Fundamentals
of Aquatic Chemistry, Chemical Equilibrium &
Chemical Solubility
Reading Assignment: Read the material in this lecture and review acid base and equilibrium calculations, and finish reading chapter 3 in Manahan.
Homework: HW-3, Due Monday, February 4.
Links
and Additional Resources:
Three things to learn from this section 1. How chemical equilibrium applies to natural systems 2. Carbonate equilibrium of Natural systems 3. Other Equilibrium that are important in Environmental Chemistry Define Acid, Base, Salt Strong Acid, Strong Base (100% ionized in water) Examples Weak Acids, Weak Bases (partially ionized in water) Examples The most important weak acid in nature is carbonic acid H2CO3
Acid Base
review at University of Akron
Acid
Base review at University of North Carolina Charolette
All gases are also in a equilibrium between air and water
Mathematical relationship known as the "Mass Action Equation" is obeyed for all equilibrium systems.
Variables change their values, depending on conditions
Constants have fixed values
This is the challenge:
Use the constants, mass action expression, and "measured or specified concentrations" to calculate concentrations for unknown species.
Calculate HCO3- , CO32- and H2CO3 when the pH is known.
Why is this important?
Carbonate equilibrium in Natural Water Systems:
When calcium is added to an aqueous solution of carbonate, we form calcium carbonate--a sparingly soluble salt.
Lets quickly review solubility rules.
An old saying is "Like dissolves like". In practice, this means that substances that dissolve in water would be expected to have a molecular structure like water. We find that substances that dissolve in water are "ionic". As a general rule, a substance that is not Ionic will not dissolve in water and a substance that is ionic does dissolve in water.
Some Examples:
Oil (a saturated hydrocarbon CH3-CH2-CH2-CH2-CH2-CH2-CH2-CH3)
does not dissolve in water.
Plastic (Polystyrene for example) does not dissolve in water.
Sodium chloride (NaCl) does dissolve in water by forming individual ions of sodium (Na+) and chloride (Cl-) that interact with the dipole of water molecules.
and 
Chemists learned long ago that it was possible to modify the structure of some molecules to make them soluble in water. This is accomplished by making part of the molecule ionic (or at least giving it a dipole moment).
Examples of organic molecules that dissolve in water are:
Acetic Acid or Vinegar
Ethyl Alcohol
Ethylene Glycol or Anti-Freeze
Compounds that are "sparingly" soluble in water are in equilibrium between an ionic form and a non-ionic form. Calcium carbonate is an example of such a substance
The equilibrium expression that governs this reaction is called the solubility product, or Ksp. Ksp is a constant value for any substance, and if known, can be used to calculate the solubility of that substance in water.
For Calcium Carbonate:
The solubility of calcium carbonate in pure
water (in a system not exposed to air and ignoring reactions with water)
can be calculated as follows:
Let X be the concentration of Ca2+
in solution--at equilibrium since the number of moles of Ca2+
is equal to the number of moles of CO32-, X is
also equal to the number of moles of CO32-.
Then:
Therefore, the solubility of calcium carbonate in water--if not open to the air, is
6.69 x 10-5 moles CaCO3/Liter.
By allowing the water to mix with air, the value for
CO32- would change to reflect CO32-
from dissolved calcium carbonate and from CO32-
from dissolved carbon dioxide. Calculating the solubility of calcium carbonate
for a system open to the air is more difficult than the simple example
we just did. Calcium carbonate in water, open to the air, is an important
system to understand because it represents the weathering process when
limestone (CaCO3) encounters fresh water. We will deal with
this in a future lecture.
Solubility versus Intrinsic Solubility
The simple solubility calculation we just completed leads to an error of as much as 100% if salt water is used instead of pure water. Since most natural waters contain at least a few dissolved salts, we should try to understand what causes this error.
The errors encountered for solubility in salt solutions results from an effect called "activity". Salts modify the properties of water by reducing an ion's activity in solution. This reduced activity results in greater than expected solubility.
Activity is related to ionic strength, and these effects can by quantified if an ion's activity coefficient is known. Activity coefficients for some important ions are shown in the following table.
| Table 8.1. Activity coefficients for aqueous solutions at 25 | ||||||
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| H+ | |
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| Na+, Cd+2, Cl-, ClO2-, IO3-, HCO3-, H2PO4- HSO3-, H2AsO4-, CO(NH3)4 (NO2)2+ | 450 | 0.964 | 0.928 | 0.902 | 0.82 | 0.775 |
| OH-, F-, SCN-, OCN-, HS-, ClO3-, ClO4-, BrO3-, IO-4, MnO4- | 400 | 0.964 | 0.927 | 0.901 | 0.815 | 0.770 |
| K+, Cl-, Br-, I-, CN-, NO2-, NO3- | |
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| Rb+, Cs+, NH4+, , Ti+, Ag+ | |
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| Mg2+, Be2+ | |
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| Ca2+, Cu2+, Zn2+, Sn2+, Mn2+, Fe2+, Ni2+, Co2+ | 600 | 0.870 | 0.749 | 0.675 | 0.485 | 0.405 |
| Sr2+, Ba2+, Cd2+, Hg2+, S2-, S2O4- | 500 | 0.868 | 0.744 | 0.670 | 0.465 | 0.380 |
| Pb2+, CO32+, MoO42- , Co(NH3)5NO2- | 450 | 0.867 | 0.742 | 0.665 | 0.455 | 0.370 |
| Hg22+, SO42-, S2O32-, S2O82-, SeO42-, CrO42-HPO42- | |
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| Al3+, Fe3+, Cr3+, Sc3+, Y3+, In3+, lanthanides |
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| PO43-, Fe(CN)63-, Cr(NH3)63+, Co(NH3)5H2O3+ | |
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| Th4+, Zr4+, Ce4+, Sn4+ | |
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The activity coefficients from the Table are now used
to calculate an "intrinsic" solubility that better approximates
the concentration for a substance in the real world. For Example, calculate
the intrinsic solubility for calcium carbonate in a closed system where
the ionic strength is 0.10 (close to sea water)
Activity = Concentration x Activity Coefficient
The intrinsic solubility of 0.10 ionic strength,
is 1.73 x 10-4 -- more than twice the solubility in pure
water.
ENV 440 - Course Topics