq/mole) (also other we will investigate later such as PO₄^2-, SO₄^2-, and Natural hydroxides {Al(OH)₃})

Stratification of a lake.

Acids and Bases (review)

The Bronsted - Lowery definition (1923) -
"an acid is a substance capable of donating a proton and a base is a substance that can accept a proton."
(Skoog or M & R)

The Lewis definition (G. N. Lewis) -
"an acid is an electron - pair acceptor and a Lewis base is an electron- pair donor"
(McQuarrie and Rock {M & R})

1. Many 3 dimensional structures of enzymes and other important biological molecules are pH dependent.
2. Many (if not most) chemical reactions are pH dependent (remember your organic synthesis H⁺ and OH^- catalyzed reactions also Lewis defined acids and bases, etc.)
3. Environmentally may leaching situations depend on the pH for movement of ions and molecules. Some we will study, complexation, precipitation, neutralization, adsorption, protonation, ionization, etc.
4. Other examples of where pH is critical to other reactions or situations. Water soil interaction, solubility of particles in air (or more precisely water in air), shifting of equilibrium.

Two competing reactions occur simultaneously:

pH a brief review

The ion - product constant of water

K_w is the "ion product constant for water" usually given at 25 °C

Why doesn't the liquid water appear in the equilibrium?
The concentration of liquid H₂O(l) is relatively constant compared to the ions and it is a liquid and liquids and solids do not appear in equilibrium constants (H₂O(l) is a constant 55.6 mole)(and their activity is = to 1). A Proof is provided on pg. 126 Skoog 6th ed.

Given: hydrogen-ion concentration of 7.9 x 10^-6

Given: pH = 5.1

Thus: (also useful relationship)

The range of pH values for common aqueous solutions.

Titration curve for the titration of 50.0 mL of 0.100 M HCl(aq) with 0.100 M NaOH(aq).

Note that the pH of the solution changes very slowly until the vicinity of the equivalence point is reached (50.0 mL of NaOH(aq)) and then changes very rapidly in this vicinity.

Natural Waters & CO₂, Equilibrium:

Carbon Dioxide and its hydrated equilibrium components have a unique relationship with natural waters

But water in liquid, ([H₂O(l)], liquids, and solids not included)

Equilibrium Constant - K (K_eq)

K _eq =  = 4.45 x 10^-7     pK _a1 = [-LogK_a] = 6.35

K _eq =  = 4.69 x 10^-11     pK _a2 = [-LogK_a] = 10.33

Which equilibrium is most likely in natural water systems?? Why??

A distribution of species diagram for CO₂, HCO₃ ^-, CO₃^2- is given:

Alpha fractional species

The specie of carbonate found in most waters depends on the pH and since most natural waters are pH 6.8-8.2 HCO₃^- is the predominant form.

Substitution of K_a1 and K_a2 from the equilibrium into  yields a relationship as functions of

[Given: from air equilibrium, 0.035% of gas in air, Henry's Law of partial pressure]
Natural water typically has an alkalinity of 1 x 10^-3 mol/L

Total alkalinity (all contributing species) becomes:

Also as the carbon source for the Bioshpere:
While the carbonate system is the buffer for natural waters it is also the carbon source and a significant oxygen source when interacting with the algae in water.

The pH of Rain Water

Assuming no anthropogenic or natural pollution to the atmosphere,
What is the pH of rain water?

Given: One buffer present from gas Henry's Law of partial pressure

Alkalinity

Alkalinity - is the capacity of water to accept H⁺, (H₃O⁺, Hydronium ion)

Alkalinity (or total alkalinity) is the buffer capacity of water for acid.

This is also an indication of which streams are susceptible to acid rain and which have a buffer capacity sufficient to resist a pH change that could damage their ecosystems.

Alkalinity then becomes an important property of natural water stabilizing the pH for aquatic plant growth and aquatic life.

[alk] = [HCO₃^-] + 2[CO₃ ^2-] + [OH^-] - [H₃O⁺]

Thus the lower the pH (< 7) the low the alkalinity typically and the lower the buffer capacity of the water and less tolerant to acids.

Reservoirs of carbonate

Carbon dioxide-calcium carbonate equilibria.

Natural water in contact with a mineral calcium carbonate

Note: over simplified equation not balanced for H⁺ form water

Assuming: that [CO₃^2-(aq) = [H ₂CO₃(aq)]

Water in contact with solid CaCO₃ would have the calculated concentration of these species

These concentrations shift as the equilibrium relating them shift

Very useful for buffer calculations (a review)
(works in 90-95% of the time)

Take the log of both sides of the equation

Put log [H₃O⁺] on the left of the equation and Log Ka on the right
so that pH = -Log[H₃O⁺] or -Log[H⁺ ] and, -LogK_a = pK_a

The equation can be used for buffer calculations when:

• Note: This automatically makes the assumption that the contribution of the natural dissociation of water into hydronium and hydroxide ions concentrations is insignificant.
• Also the concentration of the buffer components should be >10^{-6 M}
  (Why? - - the larger the less the dissociation of water is significance.)

* Dilution of a buffer does not change the pH at which it buffers within several orders of magnitude of dilution

Example (Skoog Figure 10-3 show dilution of a strong acid, weak acid, and a buffer.) Several examples of buffer titration curves.

Alternate Derivation of the Henderson-Hasselbach Equation

Start with equilibrium equation and the equilibrium constant

Take the log of both sides of the equation (when, K_a and K_b < 10^-3 M)

Isolating the hydronium ion (it will become the pH)

(inversion of the Log changes the sign and it is usually written thus)

The equatiion can be used for buffer calculations when K_a of the acid involved is small (K_a < 10^-3 ).

Refs.
General Chem., by McQuarrie and Rock, Freeman Press, NY, pg. 677, 1991.
Analytical Chem., by Skoog et. al. 6th ed., Saunders, pg. 221, 1992.
Quantitative Analysis, by Day and Underwood, 6th ed., Prentice Hall, pg. 109, 1991.

Buffers

To have a buffer both the weak acid and its conjugate base must be present to constitute that buffer pair.

Buffer Calculations

There are several general forms for buffer calculations.

Buffer Capacity

Buffers resist changes in pH when significant quantities of both weak acid and conjugate base are present.

The amount of acid or base that can be added without changing significantly the pH depends on the concentration of the Buffer Pair.

The greater the concentration the more acid or base that can be tolerated.

Definition - Buffer capacity is the number of moles of strong acid or base that result in a 1 pH change of a 1.0 L buffer solution.
(Example Skoog - Figure 10-4 )

The point at which the buffer can take the maximum amount of added acid or base is where there is an equal amount of both acid and conjugate base in solution. This corresponds to the pK_a of the acid.

Buffers are generally calculated to buffer around their pK _a where the maximum tolerance to adjustment is. They are then adjusted with either their pair concentration or with acid or base to attain the exact buffer pH for specific purposes.

Rule of thumb- Choose a buffer with a pK_a within +/-1 pH unit of the desired buffer pH. If you want it to buffer there.

The shape of the titration curve is controlled by the completeness of the dissociation and the reaction with the strong acid or base.

Natural water in contact with a mineral calcium carbonate

CaCO₃(s) (calcite, limestone, shells, coral...)

This is mediated by the solubility of CaCO₃ K _sp = 4.8 x 10^-9

This is why the pH of sea water is around 8.2

What does concentration provide if not pH??
(Buffer Capacity or Total Alkalinity)

Buffers and their Importance in Natural Waters and Environmental Systems

What is a buffer? - - Buffers are solutions that contain both a weak acid and its conjugate base in solution at the same time.

If the first equilibrium lies far to the right and is dominant then the solution will be acidic.

If the second is dominant then the solution is basic.

Depending on the strength of these equilibrium, the solution will be acidic or basic.

The effect of dilution on the pH of buffered and unbuffered solutions. The dissociation constant for HA is 1.00 x 10^-4. Initial solute concentrations are 1.00 M.

Titration curve for the titration of 50.0 mL of 0.100 M CH₃COOH(aq) with 0.100 M NaOH(aq). The indicator used in this case is phenolphthalein, which changes from colorless to pink around pH=9. The relatively flat portion of the curve between 5 and 40 mL of added base illustrates graphically the CH₃COOH(aq)/CH₃COO^-(aq) buffer action.

Titration curve for the titration of 0.100 M H₃PO₄(aq) with 0.100 M NaOH(aq). Note the three equivalence points. The third equivalence point is not sharply defined because HPO₄^2-(aq) is a very weak acid, so HPO₄^3-(aq) can compete with H₂O(l) for protons. Note that the curve has three distinct sections: 0 to 50 mL, 50 mL to 100 mL, and 100 mL to 150 mL. In each section the behavior of pH versus volume of base is analogous to a titration curve for a monoprotic acid.

Titration curve for the titration of 50.0 mL of 0.100 M NH₃(aq) with 0.100 M HCl(aq). The pH at the equivalence point is acidic because at that point we have a 0.050 M solution of NH₄Cl(aq), and NH₄⁺(aq) is an acidic ion while Cl^-(aq) is a neutral ion.

Curve for the titration of acetic acid with sodium hydroxide. A: 0.1000 M acid with 0.1000 M base. B: 0.001000 M acid with 0.001000 M base.

Titration of 20 mL of 0.100 M H₂A with 0.100 M NaOH. For H₂A, K₁ = 1.00 x 10 ^-3 and K₂ = 1.00 x 10 ^-7. Method of pH calculations is shown for several points and regions on the titration curve.

Carbon Dioxide and Carbonate Species, pH and Alkalinity (cont.)

The term Acidity in natural waters refers to the ability of the water to neutralize OH^-.

Since the carbonate ion the bicarbonate ion and carbonic acid are week acid and salts of week acids they constitute a buffer and are largely responsible for the pH of natural waters.
(CO₂, HCO₃^-, CO₃ ^2-)

Other molecular species and ions contributing to alkalinity and pH:
NH₃, H₃PO₄, R-Carboxylic acids, H ₄SiO₄ (Silicic acid).

Each of these species can act as buffers and each produce pH stability and buffer capacity and contribute to the total alkalinity of the water solution.

When taken a group the term is frequently referred to as -
Total Alkalinity

Carbon dioxide-calcium carbonate equilibrium.

Strong and weak acid - base ionizations, buffer behavior, solubility and precipitation.

Acid and bases
Conjugate base is the entity that remains following the donation of a proton and a conjugate acid is what results after accepting a proton. (c- = conjugate)

Strong acid or bases dissociate in water completely and leave no undissociated molecules.

Weak acid or bases do not dissociate completely and leave varying amounts of both species in solution.

In an equilibrium that can easily be shifted

Remember LeChatelier Principle

What will shift this to more CH₃COOH form ??

Water is an Amphiprotic solvent acting as either a proton donor or acceptor.
H₃O⁺ , OH^-

Chemical Equilibrium:

(A brief review)

Equilibrium Constant - K (K_eq)

Notes:
Solution problems usually use concentrations of 1 M or less.
As the term implies the reactions go in both directions
Direction of the reaction and the equilibrium constant -

K_eq is a temperature dependent constant.

At dilute concentrations ( < or = to 1 x 10^-3 Molar) -

Did the Previous example meet these criteria ??

Review of important and common types of equilibrium constants.

Dissociation - Ion-Product constant
Example for Water:

Dissociation of a weak acid or base
Example K_a or K_b

Review Polyprotic Acids

(pH calculations simplified)

Assumptions of K₁ dominance are usually acceptable when the K₁ and K₂ are separated by 10³ or when K₁/K₂ 10³.

The pH of the weak acid can be calculated by the following
(ref. Skoog p 247).

This equation will take into account the influence of the hydrolysis of water.

Titration of polyprotic acids and bases can be treated as buffers between equivalence points.

An example of a diprotic acid is maleic acid titrated with NaOH.

The theoretical diprotic acid couple H₂A and HA ^- are considered using the derived relationship for polyprotic acids.

Titration of 20 mL of 0.100 M H₂A with 0.100 M NaOH. For H₂A, K₁ = 1.00 x 10 ^-3 and K₂ = 1.00 x 10 ^-7. Method of pH calculations is shown for several points and regions on the titration curve.

The derived relationship can be used in this example until the second equivalence point is reached and a hydrolysis problem is created.

Diprotic acids have 2 buffering pairs and triprotic have 3.

(End of review and extension of buffer calculations)

Dissociation constants for acids.

Zwitterions, Amino acids, and the Isoelectric point

Zwitterions - are ions containing both a positive and a negative charge simultaneously.

Amino acids are a particularly important class of amphiprotic compounds containing both weak acids (carboxylic acid) and weak bases (amines) (functional groups).

An example is glycine

three equilibria exist:

This is an example of a Zwitterion.

Isoelectric point is the point where the ion has no attraction in an electromagnetic field. This occurs when the pH of the solution creates equal ionization on the molecule.

At the isoelectric point [NH₂CH₂COO^- ] = [NH₃⁺CH₂COOH]

The relationship can be derived to predict this hydronium ion concentration and thus the pH:

example for glycine:

The composition of a solution of a Polyprotic Acid as a function of pH can be described.

= relative concentration of the free acid and free anion as a function of pH

c_T = sum of molar concentrations containing various species of the polyprotic acid.

₀ 

₁ 

₂ 

1 = ₀ + ₁ +  ₂

Solving each of the equilibrium for the species of interest and plugging them in, results in alpha values that are independent of the total concentration.

Solubility and Precipitation in Water

Ions are converted to an insoluble precipitate the precipitate is in equilibrium with the ions and a constant concentration is maintained at a given temperature and with constant ion concentration as long as there is a solid precipitate present. The precipitate molecule and ions are in a saturated solutions defined by the equilibrium

Example
K_sp - solubility product constant
usually written from the solid side of the equilibrium

Stoichiometry is the precise ratio of each of the elements (or complex ions) forming the molecule or solid precipitate.

1. Precipitation example assume - pH 7.0 in water

Anion: SO₄^-2

similar reactions for Ba, Ca, Pb, Sr, etc., Cations

Anion: CO₃^-2

similar reactions for Ba, Cd. Ca, Pb, Mg, Ag

Anion: OH^-

similar reactions for Al, Cd, Fe (II) & (III), Pb, Mg, Zn

Solubility and Precipitation

1. Precipitation example assume - pH 1.0 in water

What happens to the equations?

Why?

Remember LeChatelier Principle

Anion: SO₄^-2

SO₄^-2 is removed from solution which way will the reaction be driven?

Anion: CO₃^-2

CO₃^-2 is removed from solution which way will the reaction be driven?

For anion: OH^-

What happens to Al³⁺ ions and others like it?

 Similar reactions for Al³⁺, Cd²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Mg²⁺ , Zn²⁺

Precipitation Mechanisms

Three stages of precipitation

Nucleation - the formation of a small aggregate of molecules in a supersaturated solution

Crystal Growth - the addition of molecules or ions to nuclear aggregate particles.

Aggregate Particle Growth - the joining of large numbers of crystals to form large aggregate particles that settle out of solution.

Colloidal suspensions can form with particle sizes of 10^-6 to 10^-4 mm diameter. They are to small too filter using conventional filter material.

Colloidal particles are not retained by conventional filter paper and do not settle from solution but remain in suspension due to brownian movement.

Particles 10^-4 to 10^-1 mm tend to settle from solution as precipitates spontaneously.

Non-homogeneous composite solutions are refereed to as crystalline suspensions until they settle as precipitates.

Particle size

(not in text)

Particle size has been related to the relative supersaturation of the solution where:

High Relative Supersaturation tend to produce colloidal suspensions
(small crystals, colloidal suspensions - particle size < 10^-4 mm)

Low Relative Supersaturation produce more crystalline particles with greater diameter
(particle size 0.1 mm)

The rate of nucleation increases exponentially with RS while the rate of crystal growth increases linearly

Note:
All generalities assume that we are using WATER as the solvent!

Stability of colloid

(not in text)

Stability of colloids can be explained by associated charge and is related to the ionic strength of the solution.

Charged ions associated with each colloidal particle repel other particles tending to keep them small enough to be influenced by brownian movement.
(Skoog fig. 4-2 illustrates this phenomenon)

To optimizing precipitation conditions -
Minimize Q (concentration of solution) using more dilute solutions

S is Maximized when precipitation is from hot water solutions or with specific pH change

Increasing the total ionic content of a solution often coagulates the colloid and permits it to percipitate

Other Relevant Precipitation Events -

(not discussed in text)

Coprecipitation

Coprecipitation is the inclusion in the precipitate of the insoluble compound of normally soluble substances.

Precipitates tend to carry down other normally soluble components from the solution.

Adsorption is the process of association on the surface such as the ions about a colloidal particle. This is a surface effect and surface area influences this effect.

When particles are included into the lattice sites of the crystal the mechanism is referred to as Inclusion.

Physical entrapment of other material (not in the crystal lattice) is Occlusion.
(Skoog Figure 4-5 illustrates occlusion)

Adsorption is very important in natural water especially due to the amount of iron {Fe(OH)₃ and Fe(OH)₂}and aluminum hydroxide {Al(OH)₃}that are in the environment.

Iron and Aluminum hydroxides are flocculent (fluffy non-crystalline) precipitates that scavenge by adsorption most other metal ions.

Examples of Iron Hydroxide are used to illustrate the phenomenon.
(References Skoog and Dzombak and Morel)

The solubility of Al(OH)₃(s) as a function of pH. The amphoteric nature of Al(OH)₃ is clearly shown by its solubility in both highly acidic and highly basic solutions. Note that Al(OH)₃(s) is essentially insoluble over the pH range 4 to 10.

Solubilities of Fe(OH)₃(s) and Zn(OH)₂(s) as a function of pH. Note that a much lower pH is required to dissolve Fe(OH)₃(s) than to dissolve Zn(OH)₂(s). Therefore, at pH = 4.8, for example, Fe(OH)₃(s) precipitates and Zn²⁺(aq) remains in solution. The Fe(OH)₃(s) can be filtered off, thereby separating Fe³⁺(aq) from Zn²⁺(aq).

Solubility of ZnS and FeS in water saturated with H₂S(atm) at various pH values. We can separate Zn²⁺(aq) from Fe²⁺(aq) by saturating a solution buffered at about pH = 1.5 with H₂S(g). The ZnS precipitates and Fe²⁺(aq) remains in solution.

Solubility-product constants for various salts in water at 25 degrees C.

Solubility of silver benzoate in water as a function of pH. The addition of H₃O⁺(aq) shifts the solubility equilibrium to the right (increase in solubility) as H₃O⁺(aq) combines with C₆H₅COO^-(aq) to produce C₆H₅COOH(aq) and thereby decrease [C₆H₅COO^-].

Solubility

Solubility Products

EXAMPLE (Simple calculation and later Activity will be compared)

Molar concentration method -

Solubility of a slightly soluble salt in water.

How many grams of Ba(IO₃)₂(s) (fw = 487 g ) can be dissolved in 500 mL of water at 25 ýC ?

amount of Ba(IO₃)₂ = 7.32x10^-4 M x 0.5L = 3.66 x 10^-4 moles of Ba(IO₃) ₂

3.66 x 10^-4 moles of Ba(IO₃)2 x 487{[g Ba(IO₃)₂]/[mole Ba(IO₃)₂ ]} = 0.178 g

Activity

Molar concentration expresses the number of a particular species present per liter of solution but the effective concentration may be different.

WHY?

This difference is call the activity - arising form electrostatic interactions in solutions.

It is usually only significant for concentrated solutions ~ 1M.
This is why the assumption of [concentration, M] is appropriate below M < 10<sup-3

Cations and anions in solutions neutralize opposite charges and in concentrated solutions there is a greater probability that at any point in time more of them are interacting. In dilute solution the probability is decreased.

Activity - ax = f[X]

where
f - is called the activity coefficient
[X] - is the ionic concentration in solution

as [X] approaches lower concentrations, f approaches 1.

(in other words - as the solution gets more dilute f is less significant)

Therefore in dilute solutions activity - a - approaches ionic concentration [X] and ionic concentration is used to represent activity.

How significant can activity actually be? Compare at various concentrations.

Example Table 7-1 (Skoog) Activity Coefficient for Ions.

The activity coefficient f can be related to the charge on an ion and the ionic strength of the solution through the Debye Huckel relationship (1923)

Where:
f_i is the ionic strength, [&micro;] that the Debye Huckel relationship depends on can be calculated in the following manner:

The ionic strength [&micro;] that the Debye Huckel relationship depends on can be calculated in the following manner:

Where:
&micro; - is the ionic strength
C_i is the molar concentrations of ion i

 Z _i is the charge on the ion i
[[Sigma]] is the sum of (of each ion in solution)
[[alpha]] is the effective diameter of the hydrated ion i in Angstrom units

Also written: &micro; = 1/2([A]Z_A² + [B]Z_B² + [C]Z_C² +...)

How to apply activities in Chemical Equilibrium:

Equilibrium Constant - K

K is the thermodynamic equilibrium constant

{Because - Activity = a_x = f[X]; f is the activity coefficient}

K' is the concentration equilibrium constant

So:
Why do we use concentration and ignore activity 99% of the time?

* These relationships are important, know these relationships, 

these are conceptual and are what you should not forget so that 

the concept belongs to your understanding.

Activity coefficients for ions at 25 degrees C.

Activity coefficients for dissolved ionic species.

EXAMPLE:
Ionic Strength calculation

Calculate the Ionic Strength of (a) a 0.1M KNO₃ and (b) a 0.1M solution of Na₂SO₄.

using:

also

Where:
u - is the ionic strength
C_i is the molar concentrations of ion i

 Z _i is the charge on the ion i
A, B, C - represent ions

(a). the concentration of K⁺ and NO₃ ^- are 0.1 M

(b). the concentration of Na⁺ and SO₄ ^-2 are 0.1 M

There are 2 Na⁺ for each SO₄^-2 so that the Na⁺ conc. is 2 x 0.1 M or 0.2 M and
Zi = 2 for SO₄^-2 so Zi² = 2²

EXAMPLE:
Activity Coefficient calculation
(reference Skoog chapter 6 & 7)

Calculate the activity coefficient of Hg²⁺ in a solution that has an ionic strength of 0.085. Use 5 angstroms for the effective diameter of the ion. Compare the calculated value with f_Hg²⁺ in the table of activity coefficients for ions at 25 °C. (Skoog table 7-1).

from the Debye-Huckel equation

-Log f_Hg²⁺ = [(0.51)(2)² (0.085)^1/2]/[(1 + (0.33)(5)(0.085)^1/2]= 0.4016

f_Hg²⁺ = 10^-0.4016 = 0.397 = 0.40

From the Table of activity coefficients for ions at 25 °C. (7-1, Skoog)
We interpolate from the table the activities for f _Hg ²⁺ thus:

[(0.10 - 0.085)/(0.10 - 0.05)) x (0.46 - 0.38)] + 0.38 = 0.40

which agrees with our calculation for the Debye-Huckel equation

{Note: interpolation is a way to obtain a number between two other
numbers on a table once popular due to log table calculations}

Example of Solubility calculation using activity coefficients
(without common ion effect)

Reducing the concentration by an order of magnitude and the ionic strength with it.

Problem:
Using activities instead of concentration to calculate the solubility of Ba(IO₃)₂ in a 0.033 M solution containing MgCl₂.
The thermodynamic solubility product (K_sp), for Ba(IO₃)² (barium iodate) is 1.57 x 10^-9

(K'_sp is the concentration constant corrected for activity)

Replacing activities in this equation by activity coefficients (f) and concentrations [X] yields -(from Activity -a _x = f[X] );

K_sp = [Ba²⁺] f_Ba²⁺ x [IO₃^-]² f²_IO3 ^-

K_sp = [Ba²⁺] [IO₃^-] ² (f_Ba²⁺ X f²_IO3 ^-)

Rearranging

(K'_sp the concentration-based solubility product, corrected for activity)

K_sp divided by the activity coefficients is equal to the concentration based solubility product constant, K'_sp that is not constant but is dependent on the specific ionic strength of the solution.
Both (f_Ba²⁺ X f²_IO ^3-) are dependent on the ionic strength of this particular solution and the concentration of charged particles and their specific character.

Example of Solubility calculation using activity coefficients
(Cont. without common ion effect)
Using activities instead of concentration to calculate the solubility of Ba(IO₃)₂ in a 0.033 M solution containing MgCl₂

Ionic Strength due to MgCl₂:
(ion effect assumes that ionic concentration is from soluble salt and not from sparingly soluble precipitate)

We must first adjust the thermodynamic K_sp for the concentration and type of ionic strength and ion concentration of our solution and turn it into the K'_sp concentration specific for this unique ionic strength set of condition the molecule in question barium iodate is being calculated for.

Using the Table of Activity Coefficients (at 25 °C) for Ionic Strengths:
From the table in Skoog pg. 152 an ionic strength of 0.1

continuing: Correcting for activity effect in the constant
[Ba²⁺][IO₃^-]² = K _sp/( f_Ba2⁺ x f²_IO3 ^- ) = 1.57x10^-9/(0.38)(0.78)² = 6.8 x 10^-9 = K'_sp

K'_sp for Ba(IO₃)₂ Barium Iodate corrected for ionic activity in solution for this specific set of ions MgCl₂ Magnesium Chloride contributing to the ionic strength at their present concentrations of 0.033M in solution.

Considering ionic strength and activity:

Using the corrected solubility for activity and ionic strength of the MgCl₂ solution.

K'sp for Ba(IO₃)₂ Barium Iodate corrected for ionic activity in solution for this specific set of ions MgCl₂ Magnesium Chloride contributing to the ionic strength at their present concentration of 0.033 M in solution.

Correcting for activity effect in the constant
[Ba²⁺][IO₃]² = Ksp/ (f_Ba2 + x f2_IO3-) = 1.57 x 10^-9/(0.38)(0.78) ² = 6.8 x 10 ^-9 = K'sp

K'_sp = [Ba²⁺][IO₃^-] ² = 6.8 x 10^-9

[Ba²⁺] (2[Ba²⁺])² = 4[Ba ²⁺]³

[Ba²⁺][IO₃^-]² = [X] 2[X]² = 4[X]³ = 6.8 x 10^-9

X³ = 6.8 X 10 ^-9/4 = 1.7 X 10^-9

Soluble concentration of [Ba²⁺] = 1.2 x 10^-3 M, Considering activity
 

Example of Solubility calculation using activity coefficients

Using concentration instead of activity to calculate the solubility of Ba(IO₃)₂ in a 0.033 M solution containing MgCl₂. MgCl₂ is not considered in the calculation.

Neglecting ionic strength and activity:

K_sp = [Ba²⁺][IO₃^-] ² = 1.57 x 10^-9

[Ba²⁺] (2[Ba²⁺])² = 4[Ba ²⁺]³

[Ba²⁺][IO₃^-]² = [X] 2[X]² = 4[X]³ = 1.57 x 10^-9

4[X]³ = 1.57 X 10^-9 = K_sp

X³ = 1.57 X 10 ^-9/4 = 3.9 X 10^-10

Neglecting activity X = [Ba²⁺] = Ba(IO₃)₂ = 7.32 x 10^-4M

Soluble concentration of [Ba²⁺] = 7.32 x 10^-4 M, Neglecting activity

This is the same as on page 31 where activity was neglected also but as you can see it is not correct if there are other ions present in solution with discrete charges that will influence the dissociation of this molecule.

Compare the two calculations

Neglecting activity conc./Considering activity conc. x 100 = % difference

(7.32 x 10^-4)/(1.2 x 10^-3) x 100 = 61% difference in the two calculations

The Activity calculation is 39% different and 39% more accurate when Activity is also evaluated.

39% more barium iodate, Ba(IO₃)₂ is actually in solution than would be calculated using traditional calculations when these other ions are present in a very conservative example with low additional salt concentrations.

More of the (+) and (-) ions are associated with the other charges in solution and are thus not free to react to precipitate as would normally be predicted not considering other salts in solution. This is what makes environmental calculations very complex and requires the use of a more sophisticated approach.

Soluble ions like sodium chloride are electrostatically attracted to other ions and"...ions may be so close together that they behave as a neutral molecule..." and do not react as would be simply predicted by traditional equations.

Ba²⁺ (Cl^-) behave as neutral and Na⁺ (IO^3-) behave as neutral

Therefore the solubility

Ba(IO₃)_2(s) <---> Ba²⁺ + 2IO^3-

does not hold as a simple relationship in salt solutions.

( Ref. Hargis, Analytical Chem. Principles and Techniques; or Skoog ch.7, 6th ed.)

Titration of 20 mL of 0.100 M H₂A with 0.100 M NaOH. For H₂A, K₁ = 1.00 x 10 ^-3 and K₂ = 1.00 x 10 ^-7. Method of pH calculations is shown for several points and regions on the titration curve.

Review- Solving general equilibrium problems

Equilibrium

If you assume all H⁺ came from HA

Method 1. Successive Approximation

Method 2. Quadratic Formula

Use 5% rule to make assumption about denominator -x.

EXAMPLE
Conjugate base hydrolysis problem of the general form

Calculate the hydronium ion concentration and the pH in a 0.0100 M sodium hypochlorite (OCl^-) solution at 25 °C.
Given K_a = 3.0 x 10^-8.

pH
recall

Titration curve for 25.00 mL of 0.100 M maleic acid, H₂M, with 0.1000 M NaOH.

Curves for the titration of polyprotic acids. A 0.1000 M NaOH solution is used to titrate 25.00 mL of 0.1000 M H₃PO₄ (A), 0.1000 M oxalic acid (B), and 0.1000 M H₂SO₄

Composition of H₂M solutions as a function of pH.

Hydrated Metal Ion

(frequently form acid in nature)

Metal ions in water are usually hydrated by water to stabilize the charge. They can also be coordinated with other anions or proteins or organic molecules to stabilize them. We will study these effects individually.

Metal ions can act as acids
Hydrated metal ions can act as acids as they are hydrated

Example: Fe(III) in water
(Same example as in text)

This metal ion is a strong acid compared to organic acids

Metal ion acids increase in acidity with increasing positive charge and decreases with increasing ionic radius.

Many metal ions form polymeric species with OH^- as a bridging group. This is also another form of a complex formed with the hydroxide ions - 2(Fe(III) and 2(OH^-) for a net charge of 4+ .

This is a process known as dehydration-dimerization

Other 2⁺, 3⁺, 4⁺, 5⁺ ,6⁺ ions form polymeric species with OH^-

Al(III), Be(ii), Bi(III), Ce(IV), Co(III), Ga(III), Mo(V), Pb(II), Sc(II), Sn(IV) and U(VI)

Metal Ions Produce Acid Aqueous Environments

Metal ions can also shift equilibrium and produce acidic aqueous environments such as in the case of Fe. The following is the reaction that occurs in the situation known as acid mine drainage.

Why is this significant?

Hint:

Example:

When water at low pH from acid mine drainage, (say pH 2.5 to 3) flows into an aerated stream at a higher pH, (say pH 6 to 8.2) then the iron that was complexed as a hydrate precipitates out as a hydroxide. Such as "yellow breaches creek" or "Yellow Creek" in Indian County.

Calcium, Water "Hardness"

Calcium Deposits, and Carbonates related theme

Minerals containing Ca and Carbonate in contact with water

Water with high CO₂ levels dissolve higher quantities of Ca minerals from geological materials in contact with the water. (formation of caves)

When the equilibrium is reversed large quantities of CaCO_{3 (s)} are deposited thus giving rise to the term hardness.
(heating of water eliminating CO₂ deposits Calcium carbonate in pipes and boilers)

Species in water

Many different species of the same element can exist in solution.

Different species have different chemical reactivity.

Generally -- species are different forms of an element that are chemically distinguishable or that can react in a different chemical manner. These include different oxidation states, metal complexes, organometallic compounds, hydrates, etc.

Table 2.2 form text (Manahan) gives some general species significant in water systems. Think also about Fe(II), Fe(III), Cr(III), Cr(VI), these are examples of oxidation state species. Different ligand or complexed species exist as well as different organic and organometallic species.

We will discuss this topic more extensively as a separate topic also and expand on this theme.

See table of species as examples:

Complexation and Chelation

82% of the 108 known elements are metals.

30 elements are of key importance and are usually evaluated.

100,000 different organic molecules are produced by industrialized countries. Each may have several environmentally derived conversions into metal binding compounds.

Thus several million species formed as complex ligands with other elements, ions, and organic ligands may result form the combination of these two sources.

Competing reactions from precipitation, multiple complexation and ligand formation and molecular bond formation are all being applied simultaneously.

Solvent and transport medium -

All of the fundamental environmental chemistry is water related as water forms the universal solvent for the planet and is the transport medium as well. It also enables most species formation.

Complexation, Ligands, and Complexometric Reactions

Chelation is based on the formation of complex ligands

Metal Ions react with electron-pair donors to form:

Ligands are ions or molecules that form covalent bonds with cations by donating a pair of electrons which are then shared by the ligand and the cation.

The number of bonds formed per molecule is referred to a "Dentate" meaning tooth-like.

Unidentate reagents form a single bond per molecule.

Radioactive elements, and essential and toxic elements in the environment.

Occurance of organometallic compounds in the environment.

Results of an equilibrium computation for freshwater system containing four organic ligands

Some chelates are built up in a stepwise manner with successive additions of ligand bonds and successive stability constants that generally decrease in stability with the number of ligands.

Example - (form text pg. 51)

(For stepwise formation constants see Skoog Appendix 5 page A-11)

Rule of thumb - Take note. -
Complexes with monodentate ligands are usually less stable than those with multidentate ligands.

The formation of complexes occur when electron pair donors contribute electrons to a metal ion through the ligand bond filling empty orbitals of the metal ion.

Chelating Agents

(pronounced kee'late)

A chelate is a cyclic complex formed by a cation bonded by two or more donor groups contained in a single ligand.

One of the most important group of complexometric agents are the Aminopolycarboxylic and Polyprotic Acids

These are large organic compounds that posses amines and carboxylic acid groups.

They tend to form 1:1 complexes with metal ions and usually form 4 or 6 ligand bonds (tetra- or hexa- dentates).

Example of the most studied of these is EDTA or Ethylenediaminetetraacetic acid.

It is a Hexadentate ligand with Six potential bonding sites two amino groups and four carboxyl groups.

EDTA is water soluble and has four ionizable groups

It is a zwitterion of the form H₄Y

It is usually used in the salt form: Na₂H₂ Y&#128;H₂O

The Ionization of EDTA is pH dependent

The dissociation constants are:

EDTA combines in a 1:1 ratio with all metals regardless of the charge of the particular cation.

What about Na and H then (with 2 each) ?

Note: chelators can also be ion exchangers

It forms a stable complex with almost all metal cations (group or= II)

The steps in its progressive ionization are illustrated in
Figure 13-2, 3, 4. Fundamentals of Analytical Chemistry, Skoog, West, Holler, Saunders College Publishing, NY, NY, pg. 288-289, 1992.

Structure of a metal/EDTA chelate

Complexes of EDTA and Metal Ions

EDTA forma 1:1 complex with metal ions regardless of the charge on the cation.

The Formation Constants K_MY are given in Table 11-1 (13-1) for many metal ions.

 For

Figure 13-8 (Skoog) Depicts the pH range that EDTA functions over and the ions it complexes from pH 2 to 11.

Calculated titration of a model freshwater with cysteine (a) Speciation of the metals. (b) Cysteine speciation.

The hexadentate EDTA complex surrounds the ion and shields it from solvent interaction

Figure 13-4 (Skoog)

These laboratory complexes are pH dependent and must be carried out in Buffered solutions.

Natural complexes are also pH dependent and rely on naturally occurring buffers in natural water such as the carbonate or phosphate systems. (in biological systems such as Fe in hemoglobin)

Because the complexation is pH dependent a new term the Conditional or Effective Formation Constant (dependent on pH) apply at a single pH.

They modify

Structure of H₄Y and its dissociation products.

They take the following form

Where:

Therefore solving for the product of the two constants yields:

where:

K'_MY is the conditional formation constant at a specific pH where  ₄ applies.

₄ is calculated by applying the following relationship involving the pH through the hydronium concentration.

₄ = 

₄ has values of 3.7 x 10^-14 at pH 2.0 to 9.8 x 10^-1

as given in Table 13-2 (Skoog).

Composition of EDTA solutions as a function of pH.

Minimum pH needed for satisfactory titration of various cations with EDTA

Example calculations involving EDTA

1. Example of titration curve using Ca²⁺

General rule of thumb for the calculation of concentration pM ⁿ⁺ for a complexometric titration.

Figure 13-5 (Skoog) Is the titration curve from Example 1 (Ca).

Ca has a larger Formation Constant than Mg and this shows up in the titration curve just as the titration of a weaker acid would in an acid base titration.

This illustrates the selectivity of a complexing agent like EDTA for specific ions in solution.

The variability of the complex with different metal ions at different pH values results in a different formation constant at different pH values. A minimum pH needed for an acceptable complex of various metal ions is in the range of pH 6-8.

Besides competing for different ions other competing reactions must also be considered in environmental complexation reactions.

Competing reactions and the complexities of real environmental situations.

Multiple competing reactions are occurring.
Hydroxides of iron and aluminum start forming at a pH of 3.5.

Multiple complexing agents such as monodentate Cl^-, F^-, OH^-, NH₃, or H₂O

Chelating agents are found in natural waters, as part of industrial waste waters and sewage. They are ubiquitous in nature.

What are some of them?
Humic acids, Fulvic acids, Phosphates, anions, biological waste products (NH₃), Industrial molecules, etc, and water itself.

Example - Case study of EDTA at Oak Ridge National Laboratory -

Oak Ridge National Laboratory in Oak Ridge, TN used EDTA to chelate metals to make them more soluble in the clean-up and disposal process and pumped them into disposal pits and disposal wells at concentrations of 10^-7 M. Some 12-15 years after burial EDTA is making metals mobile in the environment that would not be mobile. This case study shows the slow degradation rate of EDTA and other similar chelators and the role they play in transport of metals in the environment. (The use of specific isotopes permitted the positive identification of these metal ions in the environment and identified them as anthropogenic and coming form the study)
{pg. 66 text, 6th ed. Env. Chem., by Manahan}

Others involving detergents (water softeners), Natural materials - later.

Environmental Examples of Applications of EDTA

A Common Water Hardness Testing

Calcium Carbonate equivalency is calculated by adding drops of EDTA solution to a measured Buffered solution of the water to be tested. The standard usually used is 1 drop per 0.065 g (1 grain) of CaCO₃.

Epilog (as an analytical tool)

Several decades ago the titration of common metal ions was performed using EDTA on a routine basis. Many of the complexometric procedures are 30 to 100 years old. In the past decades most of the tests have been transferred to spectroscopic methods.

Another form of complex formation -
Many metal ions form polymeric species with OH^- as a bridging group. This is also another form of a complex formed with the hydroxide ions - 2(Fe(III) and 2(OH^-) for a net charge of 4+.

A man made chelator in natural water (NTA)

Case Study -
The sodium salt of NTA is used as a detergent or detergent additive and is a substitute for phosphate detergents. It is present in natural waters as a contaminant in mg and greater quantities. It has a relatively long environmental lifetime in the environment being broken down by a dominantly bacterial mechanism almost exclusively.

The text goes through the ionization of Nitrilotriacetic acid (NTA).

NTA has 3 hydrogens to dissociate with K_as and pK _as as follows.

K_a2 dominates the species range for all normal natural waters due to its pH range. Figure 3.4 (Manahan) demonstrates this with the display of HT^2- as the predominant species from pH 3 to pH 10.

Example from the book of a series of complex interactions of -
K_sp, K_a3, K_f and K_w.

Hydroxides such as Pb(OH)₂ are solids and form with many metals at natural water pH's (usually at relatively low pHs).

Many metal ions are suspended in water or precipitated from water as hydroxides. An examination of the relationship of a chelate and a hydroxide.

Which are dominant at pH 8.0 in natural water?

Fractions of NTA species at selected pH values.

Example:

(Env. Chem. 6th ed. text, pg. 72-73)

Pb(OH)₂ is a species in suspension in a fresh natural water at pH 8.00 NTA is present and at pH 8 is predominantly HT^2- ion.

Pb(OH)₂(s) + HT^2- <---- PbT^- + OH^- + H₂O

Equations describing the solubilization are:

adding them algebraically

A question - which is the predominant species? - complexed lead as PbT^- or uncomplexed chelating agent NTA as HT^2- .

Thus  is 20/1 in favor of the complex over the free form of the chelating agent.

Also note that the hydroxide form of the solid precipitate was solubilized.

What does this tell you about the interaction of chelation in natural waters?

The solubility of this complex decreases with increasing pH

Apply Le Chatelier's principles to the overall reaction -

What would increasing [OH^-] do to the [PbT^- ]?

What would increasing [HT^2-] do to the [PbT^- ]?

What would decreasing [OH^-] do to the [PbT^- ]?

Example: What about for a Carbonate and NTA?

This is another very important ubiquitous component in natural waters.

A second Example with the chelator NTA.

Now lets do the same examination using lead carbonate (PbCO ₃ ).

PbCO₃ is in equilibrium with the CO₂, HCO ₃, CO₃2- system and the NAT species. At pH 7.0 the HT^2- species is again the dominant species.

K_eq =  = 4.45 x 10^-7 pK_a1=[-LogK_a ]= 6.35

K_eq =  = 4.69 x 10^-11 pK_a2=[-LogK_a ]= 10.33

Components of equilibrium (problem from text):

Given:
In natural waters HCO₃^- frequently has a concentration of 1.0 x 10-3 M/L

Thus:

Thus  is thus a 41 to 1 lead complexed as (PbT^-) to uncomplexed free NTA from of (HT^2-).

As can be seen from Le Chatelier's principle as applied to the overall equation

the higher the HCO₃^- concentration the less of the complexed lead (PbT^- ) and the lower the HCO₃^- the greater the complexed lead (PbT^- ).

Thus in this case the complex depends on the HCO₃ ^- concentration but both are dependent on the pH. Go back to the alpha table and adjust the pH and look at the effect of pH on complexation of NTA and see that the lower the pH the more the solubility and the higher the concentration of the complex PbT^-.

Competing Metal Ions in Complexation

Example using Ca , Pb and (NTA)

Since competing complexes are always present lets examine one of the most common ions, calcium, and its influence on the NTA complex of lead.

Calcium also forms the NTA complex from the dominant ionic form of NTA (HT^2- ). (given a pH of 7.0)

K' is the product of K_f for CaT^- and the dissociation of NTA K_a3 as before in the Pb example. Again in water Ca²⁺ is typically in natural fresh water at a concentration 1.00 x 10^-3 M (remember pH at 7.0).

Thus [CaT^-]/[HT^2-] is 76/1 thus a 76 to 1 calcium complexed as (CaT^- ) to uncomplexed free NTA from of (HT^2- ).

What about the competition the carbonate?
Combining equilibria:

Thus the distribution between the Pb complex and the Ca complex is described by

So twice the Ca chelate is present as Pb chelate. Thus the amount of Ca will influence the amount of PbCO₃ solubilized form the solid and we have seen that many other equilibrium effect the amount of an element mobilized. In summary influences from

Thus many thing effect the complexation and equilibrium in water.

One final example of competing complexes, solubility's, metal chelates, pH and other equilibrium -

Graphical representation of the developed example -

This is one reason why we will need mathematical models to assist in combining all these competing equilibrium simultaneously.

Phosphate and Plyphosphates in Water

Orthophosphate ion PO₄^3-

Phosphoric acid H₃PO₄

Phosphate is one of the essential molecules of living things and of industry and agriculture.

Orthophosphate (tetrahedral) (a resonance structure with one double bond between phosphorous and oxygen in resonance),
(McQurrie and Rock pg. 398)

Holds Calcium in solution by chelating it and solubilizing it from CaCO₃ geological formations.

It is also a buffer with a pK_a at 7.31 and adds to the total alkalinity or total buffer capacity of a natural water system.

Phosphate in many forms is a complexing agent and a buffer.

Different forms have very different K_a1 , K_a2 ... K_an for polyphosphates

Phosphate can exist in straight chain polymeric forms of commonly 4 to 18 Phosphate monomers such as the:

Pyrophosphate ion P₂O₇^4- or Pyrophosphoric acid H₂P₂O₇
Triphosphate ion P₃O₁₀^5- or Triphosphoric acid H₅P₃O₁₀

Phosphate also forms ringed polyphosphates such as-
Trimetaphosporic acid and Tetrameaphosporic acid

These acids have different acid character from the monomer of phosphoric acid.

In natural waters all polymeric phosphates act as if they have two distinct hydrogen types that have specific pK_a values of
pK_a1 at pH 4.5 and
pK_a2 at pH 9.5.

at pH 4.5 all the linear-chain hydrogens are lost (3H⁺ , hydronium ions)

at pH 9.5 all the end-chain hydrogens are lost (2H⁺, hydronium ions)

All of these polymers and ring structures eventually hydrolyze in natural waters to orthophosphoric acid.

Phosphoric acid linear- chain polyphosphates and ring phosphates along with its various salts are complexing agents along with SO₄^2-, NH₃, Cl^- , CN^-, EDTA, NTA, F^-, OH^-, CH₃COO^- and other naturally occurring molecules such as humic acids.

All function simultaneously in natural water systems.

Complexation by Humic Acids and Other Organic Matter.

One of the most important types of complexing agents that are naturally occurring are the humic acid and fulvic acids which are a class of compounds of similar origin and similar structural design.

These acids are not single molecules but a Class of Molecules.

Fulvic acids is generally accepted to be the soluble form of a large complex molecule class of 500 to 2,000 molecular weight.

Humic acids are generally accepted to be less soluble and exist in solid or particulate form in water and soil.

They have been recognized since 1800 but have only been recently studied extensively (since ~ 1970), due to there role in the formation of trihalomethanes {THMs} during chlorinating of water in purification.

Both forms ion exchange and complex metal ions accumulating and releasing these metals in complex interactions with pH, ionic strength, temperature, specific ionic content and other interactions.

Routine composition overviews suggest compositions of -

45-55% C; 30-45% O; 3-6% H; 1-5% N; and 0-1% S

Examples of these ions show their very complex structure.

You can recognize the usual chelating functional groups in abundance.

A possible schematic pathway for the formation of terrigenous humic acid from lignins.

Three proposed structures for terrigenous humic acids.

A possible pathway for the formation of marine humic acids.

Complexing functionalities of some biogenic ligands. (a) hydroxamate siderophores; (b) catechol siderophores; (c) phytochelatins.

The binding characteristics of the fulvic and humic acids follow roughly this pattern:

Charge, ionic radius, pH and other factors effect the complexation of these chelating agents.

Lignite is high in humic acids and is an example of a largely nonsoluble humic acid type. It is known for ion exchange of metals from percolating water.

In the environment all of the effects discussed so far and several more to come are acting simultaneously and therefore the most complex interactions known are what we are trying to study.

The complexity of natural phenomenon - a perspective.

In living systems (a group of cells, for example) certain patterns are set up, maintained by energy input by the system, ordered and others prohibited by the system (while it is alive) and patterns are discovered and identified. In each living system we looked for and identified these patterns in a particular order and establish the norm for this healthy living system.

In nature patterns are governed by the interaction of all equilibrium simultaneously and the results are up to us to predict and understand. The health of the system has many more natural patterns and is more difficult to study as it has more variables. If you are looking for a challenge you have found it.

References

References used in preparation of these lectures.
These are additional references of interest and for future inquiry.

Aquatic Chemistry (An Introduction Emphasizing Chemical Equilibria inNatural waters), 2nd ed., Warner Strumm and James Morgan, John Wiley & Sons, Inc., NY, NY, 1981.

Environmental Inorganic Chemistry, Eds Kurt Irgolic and Arthur Martell, VCH Publishers, Deerfield Beach, FL, pgs. 1, 2, 4, 1985.

Principles and Applications of Aquatic Chemistry, Francois Morel and Janet Hering, John Wiley & Sons, Inc., NY, NY, pgs. 3, 369, 377-379, 394, 1993.

Surface Complexation Modeling - Hydrous Ferric Oxide, David Dzombak and Francois Morel, John Wiley & Sons, Inc., NY, NY, pgs. 104-105, 192, 1990.

Environmental Chemistry, Nigel Bruce, Wuerz Publishing Ltd. Winnipeg, Canada, 1991.

General Chemistry, Donald McZuarrie and Peter Rock, 3rd ed. Chapter, 20 (Solubiligy & Precipitation Reactions), Intrerchapter I & J, W. H. Freeman and Company NY, 1987.

Fundamentals of Analytical Chemistry, Skoog, West, Holler, Saunders College Publishing, NY, NY, pg. 288-289, 1992.
 

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Revised 4/23/99.