Lecture 1-13-11

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
Ocean Salinity
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 ≈
Rocks and Minerals
Mineralogy of Igneous Rocks:
Bowen’s Reaction Series
At/Near Earth’s
Less Stable
More Stable
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
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
Mineralogy of Metamorphic Rocks
• Mineral composition reflects parent rocks
– e.g. marble from limestone (CaCO3)
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
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
Mineralogy of Clays
• Mafics (high Mg, Ca, Fe)
– 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
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
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
Carbonates and Sulfates
• Carbonates
– 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
Other sources of material to the
• Dust
• Organic matter
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)
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
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,
Balancing Chemical Reactions
Balancing Reactions
1. Determine the species involved (reactants
and products) and the state (solid, liquid, gas)
that each species is in
2. Write an unbalanced equation that
summarizes information from step 1
3. Balance: molar amounts on each side of
reaction must be equal
– Stoichiometry: refers to the mathematical
balancing of reactants and products (moles)
Balancing Reactions
• Balancing (determining stoichiometric
– 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
• Examples…
Reactions and Equilibrium
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
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?
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
• At equilibrium, the rate of NaCl dissolution = rate of
NaCl precipitation
– NaCl(s)  Na+ + Cl- and Na+ + Cl-  NaCl(s) have the same
– NaCl(s)
Na+ + Cl30
• 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
– That is the preferred state
– Disequilibrium is inherently unstable
– But, there are plenty of reactions that are in
disequilibrium near the Earth’s surface
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
Saturated solution:
can’t dissolve any more
• 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)
• Need equilibrium to define solubility
• e.g., CaCO3 + 2HCl  Ca2+ + 2Cl- + 2H2O + CO2
– In a closed system, reaction will (eventually) reach
– 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
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)
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)
– kf = Keq = the equilibrium constant
– Keq values have been determined in the
lab for many reactions
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), (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
Solubility Index Calculations
• We can use the mass action equation to
estimate the equilibrium status of a particular
– Is the solution undersaturated, oversaturated, or
at equilibrium with respect to a mineral?
– i.e., do we expect it to dissolve or be precipitated?
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

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