1. Some Basic Terminology
Solute: chemical species that is dissolved in water; element or compound, charged or neutral
Total Dissolved Solids (TDS): total amount of solids remaining when a water sample is evaporated to dryness (mg/L)
Related to salinity
Fresh water: < 1000 mg/L
Drinking water standard = 500 mg/L
Brackish water: 1000 – 20,000 mg/L
Saline water: > 20,000 mg/L
Seawater = 35,000 mg/L (on average)
Brines: > 100,000 mg/L

2. Ocean Salinity

3. Solute Concentration Terminology
(mass) concentration: mass of solute / unit volume of solution (e.g. mg/L)
Molarity (M): (mass of solute/formula weight) / liters of solution = moles/L
Molality (m): moles of solute / kg of solution
Equivalents/L: (moles of solute x valence of solute) / L of solution = eq/L
Used for ion balances, piper diagrams
Normality (N): eq/L (primarily used for acids)
Parts per million (ppm) by weight: 1 millionth g solute per g of solution = mass of solute / 106 g water
commonly used outside academia
for most dilute water solutions, mg/L ≈ ppm and molarity ≈ molality

4. Rocks and Minerals

5. Mineralogy of Igneous Rocks:
Bowen’s Reaction Series
At/Near Earth’s
Surface:
Less Stable
More Stable

6. Mineralogy of Igneous Rocks
Mafic
Olivines:
fayalite (Fe2SiO4)
forsterite (Mg2SiO4)
Pyroxenes:
wollastonite (CaSiO3)
enstatite (MgSiO3)
augite ((Ca,Na)(Mg,Fe,Al)
Plagioclase feldspars:
anorthite (CaAl2Si2O8)
Amphiboles:
hornblende ((Na,K)0-1Ca2(Mg,Fe,Al)5Si6-7.5AlO(OH)2

7. Mineralogy of Igneous Rocks
Felsic
Feldspars:
Albite (NaAlSi3O8)
orthoclase (KAlSi3O8)
anorthite (CaAl2Si2O8)
Micas:
muscovite (KAl2[Si3AlO10](OH)2
biotite (K(Mg,Fe) 3AlSi3O10(OH)2)
Quartz: SiO2

8. Mineralogy of Metamorphic Rocks
Mineral composition reflects parent rocks
e.g. marble from limestone (CaCO3)

9. Mineralogy of Sedimentary Rocks
Shale: low reactivity
Quartz, clay/mica, feldspars, calcite, organic matter
Sandstone
Quartz, K/Na feldspars, micas, many other minerals
Carbonates
Calcite: CaCO3
Dolomite: (Ca,Mg)(CO3)2
Chert (SiO2)
Clastic sediments
Evaporites
Gypsum: CaSO4·2H2O

10. Mineralogy of Secondary Minerals
Form by chemical reactions (weathering/ diagenesis) between water and primary/ secondary minerals
Incongruent dissolution (solid products)
Stable, may be in equilibrium with H2O
Clays:
mineral form depends on reaction type and climate (precipitation, temperature)
garbage minerals: chemical composition highly variable, not well known

11. Mineralogy of Clays Mafics (high Mg, Ca, Fe) Felsic (high Na, K, Si)
Chlorite (Mg,Fe)3(Si,Al)4O10(OH)2·(Mg,Fe)3(OH)6  vermiculite  smectites (Na,Ca0.5)0.7(Mg,Fe,Al)4(Al,Si)8O20(OH)4  kaolinite Al2Si2O5(OH)4  Fe/Al oxides [Al(OH)3, Fe(OH)3]
 decreasing cation content
Felsic (high Na, K, Si)
Illite K15(Mg0.5Al3.5)(AlSi7)O20(OH)4 (close to muscovite)  smectites  kaolinte  Fe/Al oxides
 degree of weathering
80% muscovite, 20% smectite main clay in shales
These are general trends observed. Often, several clay minerals exist, but dominant type can be predicted based on rock type and amount of precipitation
Illite, kaolinite, chlorite, vermiculite, and montmorillonite common in Illinois

12. Silica Minerals
Most Si released by weathering of primary silicates (feldspars, etc.)
Quartz (SiO2): not dominant secondary mineral, very resistant to weathering
Secondary Si minerals (SiO2)
Amorphous silica: non-crystalline
Chalcedony (agate): cryptocrystalline
Chert: nodules/beds
Opal: gem quality
Secondary SiO2 is found in all rock types

13. Oxyhydroxides of Fe/Al
Also found in all rock types
Al: very low solubility (more soluble at high and low pH)
Amorphous Al(OH)3
Gibbsite: Al(OH) 3
Fe: Ferric iron (Fe3+), low solubility
Amorphous Fe(OH) 3
Goethite: FeOOH
Hematite: Fe2O3
Very common in Illinois sediments

14. Carbonates and Sulfates
Found in almost all geologic environments
Formed by 2 most common dissolved ions
Calcite: CaCO3
Also dolomite [(Ca,Mg)(CO3)2], siderite [FeCO3]
Sulfates
Gypsum: CaSO4·2H2O; layers or disseminated
Anhydrite: CaSO4

15. Other sources of material to the subsurface
Dust
Organic matter

16. Dissolved Molecules
Dissolved is (somewhat) arbitrarily defined as anything that passes through a 0.45 μm filter
> 0.45 μm defined as “suspended” material
Some colloids and all nano-particles are < 0.45 μm
Aqueous Species refers to any molecule dissolved in water
Ions are charged molecules only
HCO3- is an ion, H2CO3 is not (uncharged)

17. Dissolved Molecules Major ions defined as being typically > 5 mg/L
Cations: Ca2+, Mg2+, Na+
Anions: HCO3-, SO42-, Cl-
H4SiO4: almost all dissolved Si in this form
In Groundwater, these typically account for > 90% of TDS regardless of TDS value

18. Dissolved Molecules
Minor constituents typically between 0.01 and 10 mg/L, higher in unusual situations
K+, Fe2+, Mn2+, NO3-, F-, B(OH)3, CO32-
Trace constituents usually < 0.1 mg/L
Most metals
Rare earth elements
Arsenic, bromide, iodide, phosphate, radium, uranium

19. Balancing Chemical Reactions
Reactants Products

20. Balancing Reactions
Determine the species involved (reactants and products) and the state (solid, liquid, gas) that each species is in
Write an unbalanced equation that summarizes information from step 1
Balance: molar amounts on each side of reaction must be equal
Stoichiometry: refers to the mathematical balancing of reactants and products (moles)

21. Balancing Reactions
Balancing (determining stoichiometric coefficients)
Start with most complicated species
Cations first, then Si, then C, then O, then H
Check charge balance
Assume that H2O, H+, and OH- are readily available since it’s an aqueous reaction
In most weathering reactions, acid H+ attacks the mineral
Examples…

22. Reactions and Equilibrium

23. Chemical Reactions Chemical reactions usually need water
When 2 ions or molecules approach each other closely and establish a bond, a chemical reaction has taken place
e.g., precipitating a solid
Breaking bonds is also a chemical reaction
e.g., dissolving a solid

24. Reactions and Equilibria
Let’s add salt (NaCl) to water at constant T, P; add more than can be dissolved
Salt crystals will dissolve, Na-Cl bonds are broken, ions (Na+, Cl-) in solution
What if we measured the Na+ (or Cl-) concentration in solution over time?

30. Reactions and Equilibria (NaCl)
At equilibrium, Na+ and Cl- concentrations in solution become invariant with time
This does not mean that all reactions have stopped; both NaCl dissolution and precipitation are still occurring
At equilibrium, the rate of NaCl dissolution = rate of NaCl precipitation
NaCl(s)  Na+ + Cl- and Na+ + Cl-  NaCl(s) have the same rates
NaCl(s) Na+ + Cl-

31. Equilibrium
The system maintains a constant state, and there are no observable changes in the constituents in the solution
In Nature, reactions tend to progress to equilibrium
That is the preferred state
Disequilibrium is inherently unstable
But, there are plenty of reactions that are in disequilibrium near the Earth’s surface

32. La Châtelier’s principle
When a reaction at equilibrium is disturbed, i.e., if conditions change, a new state of equilibrium will be established to counteract the disturbance
e.g., heat up NaCl solution, more NaCl will dissolve
Or, we add more water more NaCl will dissolve

33. can’t dissolve any more
Saturated solution:
can’t dissolve any more

34. Solubility
The amount of a compound dissolved to form a saturated solution
In our example, the amount of NaCl that dissolved
Measured in weight per volume (e.g., g/L)
Need to state T, P (usually 25°C, 1 atm, which is standard state)

35. Solubility Need equilibrium to define solubility
e.g., CaCO3 + 2HCl  Ca2+ + 2Cl- + 2H2O + CO2
In a closed system, reaction will (eventually) reach equilibrium
In an open system, this reaction cannot reach equilibrium because CO2 does not accumulate
Reaction will proceed until CaCO3 or HCl is used up, i.e., reaction runs to completion
Many reactions at/near Earth’s surface do not achieve equilibrium because products escape or are transported away
Also many reactions occur slowly

36. Law of Mass Action
Mathematical model that explains and predicts behaviors of solutions in equilibrium
Valid only for reversible reactions
Consider the reaction: A + B  C + D
rate of forward reaction (uf) = kf (A)(B)
kf = proportionality constant for forward reaction
(A) and (B) are in molar amounts
Rate a function of reactant concentrations
rate of backward reaction (ub) = kb (C)(D)

37. Law of Mass Action A + B  C + D At equilibrium, uf = ub =
The ratio of products to reactants is fixed
So kf (A)(B) = kb (C)(D) =
= Keq = the equilibrium constant
Keq values have been determined in the lab for many reactions
kf (C)(D)
kb (A)(B) kf kb

38. Law of Mass Action For reaction: aA + bB  cC + dD
Where a, b, c, and d are the stoichiometric coefficients for the reactants and products
Keq = [(C)c(D)d]
[(A)a(B)b]
(A), (B), (C), (D) are concentrations at equilibrium
[(C)c(D)d] is the reaction quotient (Q)
Q changes until equilibrium is achieved, then it is constant
However, strictly speaking this is only valid for ideal gases

39. Solubility Index Calculations
We can use the mass action equation to estimate the equilibrium status of a particular reaction.
Is the solution undersaturated, oversaturated, or at equilibrium with respect to a mineral?
i.e., do we expect it to dissolve or be precipitated?

40. Solubility Index Calculations
SI = log IAP – log Keq
SI = saturation index
IAP = ion activity product = measured concentration (= Q)
Keq = equilibrium constant (IAP at equilibrium)
If log IAP = log Keq, then SI = 0 (equilibrium): no net dissolution or precipitation of mineral
If log IAP < log Keq, SI < 0, solution is undersaturated with respect to that mineral: expect active dissolution
If log IAP > log Keq, SI > 0, solution is oversaturated with respect to that mineral: expect active precipitation
In practice, if SI = 0 ± 0.5, then water at or close to equilibrium with mineral phase, due to uncertainty in the thermodynamic variables