Redox, ph, pe Equilibrium? OUTLINE Finish last lecture Mineral stability Aquatic chemistry oxidation and reduction: redox Reading: White p555-563 1
Question of the day? So what about the CO 2 system? CO 2 + conjugate bases set ph for many natural waters ( CO 2 (aq), H 2 CO 3 (aq), HCO 3- (aq), CO 3 2- (aq) ), other acids/bases modify this So CO2 in H2O buffers: H 2 CO 3 <> H + + HCO 3 - K a1 = [H + ][HCO 3- ]/[H 2 CO 3 ] pk a1 = ph - log [HCO 3- ] [H 2 CO 3 ] HCO 3- <> H + + CO 3 2- K a2 = [H + ][CO 3 2- ]/[HCO 3- ] pk a2 = ph - log [CO 3 2- ] [HCO 3- ] 2
Alkalinity acid-neutralizing capacity (removing H + ) sum of all the titratable bases In nature mostly: CO 2 (aq) and OH -. [Alk] [HCO 3- ] + 2[CO 3 2- ] + [OH - ] - [H + ] ( + borates, phosphates, organic acid anions, silicates or other bases) Behavior of ions in water: Elemental distribution in water = f (interaction with H2O, OH -, H + and dissolved oxygen aka DOx ) => function of O bonding Electronegativity (desire to gain e - ) and ion size determine the bonding preference of a cation for DOx or water 3
Solubility Solubility of O-complexes depends on stability in H2O: Charge (electronegativity) and (ionic) size Elements in green are largely insoluble (med IP: next) Behavior of ions in water - Ionic Potential = charge Z / radius r (=effort to remove e-) Ionic bonds with O The intermediate ones don t dissolve Covalent bonds with O Caveat: compounds can modify this, e.g. dissolved organic carbon 4
9/12/17 Solubility in pure water, elements on the outer edges are more soluble than those in the middle Concentrations in units of molarity (mole/l), molality (mole/kg), ppm by weight (or mg/kg = μg/g) Saturation = maximum solute concentration in solution Solubility versus particles dissolved organic matter key: affects solubility of heavy metals and their mobility: When bound, part of dissolved phase. Particles also affect solubility: ions can attach to solid particles => part of the suspended phase 5
Complexes/chelates and solubility Chelates: complex with multi-dentate ligands (>1 electron to donate to a cation) Many organic molecules chelate metals ÞLeads to mobility of metals Humic Substances are naturally occurring organic chelates Chelates vs simple complexes There s a preference for multidentate chelates Example (Ni, or M) H 3 N---Ni---NH 3 vs. Ni-N very similar bond energy (ΔH ~ 0) ethylenediamine ΔS formation 2 NH 3 ligands and 1 metal ion requires more order than 1 ethylene diamine ligand and 1 metal ion ΔS reaction =+ for chelate formation: ΔG = -TΔS, ΔG = -ΔS -- favors chelate 6
TDS, or total dissolved solids The dry weight of all solutes in solution per liter or kg of solution. TDS includes ionic and covalent solutes. high TDS = lots of things in solution. low TDS ~ pure solvent. TDS affects many properties of an aquous solution density Pure water has a density of 1 kg/l at 4 o C Sea water (mean density of 1.034 kg/l) can be thought of as ~1 kg/l water and 0.034 kg/l TDS, or ~34 g TDS/L Solubility Specific solutes can be more or less soluble as in a natural water as a function of TDS Usability High TDS waters tend to be less useful for urban and industrial settings because precipitates can foul machinery and pipes. TDS related to physical and chemical processes - precipitation and evaporation - weathering (dissolution/precipitation, leaching, and ionexchange) - temperature - ph - gas solubility - biological processes 7
TDS in values for different natural waters Water "type" TDS (mg/l) Examples Fresh <1000 rain, river, lakes, drinking water Brackish 1000-10000 estuaries, lagoons, near-shore aquifers, some inland seas Saline 10000-100000 oceans, some inland seas, some geothermal waters Brine >100000 shallow tidal basins, geothermal waters Ways to dissolve based on bond-type 1. ionically bonded solids, which dissociate NaCl (s) Na + (aq) + Cl - (aq) 2. covalently bonded material (glucose) dissolve unchanged C 6 H 12 O 2 (s) C 6 H 12 O 2 (aq) 3. reactive, incongruent dissolution: leaves new phase CO 2 (g) CO 2 (aq) HCO 3- (aq) + H + (aq) (to CaCO3) MgSiO 3 (s) + H 2 O Mg 2+ (aq) + SiO 2 (aq) + 2 OH - (aq) Pyroxene and H2O reacting to dissolved Silica, Mg2+ and OH- ions 2KAlSi 3 O 8 +2H + + H 2 O Al 2 Si 2 O 5 (OH) 4 + 2K + + 4SiO 2 (aq) K-feldspar & H2O reacting to Kaolinite (clay), dissolved silica, K+ ions 8
Relationship between saturation point and K NaCl (s) Na + (aq) + Cl - (aq) 1 mole NaCl (s) -> 1 mole of Na + (aq), 1 mole of Cl - (aq). solubility = moles of NaCl (s) dissolved into x volume at saturation [NaCl(s)] = [Na + ] = [Cl - ] With K sp = [Na + ][Cl - ] Solubility x =[Na + ]=[Cl - ], and using K sp = x 2 => x = K sp ½ CaF 2 (s) Ca 2+ (aq) + 2F - (aq) Solubility = x = [Ca 2+ ] = ½ [F - ] K sp = [Ca 2+ ][F - ] 2 => K sp = x (2x) 2 = 4x 3 => x = (K sp /4) 1/3 Write out similarly for covalent bonds and reactive dissolution The common ion effect Tends to lower the expected solubility of a salt relative to that in pure water. Example.. Some equilibrium of PbI 2 ó Pb 2+ + 2I - Add (K)I drives reaction to PbI 2 So solubility depends on each other, due to the common I - ion. 9
Mineral Stability diagrams Predicting mineral stability, e.g. incongruent dissolution of K-feldspar into water to Kaolinite: 2KAlSi 3 O 8 +2H + +9H 2 O Al 2 Si 2 O 5 (OH) 4 +2K + +4H 4 SiO 4 (aq) fsp acidic water kaol dissolved K, SiO 2 K eq = [H 4 SiO 4 ] 4 [K + ] 2 take log logk eq = 4 log[h 4 SiO 4 ] + 2 log [K+] [H + ] 2 [H + ] Or: logk eq / 2 2 log [H 4 SiO 4 ] = log [K+] / [H+] b -2 * x = y or: y = -2*x + b Mineral Stability Diagrams logk eq / 2 2 log [H 4 SiO 4 ] = log [K+] / [H+] b -2 * x = y Line (slope = -2) separates fsp from kaol, describes needed chemical condition for fsp weathering 10
Adding silica Amorphous silica (s) deposits at higher concentrations (log(h4sio4)=0 to - 2.74): adds vertical line for new phase assemblages SiO 2 (s, amorph) + 2H 2 O H 4 SiO 4 (aq) K eq = a H4SiO4 = 10-2.74 at 25 C Amorph. Silica deposition depends on kinetics (metastable) fluids More minerals and examples Equilibria for K-mica, gibbsite and quartz, and example water compositions K+ also affected by sorption onto charged particles, modifies [K+]/[H+] Can also do this for other elements (e.g. Mg) to consider other minerals 11
1. Oxidation- Reduction: Together with ph, redox controls composition of natural H2O Measure of charge on an atom in any chemical form, think of balance between e- and p+ For monatomic species, it is equal to the ionic charge (e.g., Fe +3 ox. state = +3, Cl - ox. state = -1). 1. Oxidation- Reduction: Review of "oxidation state" concept. oxidation state in molecules calculated with 6 simple rules: a. sum up the oxidation states of individual atoms (e.g., for SO 2-4 : ox state O= -2, ox state S= +6) 12
Rules continued b. covalent bond has oxidation state of 0 (e.g., N, O and H in O 2, N 2 and H 2, respectively). c. oxidation state of O is -2, except peroxides (H 2 O 2, Na 2 O 2 ) d. H is +1 except: 1. Covalent (0) in H 2 2. -1 in metal- H bonds ("hydride ) e. halogens are -1 in compounds lacking H and O, except: 1. Covalent (0) in e.g. Cl 2 f. Assume alkali metals (e.g., Ba), alkaline earths (e.g., Ca) are+1 and +2, respectively. Caveat Oxidation state for ionic bonds, pretty close to reality Example: NaCl Cl = -1, Na = +1, highly ionic bond; ox. states balance molecule but covalent. Example: N-O O = -2, N = +2 No "donation" is really taking place, but ox. states used to balance the net flow of electrons in redox reactions 13
Redox and Half Reactions Balance elements & electrons: NO + ClO - NO 2 + Cl - From table of half rxns: NO + H 2 O NO 2 + 2e - + 2H + "L E O : oxidation ClO - + 2H + + 2e - Cl - + H 2 O "G E R : reduction H 2 O and 2H + cancel out (sometimes omitted for ½ rxn) 14