Sustainability & Environment (Agri-Food Systems) · Foundations · also Chemistry of the Environment, Aquatic and Atmospheric Chemistry, Environmental Chemistry & Analysis
The chemistry of water, soil and air that pollution, treatment and climate subjects assume: equilibria, the carbonate system, redox, gas solubility, ozone and smog chemistry. GATE ES Section 2 in full.
In the exams: GATE ES Section 2; UGC NET Environmental Sciences Unit 2; ICAR AIEEA-PG Physical Science Unit II (chemistry)
Environmental chemistry studies the sources, reactions, transport, effects and fate of chemical species in water, soil and air. It supplies the quantitative tools that pollution control, water treatment, soil management and climate science assume: concentration units and activities, equilibrium and kinetics, acid–base and carbonate chemistry, gas solubility, precipitation and complexation, redox chemistry, speciation of contaminants, the measurement of water quality, soil exchange chemistry, atmospheric composition, ozone and smog chemistry, stratospheric ozone depletion, and radioactivity.
This manual concentrates on the chemistry and its standard calculations. Applications are developed in later subjects of the track: water quality and nonpoint pollution, water and wastewater treatment, air and noise pollution, soil science for sustainability, soil contamination and remediation, pesticides and the environment, and climate system science. Where a formula has variants, such as Henry's law constants expressed in different directions or alkalinity computed with or without minor species, the stem of each question states the form used.
The subject is examined in full in GATE Environmental Science and Engineering Section 2, with numerical answer type questions on pH, alkalinity, hardness, ionic strength, Henry's law, solubility products, BOD and unit conversions. UGC NET Environmental Sciences Unit 2 and CSIR-style papers test principles and reactions; ICAR AIEEA-PG and ICAR AICE ask for soil and water chemistry relevant to agriculture, such as CEC, SAR and water quality standards. Constants are given at 25 °C unless stated.
Environmental behaviour follows from bonding. Ionic bonds hold salts such as sodium chloride, which dissociate in water into hydrated ions. Covalent bonds join atoms in molecules; when electrons are shared unequally, bonds are polar, and water's bent, polar structure makes it an excellent solvent for ions and polar molecules and gives it hydrogen bonding, high heat capacity and high surface tension. Non-polar organic compounds such as oils and many pesticides dissolve poorly in water and partition into organic matter and fat, which governs their transport and bioaccumulation. Concentrations are expressed in several ways. Mass per volume (mg/L) is standard for water; in dilute water, 1 mg/L ≈ 1 part per million (ppm) by mass, and 1 µg/L ≈ 1 ppb. Molarity (mol/L) counts moles of solute per litre. Normality counts equivalents per litre, where an equivalent is the amount that supplies or reacts with one mole of charge (or H⁺); equivalent weight = molar mass ÷ charge (or number of replaceable H⁺ or OH⁻). Milliequivalents per litre (meq/L) are used for major ions so that cations and anions can be balanced: in a correct analysis, Σ cation meq/L ≈ Σ anion meq/L. Hardness and alkalinity are conventionally expressed 'as CaCO3': the mass of calcium carbonate that would contain the same number of equivalents. Because the equivalent weight of CaCO3 is 100/2 = 50, mg/L as CaCO3 = (mg/L of species) × 50 ÷ (equivalent weight of the species). For example, 40 mg/L of Ca²⁺ (equivalent weight 20) is 40 × 50/20 = 100 mg/L as CaCO3. In air, gas concentrations are given by volume (ppm, ppb) or by mass per volume (µg/m³), converted at 25 °C and 1 atm by µg/m³ = ppb × molar mass ÷ 24.45.
In concentrated or salty solutions, ions interact electrostatically, so their effective concentration, the activity, is less than their analytical concentration. Activity a = γ × c, where γ is the activity coefficient (1 in very dilute solution). Equilibrium constants are strictly defined with activities. The ionic strength of a solution measures its total ionic concentration weighted by charge: I = ½ Σ cᵢzᵢ², where c is molar concentration and z the charge. For 0.01 M CaCl2, [Ca²⁺] = 0.01 M with z = 2 and [Cl⁻] = 0.02 M with z = 1, so I = ½(0.01 × 4 + 0.02 × 1) = 0.03 M. Freshwaters typically have ionic strengths of about 0.001–0.01 M and seawater about 0.7 M. Activity coefficients are estimated with the Debye–Hückel limiting law, log γ = −A z² √I, with A ≈ 0.51 at 25 °C, valid for I below about 0.005 M; the extended Debye–Hückel and Davies equations extend the range to about 0.1 and 0.5 M. Using the limiting law for Ca²⁺ at I = 0.03 gives log γ = −0.51 × 4 × √0.03 ≈ −0.353, so γ ≈ 0.44, showing how strongly divalent ions are affected. An empirical estimate of ionic strength from electrical conductivity or total dissolved solids (I ≈ 2.5 × 10⁻⁵ × TDS in mg/L) is often used for natural waters. Activity corrections matter for calculating mineral saturation, such as calcite scaling and corrosion indices.
Whether a reaction can proceed is determined by thermodynamics. The Gibbs free energy change ΔG = ΔH − TΔS combines enthalpy (heat) and entropy (disorder); a reaction is spontaneous when ΔG is negative. The standard free energy change is related to the equilibrium constant by ΔG° = −RT ln K, where R = 8.314 J mol⁻¹ K⁻¹. A reaction with K = 1000 at 25 °C has ΔG° = −8.314 × 298.15 × ln 1000 ≈ −17.1 kJ/mol. At any moment, ΔG = ΔG° + RT ln Q, where Q is the reaction quotient; if Q < K the reaction proceeds forward, if Q > K it proceeds backward, and at equilibrium Q = K and ΔG = 0. Le Chatelier's principle states that a system at equilibrium responds to a disturbance by shifting so as to counteract it: adding a reactant shifts the equilibrium towards products; raising temperature favours the endothermic direction. The van 't Hoff equation quantifies the temperature dependence of K: ln(K2/K1) = −(ΔH°/R)(1/T2 − 1/T1). Many environmental equilibria, such as gas solubility, acid dissociation and mineral solubility, change significantly with temperature. Environmental systems are often not at equilibrium, because reactions are slow, inputs and outputs are continuous, or biological processes drive them away from equilibrium. Equilibrium calculations nevertheless show the direction in which a system will tend to move and are the basis of chemical speciation models.
Kinetics describes how fast reactions proceed. For a reaction A → products, the rate law r = −d[A]/dt = k[A]ⁿ defines the order n. In zero-order reactions the rate is constant (C = C0 − kt), as when an enzyme is saturated. In first-order reactions the rate is proportional to concentration: C = C0e^(−kt), with half-life t½ = ln 2/k = 0.693/k, independent of concentration. Radioactive decay, many pesticide degradation processes and BOD exertion are described as first order. In second-order reactions (rate = k[A]²), 1/C = 1/C0 + kt, and the half-life depends on initial concentration. Pseudo-first-order kinetics arise when one reactant is in large excess. For example, a pesticide with a half-life of 10 days decays first-order to 100 × (½)³ = 12.5 per cent of its initial concentration after 30 days; its rate constant is 0.693/10 = 0.0693 per day. Temperature strongly affects rates. The Arrhenius equation, k = A e^(−Ea/RT), relates the rate constant to the activation energy Ea. In environmental engineering a simpler form, kT = k20 θ^(T−20), with θ often about 1.02–1.10 (1.047 is a common value for BOD), adjusts rate constants from 20 °C to other temperatures. Catalysts, including enzymes, lower the activation energy without changing the equilibrium. Kinetics determines persistence of contaminants, the design of treatment reactors, and the rates of natural processes such as weathering and decomposition.
Water self-ionises: H2O ⇌ H⁺ + OH⁻, with Kw = [H⁺][OH⁻] = 10⁻¹⁴ at 25 °C, so neutral water has pH 7 at that temperature (pH = −log[H⁺] or, strictly, −log of the hydrogen-ion activity). Kw increases with temperature, so neutral pH is slightly below 7 in warm water. pH + pOH = 14 at 25 °C. Strong acids and bases dissociate completely; weak acids such as acetic acid, carbonic acid and hypochlorous acid dissociate partially, with acid dissociation constant Ka and pKa = −log Ka. At pH = pKa, the acid and its conjugate base are present in equal amounts; at pH one unit above pKa, the base form is about 91 per cent; at two units above, about 99 per cent. Hypochlorous acid (pKa ≈ 7.5) illustrates why this matters: below pH 7.5 the more effective disinfectant HOCl dominates, above it the weaker OCl⁻. A buffer resists changes in pH on addition of acid or base and consists of a weak acid and its conjugate base. The Henderson–Hasselbalch equation gives its pH: pH = pKa + log([A⁻]/[HA]). An acetate buffer with 0.2 M acetate and 0.1 M acetic acid (pKa 4.76) has pH = 4.76 + log 2 ≈ 5.06. Buffers are most effective within about one pH unit of the pKa. Natural waters are buffered mainly by the carbonate system, and soils by clays, organic matter and carbonates.
Alkalinity is the acid-neutralising capacity of water: the amount of strong acid needed to lower the pH to the endpoint of carbonic acid (about 4.3–4.5). For most natural waters, alkalinity (in eq/L) ≈ [HCO3⁻] + 2[CO3²⁻] + [OH⁻] − [H⁺], and it is usually reported as mg/L as CaCO3; borate, silicate, phosphate and organic bases add minor contributions. Acidity is the opposite: base-neutralising capacity. Alkalinity is important for coagulation (which consumes alkalinity), corrosion control, nitrification (which consumes about 7.1 mg alkalinity as CaCO3 per mg ammonia-N oxidised) and the buffering of lakes against acid rain. Alkalinity is measured by titration with standard acid. Phenolphthalein alkalinity (P) is measured to about pH 8.3 and total alkalinity (T) to about pH 4.5 using methyl orange or a pH meter. Assuming no other significant bases, the forms can be inferred: if P = 0, all alkalinity is bicarbonate; if P < T/2, carbonate = 2P and bicarbonate = T − 2P; if P = T/2, all is carbonate; if P > T/2, hydroxide = 2P − T and carbonate = 2(T − P); and if P = T, all is hydroxide. For example, with P = 40 and T = 200 mg/L as CaCO3, carbonate = 80 and bicarbonate = 120 mg/L as CaCO3. Hardness is the concentration of multivalent cations, mainly calcium and magnesium, expressed as mg/L as CaCO3. Carbonate (temporary) hardness is the part equivalent to the alkalinity, removable by boiling; non-carbonate (permanent) hardness is the rest. When total hardness exceeds alkalinity, carbonate hardness equals alkalinity; when alkalinity exceeds hardness, all hardness is carbonate hardness. Waters are commonly classed as soft (below about 75 mg/L), moderately hard (75–150), hard (150–300) and very hard (above 300 mg/L as CaCO3).
Carbon dioxide dissolves in water and reacts: CO2(g) ⇌ CO2(aq); CO2(aq) + H2O ⇌ H2CO3; H2CO3 ⇌ H⁺ + HCO3⁻; HCO3⁻ ⇌ H⁺ + CO3²⁻. Dissolved CO2 and true carbonic acid are usually lumped as H2CO3*. At 25 °C in fresh water, pKa1 ≈ 6.35 and pKa2 ≈ 10.33. A distribution diagram of the fractions α0 (H2CO3*), α1 (HCO3⁻) and α2 (CO3²⁻) against pH shows that H2CO3* dominates below pH 6.35, bicarbonate between 6.35 and 10.33, and carbonate above 10.33. At pH 8.3, bicarbonate makes up about 98 per cent of the total, which is why phenolphthalein alkalinity endpoints are set near that pH. The total dissolved inorganic carbon is CT = [H2CO3*] + [HCO3⁻] + [CO3²⁻]. In an open system in contact with the atmosphere, H2CO3* is fixed by Henry's law. With a Henry's constant for CO2 of about 10⁻¹·⁴⁷ mol L⁻¹ atm⁻¹ and an atmospheric partial pressure of about 420 × 10⁻⁶ atm, [H2CO3*] ≈ 1.4 × 10⁻⁵ M; pure water in equilibrium with this air has [H⁺] ≈ √(Ka1 × [H2CO3*]), giving pH ≈ 5.6. This is why unpolluted rain is slightly acidic, and why 'acid rain' is conventionally defined as rain with pH below about 5.6, caused by sulphuric and nitric acids from SO2 and NOx emissions. In a closed system, CT is constant and pH changes shift the proportions of species. Calcium carbonate dissolution and precipitation are coupled to this system: CaCO3 + CO2 + H2O ⇌ Ca²⁺ + 2HCO3⁻. Increased CO2 dissolves limestone, forming caves and hard water; loss of CO2 precipitates calcite, as in scale in pipes and boilers. Saturation indices, such as the Langelier saturation index, indicate whether water will deposit or dissolve calcium carbonate and are used for corrosion control.
The ocean has absorbed a large share of human CO2 emissions, as the climate system science subject explains. Seawater chemistry differs from fresh water: it has high alkalinity and ionic strength, and the apparent dissociation constants are different (pK1* ≈ 5.9 and pK2* ≈ 9.0 at typical surface conditions), so at pH about 8.1 roughly 90 per cent of dissolved inorganic carbon is bicarbonate, about 9 per cent carbonate and under 1 per cent dissolved CO2. As CO2 dissolves, it forms carbonic acid, which releases H⁺; most of that H⁺ combines with carbonate ions to form bicarbonate. The net effect, CO2 + H2O + CO3²⁻ → 2HCO3⁻, is a fall in pH and in carbonate ion concentration. Surface ocean pH has fallen by about 0.1 units since pre-industrial times (from about 8.2 to 8.1), equivalent to about a 26 per cent increase in hydrogen-ion concentration because pH is logarithmic. The saturation state of seawater with respect to a carbonate mineral is Ω = [Ca²⁺][CO3²⁻] ÷ Ksp′, where Ksp′ is the stoichiometric solubility product. Ω greater than 1 means supersaturation; below 1, the mineral tends to dissolve. Aragonite, used by corals and many molluscs, is more soluble than calcite, so it becomes undersaturated first, especially in cold, high-latitude and upwelling waters. The ocean's buffer capacity, measured by the Revelle factor, declines as carbonate is consumed, reducing its ability to absorb further CO2. Acidification therefore links atmospheric chemistry to marine ecosystems, fisheries and aquaculture, as the climate impacts subject explains.
Henry's law states that, at equilibrium and constant temperature, the concentration of a gas dissolved in a liquid is proportional to its partial pressure above the liquid. It is written in two common forms, which must not be confused: C = KH × P, with KH in mol L⁻¹ atm⁻¹ (a solubility constant, larger for more soluble gases), or P = H × C (or P = H × x, with x the mole fraction), with H in atm L mol⁻¹ or atm (a volatility constant, larger for less soluble gases). A dimensionless Henry constant, Hcc = Cgas/Cwater, is used in air stripping design. Oxygen is sparingly soluble. With KH for O2 of about 1.3 × 10⁻³ mol L⁻¹ atm⁻¹ at 25 °C and a partial pressure of 0.21 atm, the saturation concentration is 1.3 × 10⁻³ × 0.21 × 32,000 mg/mol ≈ 8.7 mg/L, close to the tabulated 8.3 mg/L at 25 °C (the difference reflects the rounded constant and water vapour pressure). Saturation dissolved oxygen in fresh water is about 14.6 mg/L at 0 °C, 9.1 mg/L at 20 °C and 7.6 mg/L at 30 °C; it is lower in salty water and at high altitude, where atmospheric pressure is lower. This decline with temperature is why warm water and thermal pollution stress fish and why rivers are most vulnerable to low oxygen in summer. Highly soluble gases, such as ammonia and sulphur dioxide, react with water and have large solubility constants. Gas transfer between air and water drives aeration in treatment plants, the stripping of ammonia and volatile organic compounds, methane release from wetlands and reservoirs, and ocean uptake of CO2.
Precipitation and dissolution of solids control the concentrations of many ions. For a salt MaXb dissolving as aM + bX, the solubility product Ksp = [M]^a[X]^b at saturation. For a 1:1 salt such as CaCO3, the molar solubility s = √Ksp; for a 1:2 salt such as CaF2 (Ksp ≈ 3.9 × 10⁻¹¹), Ksp = s(2s)² = 4s³, so s = (Ksp/4)^(1/3) ≈ 2.1 × 10⁻⁴ M, about 17 mg/L of CaF2. The common-ion effect lowers solubility when one of the ions is already present: adding fluoride to calcium-rich water lowers calcium fluoride solubility. If the ion activity product exceeds Ksp, precipitation tends to occur. Precipitation is used in water treatment: lime softening raises pH to precipitate CaCO3 and Mg(OH)2; heavy metals are removed as hydroxides, carbonates or sulphides by adjusting pH, each metal having a pH of minimum solubility; phosphorus is removed by precipitation with aluminium or iron salts or with lime; fluoride is removed by precipitation and adsorption (the Nalgonda technique, developed in India, uses alum and lime). Amphoteric hydroxides, such as those of aluminium and zinc, redissolve at high pH. Complexation occurs when a central metal ion binds ligands, molecules or ions that donate electron pairs, such as hydroxide, chloride, carbonate, ammonia, and organic ligands such as humic substances and EDTA. Chelates are complexes in which one ligand binds through several atoms, forming rings; they are very stable. Complexation increases the solubility and mobility of metals, can reduce their bioavailability and toxicity (because the free ion is usually the most toxic form), and affects treatment: EDTA-complexed metals resist precipitation. The hard and soft acids and bases concept explains preferences: hard metals such as Ca²⁺ and Al³⁺ bind oxygen ligands, while soft metals such as Hg²⁺ and Cd²⁺ bind sulphur strongly.
Oxidation is the loss of electrons and reduction the gain; every redox reaction pairs both. Oxidation states track electron distribution: nitrogen ranges from −3 in ammonium (NH4⁺) through 0 in N2, +3 in nitrite (NO2⁻) to +5 in nitrate (NO3⁻); sulphur from −2 in sulphide to +6 in sulphate; carbon from −4 in methane to +4 in CO2; iron is +2 in ferrous and +3 in ferric compounds; manganese +2 to +4; chromium +3 or +6; arsenic +3 (arsenite) or +5 (arsenate). Redox potential is measured as E (volts, relative to the standard hydrogen electrode) or expressed as pe, the negative logarithm of electron activity: pe = E/0.0592 at 25 °C, so E = 0.4 V corresponds to pe ≈ 6.8. The Nernst equation relates potential to concentrations: E = E° − (0.0592/n) log Q at 25 °C. High (positive) pe indicates oxidising conditions; low (negative) pe reducing conditions. pe–pH (or Eh–pH, Pourbaix) diagrams show which species of an element is stable under given redox and pH conditions, bounded by the stability limits of water (oxidation to O2 at the top, reduction to H2 at the bottom). For iron, oxidising, near-neutral conditions favour insoluble Fe(III) oxyhydroxides, while reducing or acidic conditions favour soluble Fe²⁺; this explains why anoxic groundwater can carry dissolved iron that precipitates as orange deposits when exposed to air.
Microorganisms oxidise organic matter using the electron acceptor that yields the most energy among those available. When oxygen is depleted, they use others in a characteristic sequence, the redox ladder: aerobic respiration (O2 reduced to water); denitrification (nitrate reduced to N2 and N2O); manganese reduction (Mn(IV) to Mn²⁺); iron reduction (Fe(III) to Fe²⁺); sulphate reduction (SO4²⁻ to sulphide); and finally methanogenesis (CO2 reduced to, or acetate split into, methane). Energy yield declines down the ladder. This sequence produces redox zonation in space and time: in lake sediments and stratified water bodies, in aquifers downstream of organic pollution plumes such as landfill leachate, in wetlands, and in flooded rice paddies after flooding. In a paddy soil, oxygen disappears within a day or so, then nitrate is lost by denitrification, then manganese and iron are reduced (increasing soluble iron, which can cause iron toxicity in rice), then sulphate is reduced, and finally methane is produced, which is why rice paddies are a methane source and why draining them periodically (alternate wetting and drying) cuts methane. Redox conditions control the mobility and toxicity of many contaminants. Arsenic is mobilised in reducing aquifers, such as those of the Bengal basin and the Indo-Gangetic Plains, when iron oxyhydroxides that hold arsenic are reductively dissolved. Chromium(VI), mobile and toxic, can be reduced to less mobile and less toxic Cr(III) by reducing agents, a basis for remediation. Mercury is methylated by sulphate-reducing and other anaerobic bacteria to methylmercury, which bioaccumulates. Nitrate is removed by denitrification in anoxic zones such as riparian wetlands.
Speciation is the distribution of an element among its different chemical forms: free ions, complexes, oxidation states, adsorbed and precipitated forms, and organometallic compounds. Total concentration alone is a poor guide to mobility, bioavailability and toxicity, because these depend on the species present. Key examples: chromium(VI), present as chromate and dichromate anions, is highly soluble, mobile and carcinogenic, whereas chromium(III) is less soluble and an essential trace nutrient at low levels; tannery and electroplating wastes are major sources. Arsenite, As(III), is more toxic and generally more mobile than arsenate, As(V), which adsorbs more strongly to iron oxides. Mercury occurs as elemental Hg⁰ (volatile), inorganic Hg²⁺ and methylmercury (CH3Hg⁺); methylmercury is the most toxic and bioaccumulative form and caused Minamata disease in Japan in the 1950s. Lead, cadmium and copper are mostly bound to particles and organic matter at neutral pH but become more mobile under acidic conditions. Free metal ion activity, rather than total metal, often best predicts toxicity to aquatic organisms (the free ion activity model and the biotic ligand model). Nutrients are also speciated: nitrogen as ammonium, ammonia, nitrite, nitrate and organic N; un-ionised ammonia (NH3) is far more toxic to fish than ammonium (NH4⁺), and its share rises with pH and temperature (the pKa of ammonium is about 9.25 at 25 °C). Phosphorus occurs as orthophosphate species, polyphosphates and organic P, with orthophosphate the form directly available to algae.
Organic contaminants, such as pesticides, polycyclic aromatic hydrocarbons, polychlorinated biphenyls, pharmaceuticals and industrial solvents, are distributed among water, air, soil, sediment and organisms according to their properties. Water solubility, vapour pressure and the Henry's law constant govern partitioning between water and air. The octanol–water partition coefficient, Kow (usually given as log Kow), measures hydrophobicity: compounds with log Kow above about 4–5 tend to bioaccumulate in fatty tissues and sorb strongly to organic matter. Sorption to soils and sediments is described by the distribution coefficient Kd = Csorbed/Cwater (L/kg). For non-polar organics, sorption is mainly to organic carbon, so Kd ≈ Koc × foc, where Koc is the organic-carbon partition coefficient and foc the fraction of organic carbon in the soil. A pesticide with Koc = 500 L/kg in a soil with 2 per cent organic carbon has Kd = 500 × 0.02 = 10 L/kg. Sorption slows transport through soil and aquifers, described by the retardation factor R = 1 + (ρb/θ)Kd, where ρb is bulk density and θ porosity or water content: with ρb = 1.6 kg/L and θ = 0.4, R = 1 + 4 × 10 = 41, meaning the pesticide moves 41 times more slowly than the water. The pesticides and soil contamination subjects use these tools to assess leaching risk, for example through the groundwater ubiquity score. Persistence depends on degradation by hydrolysis, photolysis and biodegradation, often approximated as first-order with a half-life. Persistent organic pollutants (POPs) combine persistence, bioaccumulation, toxicity and long-range transport; the Stockholm Convention (2001) restricts them, starting with the 'dirty dozen' including DDT, aldrin, dieldrin, PCBs, dioxins and furans. Ionisable organics, such as some herbicides and pharmaceuticals, change their sorption and mobility with pH depending on their pKa.
Water quality is described by physical, chemical and biological parameters. Physical parameters include temperature, turbidity (measured by light scattering in nephelometric turbidity units, NTU), colour, taste and odour, and solids: total solids, total dissolved solids (TDS, the residue after filtering and evaporating at 180 °C) and total suspended solids (TSS, retained on a filter). Electrical conductivity (EC) rises with dissolved ions; TDS (mg/L) is roughly 0.55–0.7 times EC (µS/cm), the factor depending on the ions present. Chemical parameters include pH, alkalinity, hardness, major ions, nutrients (nitrogen and phosphorus species), dissolved oxygen, organic matter indicators, and toxic substances such as fluoride, arsenic, heavy metals and pesticides. Dissolved oxygen (DO) is measured by the Winkler (iodometric) titration or by membrane and optical electrodes. Biochemical oxygen demand (BOD) is the oxygen consumed by microorganisms decomposing organic matter, conventionally measured over 5 days at 20 °C (BOD5) in diluted samples with seed microorganisms: BOD = (D1 − D2)/P, where D1 and D2 are initial and final DO and P the decimal fraction of sample in the bottle. With D1 = 8.6 mg/L, D2 = 4.1 mg/L and 6 mL of sample in a 300 mL bottle (P = 0.02), BOD5 = 4.5/0.02 = 225 mg/L. BOD exertion is modelled as first order: BODt = L0(1 − e^(−kt)), where L0 is the ultimate carbonaceous BOD and k the rate constant (often about 0.23 per day, base e, at 20 °C for domestic wastewater). With L0 = 200 mg/L and k = 0.23 per day, BOD5 = 200(1 − e^(−1.15)) ≈ 137 mg/L. Chemical oxygen demand (COD) uses a strong chemical oxidant (potassium dichromate in acid with silver catalyst) and is completed in about two hours; it is always at least as large as BOD, and the BOD/COD ratio indicates biodegradability (above about 0.5 readily biodegradable). Total organic carbon (TOC) is measured by combustion or oxidation and CO2 detection.
Analytical methods include titrimetry (alkalinity, hardness by EDTA, chloride, DO), gravimetry (solids), potentiometry (pH and ion-selective electrodes such as fluoride), spectrophotometry based on the Beer–Lambert law (A = εbc, absorbance proportional to concentration), used for nitrate, phosphate (molybdenum blue), ammonia and many others, atomic absorption spectrometry (AAS) and inductively coupled plasma techniques (ICP-OES, ICP-MS) for metals, ion chromatography for anions, and gas and liquid chromatography (GC, HPLC), often with mass spectrometry, for organic contaminants. Quality assurance uses blanks, standards, replicates, spikes and certified reference materials; detection limits must be below the standards being checked. The Standard Methods for the Examination of Water and Wastewater (APHA, AWWA and WEF) is the widely used reference compendium. Drinking-water standards are set nationally, often based on WHO Guidelines for Drinking-water Quality. India's standard, IS 10500:2012 of the Bureau of Indian Standards, gives acceptable limits and, where no alternative source exists, permissible limits: for example pH 6.5–8.5; TDS 500 mg/L acceptable and 2000 mg/L permissible; total hardness 200 and 600 mg/L as CaCO3; fluoride 1.0 and 1.5 mg/L; nitrate 45 mg/L as NO3⁻; and arsenic 0.01 mg/L. Irrigation water quality uses EC, SAR and residual sodium carbonate, and surface waters are classified by designated best use with criteria for DO, BOD, pH and coliforms; the water quality subject develops these. Biological parameters, especially indicator bacteria such as coliforms and E. coli, are covered in the environmental microbiology subject.
Soil chemistry governs nutrient supply, pollutant retention and soil structure. Soil organic matter (SOM), including humus formed by the decomposition and transformation of plant and microbial residues, contains roughly 58 per cent carbon (the conventional van Bemmelen factor of 1.724 converts organic carbon to organic matter, though the true factor varies). SOM supplies nitrogen, phosphorus and sulphur, holds water, binds aggregates, complexes metals, sorbs organic chemicals, and has a very high cation exchange capacity. Cation exchange capacity (CEC) is the total quantity of exchangeable cations a soil can hold, expressed in centimoles of charge per kilogram (cmol(+)/kg, numerically equal to meq/100 g). It arises from negative charges on clay minerals and organic matter. Typical values are about 3–15 cmol(+)/kg for kaolinite, 80–120 for smectites such as montmorillonite, and 150–300 or more for humus (per unit mass of the component). Charge on clays comes from isomorphous substitution (permanent charge) and from pH-dependent charge on edges and organic matter, which increases with pH. Exchangeable cations are held with strength roughly in the order Al³⁺ > Ca²⁺ > Mg²⁺ > K⁺ ≈ NH4⁺ > Na⁺ (the lyotropic series), so sodium is easily leached unless it dominates the exchange complex. Base saturation is the percentage of CEC occupied by base cations (Ca²⁺, Mg²⁺, K⁺, Na⁺) as opposed to acidic cations (H⁺, Al³⁺). For a soil with CEC 20 cmol(+)/kg holding 12 Ca, 3 Mg, 0.5 K and 0.5 Na, base saturation = 16/20 = 80 per cent. Base saturation correlates with pH: acidic, leached soils have low base saturation and may have toxic exchangeable aluminium; liming raises base saturation and pH. Anion exchange capacity, generally small, occurs on iron and aluminium oxides in acidic, weathered soils, which also fix phosphate strongly.
Soil pH influences nutrient availability (most nutrients are most available around pH 6–7), microbial activity, aluminium and manganese toxicity in acid soils, and the solubility of contaminants. Acidity arises from leaching of bases, acid rain, nitrification of ammonium fertilisers and organic acids; it is corrected with lime, whose requirement depends on the soil's buffering capacity. Alkaline soils, common in arid regions, may contain calcium carbonate and, if sodic, have pH above about 8.5. Salinity is measured as the electrical conductivity of a saturated paste extract (ECe, dS/m). Sodicity is measured as the exchangeable sodium percentage (ESP = exchangeable Na ÷ CEC × 100) or, for soil solutions and irrigation water, the sodium adsorption ratio: SAR = Na⁺ ÷ √((Ca²⁺ + Mg²⁺)/2), with concentrations in meq/L. For water with 15 meq/L Na⁺, 4 meq/L Ca²⁺ and 2 meq/L Mg²⁺, SAR = 15/√3 ≈ 8.7. The United States Salinity Laboratory classification treats soils with ECe > 4 dS/m as saline and ESP > 15 (roughly SAR > 13) as sodic; saline-sodic soils have both. High sodium disperses clay, destroying structure and reducing infiltration, while high salinity causes osmotic stress. Residual sodium carbonate (RSC = (CO3²⁻ + HCO3⁻) − (Ca²⁺ + Mg²⁺), in meq/L) indicates the tendency of irrigation water to precipitate calcium and increase sodicity; values above about 2.5 meq/L are generally regarded as unsuitable, though national guidelines differ. Reclamation, as the fundamentals and land degradation subjects explain, uses leaching for saline soils and gypsum plus leaching for sodic soils. The soil science subject develops soil fertility and soil health; here the emphasis is on the exchange and solution chemistry that underlies them.
The atmosphere's major components, nitrogen, oxygen and argon, are chemically fairly inert at the surface; most of its chemistry involves trace gases. Carbon dioxide (over 420 ppm), methane (about 1.9 ppm), nitrous oxide (about 0.33 ppm), ozone, carbon monoxide, nitrogen oxides, sulphur dioxide, ammonia and volatile organic compounds (VOCs) are present at ppm to ppt levels but matter greatly for climate, health and ecosystems. The layers and physical structure of the atmosphere are covered in the atmospheric science subject. The hydroxyl radical (OH), produced mainly when ultraviolet light photolyses ozone to excited oxygen atoms that react with water vapour, is called the 'detergent of the atmosphere': it oxidises methane, carbon monoxide, VOCs, SO2 and NO2, initiating their removal. Its global mean concentration is very low (of the order of 10⁶ molecules per cm³) but its reactivity makes it the main sink for many gases. The atmospheric lifetime of a gas is its burden divided by its removal rate; lifetimes range from seconds for OH to about a decade for methane and a century for nitrous oxide. Sulphur dioxide and nitrogen oxides are oxidised to sulphuric and nitric acids, forming acid deposition and secondary particles (sulphate and nitrate aerosols). Ammonia from agriculture neutralises these acids, forming ammonium sulphate and ammonium nitrate particles, a major component of PM2.5 in agricultural regions. Air pollutant concentrations are converted between volume and mass units with µg/m³ = ppb × molar mass ÷ 24.45 at 25 °C and 1 atm; 50 ppb of SO2 (64.06 g/mol) is about 131 µg/m³.
Ozone in the lower atmosphere is a secondary pollutant formed photochemically from nitrogen oxides and VOCs in sunlight. The basic cycle is: NO2 + sunlight (λ < about 420 nm) → NO + O; O + O2 → O3; and NO + O3 → NO2 + O2. On its own this 'null cycle' sets up a photostationary state with no net ozone production. Net ozone forms when VOCs (and CO) are oxidised by OH to peroxy radicals (HO2 and RO2), which convert NO to NO2 without consuming ozone; each conversion then yields another ozone molecule when NO2 is photolysed. Peroxyacetyl nitrate (PAN), aldehydes and secondary organic aerosol are other products. Photochemical smog, first characterised in Los Angeles in the 1940s–1950s by Arie Haagen-Smit, occurs on sunny, warm days with stagnant air and high traffic emissions, and features ozone, PAN, aldehydes and particles; it irritates eyes and lungs and damages crops and materials. It differs from the 'London' or sulphurous smog of cold, damp, foggy conditions, dominated by sulphur dioxide and smoke from coal burning; the London smog of December 1952 caused thousands of excess deaths and led to the UK Clean Air Act of 1956. Ozone production can be limited by NOx or by VOCs. In NOx-limited (VOC-rich) conditions, typical of rural and downwind areas, reducing NOx reduces ozone; in VOC-limited (NOx-rich) conditions, typical of city centres, reducing NOx can briefly increase ozone because less NO is available to remove it ('titration'), so VOC controls are more effective. Ozone peaks in the afternoon and is often higher downwind of cities than within them. Ground-level ozone damages crops such as wheat, rice and soybean, as the climate impacts subject explains, and is a greenhouse gas. Methane, as an ozone precursor, links climate and air quality policy.
About 90 per cent of atmospheric ozone lies in the stratosphere, where it absorbs most solar UV-B and all UV-C radiation. Sydney Chapman (1930) explained its formation and natural destruction: oxygen molecules are split by short-wave UV (wavelengths below about 242 nm) into oxygen atoms (O2 + hν → 2O); an oxygen atom combines with O2 to form ozone (O + O2 + M → O3 + M, where M is a third molecule that removes excess energy); ozone absorbs UV and is split (O3 + hν → O2 + O), heating the stratosphere; and ozone is destroyed by reaction with an oxygen atom (O + O3 → 2O2). The Chapman mechanism overestimates ozone, because catalytic cycles involving hydrogen (HOx), nitrogen (NOx), chlorine (ClOx) and bromine (BrOx) radicals also destroy ozone. In a catalytic cycle, a radical X reacts with ozone and is regenerated: X + O3 → XO + O2, then XO + O → X + O2, with net result O + O3 → 2O2. Because the catalyst is regenerated, one chlorine atom can destroy many thousands of ozone molecules before it is locked into a reservoir compound such as hydrogen chloride or chlorine nitrate. Chlorine atoms come mainly from chlorofluorocarbons (CFCs) and other human-made compounds, which are stable in the troposphere, drift into the stratosphere over years, and are photolysed there by UV, releasing chlorine; bromine from halons and methyl bromide is even more effective per atom. Mario Molina and F. Sherwood Rowland (1974) proposed that CFCs would deplete ozone, and with Paul Crutzen, who showed the role of nitrogen oxides, they received the 1995 Nobel Prize in Chemistry.
The Antarctic ozone hole, reported by Joseph Farman and colleagues of the British Antarctic Survey in 1985, appears each southern spring (September–October). During the polar winter, a vortex isolates very cold air, and polar stratospheric clouds form below about −78 °C. On their surfaces, heterogeneous reactions convert reservoir species (HCl and ClONO2) into reactive chlorine (Cl2 and HOCl). When sunlight returns in spring, these are photolysed to chlorine atoms, which rapidly destroy ozone through cycles including the ClO dimer cycle. Denitrification of the stratosphere by sedimenting cloud particles removes NOx that would otherwise tie up chlorine. Arctic ozone losses are generally smaller and more variable because the Arctic vortex is warmer and less stable. The Vienna Convention for the Protection of the Ozone Layer (1985) and its Montreal Protocol on Substances that Deplete the Ozone Layer (1987), with later amendments, phased out the production of CFCs, halons, carbon tetrachloride, methyl chloroform and, more slowly, hydrochlorofluorocarbons (HCFCs) and methyl bromide. The Montreal Protocol is the first UN treaty ratified by every country. The ozone depletion potential (ODP) of a substance is its ozone-destroying ability relative to CFC-11 (ODP = 1); halons have ODPs of about 3–10, HCFCs about 0.01–0.1, and HFCs zero. Because many ozone-depleting substances are also powerful greenhouse gases, the Protocol has also had large climate benefits. The Kigali Amendment (2016) added a phase-down of hydrofluorocarbons (HFCs), which do not deplete ozone but are potent greenhouse gases. Scientific assessments by WMO and UNEP report that stratospheric chlorine is declining and that the ozone layer is on track to recover to 1980 values by around the middle of this century outside the polar regions and somewhat later over Antarctica, provided the Protocol continues to be implemented. Unexpected emissions of CFC-11 detected in the late 2010s, traced largely to illegal production, were subsequently reduced, showing the importance of monitoring.
Radioactive nuclides are unstable and decay spontaneously, emitting radiation. Alpha particles (helium nuclei) are heavy and highly ionising but are stopped by a sheet of paper or the outer skin; they are dangerous mainly when inhaled or ingested. Beta particles (electrons or positrons) penetrate further and are stopped by a few millimetres of aluminium or plastic. Gamma rays are high-energy electromagnetic radiation that penetrate deeply and require thick lead or concrete for shielding. Radioactive decay is first-order: N = N0e^(−λt), where λ is the decay constant, and the half-life t½ = 0.693/λ. After n half-lives, the fraction remaining is (½)ⁿ: after 3 half-lives, one-eighth remains, so 1000 Bq of a nuclide becomes 125 Bq. Activity is measured in becquerels (Bq, one decay per second); the older unit, the curie, equals 3.7 × 10¹⁰ Bq. Absorbed dose, the energy deposited per kilogram, is measured in grays (Gy, 1 J/kg); equivalent and effective doses, which weight different radiation types and tissues by their biological effect, are measured in sieverts (Sv). Everyone is exposed to natural background radiation from cosmic rays, terrestrial radionuclides (uranium, thorium and their decay products, and potassium-40), internal radionuclides, and above all radon. UNSCEAR estimates the world average natural background effective dose at about 2.4 mSv per year, varying widely by location; parts of the Kerala coast in India with monazite sands have notably high natural background. Medical exposures add a significant artificial dose in many countries.
Radon-222, a radioactive gas with a half-life of about 3.8 days, is produced by the decay of uranium-238 via radium-226 in rocks and soils. It seeps into buildings, especially basements and poorly ventilated houses on granitic or uranium-rich ground, and its short-lived decay products deposit in the lungs; radon is the second leading cause of lung cancer after smoking in many countries. Testing and ventilation reduce exposure. Anthropogenic radionuclides come from nuclear weapons testing (global fallout, peaking in the early 1960s), nuclear accidents and the nuclear fuel cycle. Important ones include iodine-131 (half-life about 8 days), which concentrates in the thyroid and caused thyroid cancers in children after the Chernobyl accident (1986), caesium-137 (about 30 years), which behaves like potassium, contaminated soils and food chains over large areas after Chernobyl and Fukushima (2011), and strontium-90 (about 29 years), which behaves like calcium and accumulates in bones and milk. Their mobility in soils depends on chemistry: caesium binds strongly to certain clay minerals (illite), limiting its uptake by plants in many soils, whereas it is more mobile in organic soils. Countermeasures after contamination include restrictions on food, potassium fertilisation to reduce caesium uptake, and removal of topsoil. Naturally occurring radioactive material (NORM) is concentrated by mining, phosphate fertiliser production (phosphogypsum contains radium) and coal combustion (fly ash). Uranium occurs in groundwater in parts of the world, including areas of Punjab in India, with chemical toxicity to the kidneys as well as radiological effects; the WHO provisional guideline for uranium in drinking water is 30 µg/L. Radioactive waste management, deep geological disposal and nuclear accidents are discussed further in the energy subject.
A student uses this subject to master the calculations that dominate GATE Section 2: mg/L as CaCO3, meq/L balances, ionic strength and activity coefficients, ΔG° from K, first-order half-lives, Henderson–Hasselbalch pH, alkalinity forms from P and T, carbonate fractions, Henry's law gas solubility, Ksp solubility, pe from E, BOD from dilution data and first-order kinetics, Kd and retardation, SAR and base saturation, ppb to µg/m³, and radioactive decay. Common traps are the two forms of Henry's law, alkalinity units, and forgetting that pH 7 is neutral only at 25 °C.
A five-year practitioner, such as a water or wastewater chemist, a pollution control board scientist or a soil testing laboratory manager, uses the subject to run and quality-control analyses, interpret water quality against IS 10500 and irrigation criteria, diagnose corrosion or scaling with saturation indices, choose coagulant and lime doses that account for alkalinity, judge arsenic or chromium risks by speciation, and advise farmers on sodic water and gypsum use.
A senior professional with 25 years' experience, such as the head of a state laboratory network, a treatment plant director or a regulator, uses environmental chemistry to set monitoring programmes and standards, decide treatment strategies for fluoride, arsenic, nitrate and emerging contaminants, assess risks from pesticides and persistent organic pollutants, and design air quality policies that account for NOx- and VOC-limited ozone regimes and ammonia's role in PM2.5.
A researcher with a PhD and 25 years' experience works on contaminant speciation and fate, geochemical modelling, redox processes in aquifers and paddy soils, atmospheric chemistry of ozone and secondary aerosol, ocean carbonate chemistry, and analytical methods for trace and emerging contaminants. Such a researcher develops models and methods used by regulators and trains analytical and modelling teams.
This manual gives equilibrium constants rounded at 25 °C; precise work needs temperature- and ionic-strength-corrected values and geochemical software. It does not derive the full carbonate system algebra, cover isotope geochemistry, or treat colloid and surface complexation chemistry in depth. Drinking-water and irrigation standards are given for India (IS 10500:2012) with WHO references; other countries' limits differ and are revised. Stratospheric chlorine trends and ozone recovery dates follow the 2022 assessment and will be updated. Analytical methods are described by principle, not by procedure; laboratory practice should follow Standard Methods or national protocols. Emerging contaminants such as PFAS and microplastics are left to the water quality and contamination subjects.
Original practice questions written to the syllabus and pattern of the exams named in each question's examTags. They are not past papers and do not reproduce any board's questions.
These notes are free to read here, with no account.
What the Laisha app adds:
Free for ten days, then a subscription. We would rather write that here than have you find it later.