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Química Ambiental · Water and soil

Química da água e do solo

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O que você vai aprender

  • Use the carbonic acid pKa values to find which carbonate species dominates at a given pH and calculate hardness and alkalinity as mg/L CaCO₃.
  • Draw the polar mechanisms of chlorine hydrolysis, chloramine formation, phenol chlorination and CO₂ hydration, and explain how pH sets the HOCl/OCl⁻ ratio.
  • Explain how disinfection byproducts form and draw the hydrolysis and substitution reactions that degrade esters and alkyl halides in natural waters.
  • Order the electron acceptors of natural waters, convert potentials to pE and calculate oxygen demand.
  • Predict how pH, redox state and ligand softness control the solubility, mobility and toxicity of metals and metalloids in water and soil.

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The carbonate system

Dissolved CO₂ reacts with water: CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ ⇌ 2 H⁺ + CO₃²⁻. Treating dissolved CO₂ and H₂CO₃ together, pKa1 ≈ 6.35 and pKa2 ≈ 10.33 at 25 °C. Below pH 6.35 dissolved CO₂ dominates, between 6.35 and 10.33 bicarbonate dominates, and above 10.33 carbonate dominates; at pH 8, typical of rivers and seawater, most inorganic carbon is HCO₃⁻ (about 97 % in fresh water, about 90 % in seawater, whose constants are shifted by its salt content). The ratio of any adjacent pair follows from pH − pKa = log([base]/[acid]). Uncatalysed hydration of CO₂ is slow (seconds to minutes): water adds to the carbon, and the zwitterion loses a proton to another water molecule. Adding CO₂ to seawater consumes carbonate, CO₂ + CO₃²⁻ + H₂O → 2 HCO₃⁻, which is why ocean acidification lowers the CaCO₃ saturation state, makes shells harder to build and dissolves them where seawater becomes undersaturated.

Hardness and alkalinity

Hardness is the total concentration of Ca²⁺ and Mg²⁺, reported as the mass of CaCO₃ with the same number of moles (M = 100.09 g/mol): hardness = (mol/L Ca²⁺ + mol/L Mg²⁺) × 100.09 g/mol. Temporary hardness is the part balanced by bicarbonate; boiling drives off CO₂ and precipitates CaCO₃ (kettle scale). Permanent hardness is balanced by sulfate or chloride and survives boiling. Alkalinity is the acid-neutralising capacity, mainly HCO₃⁻, CO₃²⁻ and OH⁻, measured by titration with strong acid to about pH 4.5 and also reported as mg/L CaCO₃ (one CaCO₃ neutralises two H⁺). Water in contact with limestone gains both: CaCO₃ + CO₂ + H₂O → Ca²⁺ + 2 HCO₃⁻, and more CO₂ dissolves more limestone.

Chlorination and chloramines

Chlorine gas reacts with water: Cl₂ + H₂O → HOCl + H⁺ + Cl⁻, a disproportionation in which one chlorine goes from 0 to +1 and the other to −1. Hypochlorous acid (pKa ≈ 7.5) is the strong disinfectant; hypochlorite ion is tens of times less effective, so disinfection is faster at pH 6.5–7 than at pH 8.5. HOCl and OCl⁻ together are free chlorine. With ammonia, HOCl transfers its chlorine to nitrogen: NH₃ + HOCl → NH₂Cl + H₂O. Monochloramine dominates near pH 8 when the Cl₂:N mass ratio is below about 5:1; it is a weaker but longer-lasting disinfectant that forms far fewer chlorinated byproducts, so many utilities use it in distribution pipes. Adding more chlorine oxidises chloramines to N₂ (2 NH₃ + 3 HOCl → N₂ + 3 H₂O + 3 H⁺ + 3 Cl⁻); the residual falls to a minimum, the breakpoint (Cl₂:N ≈ 7.6:1 by mass), and only then does free chlorine accumulate.

Disinfection byproducts

HOCl is also an electrophilic chlorinating agent. Natural organic matter contains phenols and methyl ketone units; phenoxide ions attack the chlorine of HOCl at the ortho and para positions, hydroxide leaves, and loss of the ring proton restores aromaticity. Repeated chlorination gives 2,4,6-trichlorophenol, and ring cleavage and haloform-type reactions release trihalomethanes (CHCl₃, CHBrCl₂, CHBr₂Cl, CHBr₃) and haloacetic acids. Bromide in the source water is oxidised to HOBr, so brominated trihalomethanes form too. The surest way to reduce these byproducts is to remove the organic precursors (coagulation, activated carbon) before chlorinating; raising the pH lowers the HOCl fraction but speeds trihalomethane formation, which is base-catalysed.

Hydrolysis and substitution of organic pollutants

In natural waters esters hydrolyse mainly by hydroxide attack at the carbonyl carbon (addition, then loss of the alkoxide, then proton transfer to give the carboxylate). The pseudo-first-order rate constant is kOH[OH⁻], so above about pH 7 each pH unit speeds the reaction tenfold, and t½ = ln 2/(kOH[OH⁻]). Reactivity follows the leaving group: amides are the slowest, then alkyl esters, aryl esters and anhydrides. Phthalate plasticisers such as dimethyl phthalate are hydrolysed one ester at a time. Alkyl halides undergo SN2 substitution: in anoxic porewater hydrogen sulfide ion (HS⁻), a much better nucleophile than water, converts bromomethane into methanethiol.

Redox in natural waters

The redox state of water is expressed as pE = E/0.05916 V at 25 °C; high pE is oxidising. When bacteria oxidise organic matter they use electron acceptors in order of decreasing energy yield: O₂, then NO₃⁻ (denitrification to N₂), MnO₂ (to Mn²⁺), Fe(III) oxyhydroxides (to Fe²⁺), SO₄²⁻ (to H₂S) and finally CO₂ (to CH₄). A stratified lake therefore loses O₂ in its bottom water first, then releases soluble Fe²⁺ and Mn²⁺ and finally H₂S. Oxygen is only sparingly soluble (about 9 mg/L at 20 °C, less when warmer), so the oxygen demand of organic waste matters: the theoretical oxygen demand follows from the balanced oxidation, e.g. C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O. Nitrification (NH₄⁺ + 2 O₂ → NO₃⁻ + 2 H⁺ + H₂O) and pyrite oxidation in mine waste both release acid.

Metals and metalloids in water and soil

Metal behaviour depends on speciation. Hydrated Fe³⁺ and Al³⁺ are acids (Fe(H₂O)₆³⁺ pKa ≈ 2.2), and their hydroxides dissolve as pH falls, so acid soils release toxic Al³⁺. Soils hold cations on negatively charged clay and humus sites; the cation-exchange capacity is counted in moles of charge, so one Ca²⁺ fills two sites. Soft metal ions such as Hg²⁺, Cd²⁺ and Pb²⁺ bind sulfur ligands (thiolates, sulfide) far more strongly than oxygen ligands. In anoxic sediments sulfate-reducing bacteria methylate Hg²⁺ to CH₃Hg⁺, which binds protein thiols, crosses membranes and biomagnifies in fish. Oxidation state matters as much as pH: chromate, Cr(VI), is soluble and toxic, whereas Cr(III) precipitates as its hydroxide; arsenite, As(III), is the neutral H₃AsO₃ at pH 7 (pKa ≈ 9.2), while arsenate, As(V), is anionic and adsorbs strongly to iron oxides.

What to carry forward

Natural-water chemistry is acid–base and redox equilibria plus a few polar reactions: pH − pKa sets the carbonate and HOCl/OCl⁻ ratios, HOCl transfers Cl to N, O and C nucleophiles, hydroxide and HS⁻ degrade esters and alkyl halides, the redox ladder decides which acceptor is in use, and pH, oxidation state and ligand softness decide where metals go.

Problemas de prática (40)

Praticar é grátis com uma conta.

Desenhe o mecanismo 6

Byproducts and organic pollutants

  • Bromomethane meets hydrogen sulfide ion in anoxic sedimentDraw the SN2 reaction of hydrogen sulfide ion with bromomethane and explain why HS⁻ outcompetes water even at millimolar concentration.
  • Hydroxide hydrolysis of dimethyl phthalate in natural waterDraw the addition of hydroxide to an ester group of dimethyl phthalate, the loss of methoxide and the proton transfer that gives the methyl phthalate anion and methanol.
  • Chlorination of phenoxide by hypochlorous acidDraw how phenoxide attacks the chlorine of hypochlorous acid through its para carbon and how loss of the ring proton restores the aromatic 4-chlorophenoxide.

Chlorination and disinfection

  • Hydrolysis of chlorine in waterDraw how water attacks dissolved chlorine to give hypochlorous acid, chloride and hydronium, and identify the reaction as a disproportionation.
  • Formation of monochloramine from ammonia and hypochlorous acidDraw the chlorine transfer from hypochlorous acid to ammonia that gives monochloramine and water, and explain why the rate peaks near pH 8.5.

Carbonate equilibria and hardness

  • Hydration of dissolved CO₂ by waterDraw how water adds to dissolved carbon dioxide and how a second water molecule removes a proton to give bicarbonate and hydronium.

Preveja o produto 3

Metals in water and soil

  • Aluminium in an acidified soilPredict how lowering soil pH changes the concentration of dissolved aluminium.

Redox in natural waters

  • Iron in the bottom of a stratified lakePredict how dissolved iron in lake bottom water changes when the water becomes anoxic.

Carbonate equilibria and hardness

  • Groundwater under a forest soilPredict how the CO₂ pressure of soil air changes the amount of limestone dissolved by groundwater.

Classificar 6

Metals in water and soil

  • Arsenite and arsenate at pH 7Classify the dominant forms of arsenic(III) and arsenic(V) in groundwater from their pKa values.

Carbonate equilibria and hardness

  • Inorganic carbon in a river at pH 8Identify the dominant dissolved inorganic carbon species from the pH and the two carbonic acid pKa values.
  • Which dissolved salt is removed by boiling?Classify dissolved salts as temporary hardness, permanent hardness or not hardness at all.

Chlorination and disinfection

  • Which chloramine forms at the plant?Identify the main chlorine species formed when chlorine is dosed into water with excess ammonia at near-neutral pH.

Redox in natural waters

  • What happens to ammonium fertiliser in aerated soil?Classify the conversion of ammonium to nitrate as an oxidation and predict its effect on soil pH.

Byproducts and organic pollutants

  • Spot the trihalomethaneClassify chlorination byproducts and recognise a trihalomethane by its CHX₃ formula.

Comparar 4

Chlorination and disinfection

  • Rank chlorine disinfectantsOrder hypochlorous acid, hypochlorite and monochloramine by their germicidal effectiveness.

Byproducts and organic pollutants

  • Which acyl compounds survive longest in a river?Order acyl compounds by their rate of hydrolysis in natural water from the basicity of the leaving group.

Metals in water and soil

  • Which ligand holds mercury most tightly?Compare ligands for binding to Hg²⁺ using hard–soft acid–base reasoning.

Redox in natural waters

  • Which electron acceptor is used next?Order the electron acceptors of natural waters by the sequence in which bacteria use them to oxidise organic matter.

Calcular 11

Carbonate equilibria and hardness

  • Total alkalinity from a titrationCalculate total alkalinity as mg/L CaCO₃ from a sulfuric acid titration to the pH 4.5 end point.
  • How much inorganic carbon is dissolved CO₂?Calculate the fraction of total inorganic carbon present as dissolved CO₂ from pH and pKa values.
  • Total hardness of a well waterConvert calcium and magnesium concentrations into total hardness expressed as mg/L CaCO₃.

Chlorination and disinfection

  • Theoretical breakpoint doseCalculate the theoretical Cl₂ to ammonia-nitrogen mass ratio at breakpoint from the balanced oxidation of ammonia to N₂.
  • How much free chlorine is HOCl at pH 8?Calculate the percentage of free chlorine present as hypochlorous acid from the pH and its pKa.

Metals in water and soil

  • Calcium held by a soil's exchange sitesConvert a soil's cation-exchange capacity into the mass of calcium it can hold.
  • pH of an iron(III) solutionCalculate the pH of a hydrated iron(III) solution from the pKa of the aqua ion, solving the quadratic when the acid is not weakly ionised.

Byproducts and organic pollutants

  • Half-life of an ester in a lakeCalculate the half-life of a pollutant ester from its second-order hydroxide rate constant and the pH of the water.

Redox in natural waters

  • Theoretical oxygen demand of a sugar wasteCalculate the theoretical oxygen demand of a dissolved organic waste from its balanced oxidation equation.
  • Convert a measured potential to pEConvert a measured redox potential in volts into pE at 25 °C.
  • Electrons released by pyrite oxidationCount the electrons lost when pyrite is fully oxidised to Fe³⁺ and sulfate, using the disulfide oxidation state.

Encontre o erro 1

  • Oxygen and hydrogen sulfide together?Diagnose an inconsistent equilibrium description of a water sample using the redox ladder.

Interprete as evidências 1

  • Reading a breakpoint curveInterpret the falling region of a breakpoint chlorination curve in terms of chloramine oxidation.

Trabalhe de trás para frente 1

  • Which plasticiser was it?Work back from the two hydrolysis products found in a sediment to the plasticiser ester they came from.

Planeje uma síntese 2

Metals in water and soil

  • Stopping a chromate plumeChoose a remediation strategy for chromium in groundwater from the solubility of its two oxidation states.

Byproducts and organic pollutants

  • Cutting trihalomethanes at a river plantChoose the treatment change that lowers trihalomethane formation without weakening disinfection.

Explique por quê 5

Chlorination and disinfection

  • Why switch the residual to chloramine?Explain why utilities use monochloramine rather than free chlorine as the residual disinfectant in long distribution systems.

Metals in water and soil

  • Why add phosphate to drinking water?Explain how orthophosphate dosing lowers lead in tap water through precipitation of an insoluble lead phosphate scale.
  • Why methylmercury is the dangerous formExplain how microbial methylation in sediments turns inorganic mercury into a form that accumulates in fish.

Carbonate equilibria and hardness

  • More carbon, less carbonateExplain why absorbing CO₂ lowers the carbonate ion concentration of seawater even though total dissolved carbon rises.

Redox in natural waters

  • Fish kills in a warm summerExplain why dissolved oxygen falls in warm, nutrient-rich water from gas solubility and respiration.

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