UGC NET Environmental Science Unit-II: Environmental Chemistry Syllabus

Unit-II: Environmental Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Unit-II: Environmental Chemistry in UGC NET Environmental Science, not a summary of it.

1. 1

   Fundamentals of Environmental Chemistry

   10 topics

   • Classification of Elements
   • Stoichiometry
   • Gibbs’ Energy
   • Chemical Potential
   • Chemical Kinetics
   • Chemical Equilibria
   • Solubility of Gases in Water
   • The Carbonate System
   • Unsaturated and Saturated Hydrocarbons
   • Radioisotopes
2. 2

   Composition of Air

   6 topics

   • Particles, Ions and Radicals in the Atmosphere
   • Chemical Speciation
   • Chemical Processes in the Formation of Inorganic and Organic Particulate Matters
   • Thermochemical and Photochemical Reactions in the Atmosphere
   • Oxygen and Ozone Chemistry
   • Photochemical Smog
3. 3

   Hydrological Cycle

   4 topics

   • Water as a Universal Solvent
   • Concept of DO, BOD and COD
   • Sedimentation, Coagulation, Flocculation, Filtration
   • pH and Redox Potential (Eh)
4. 4

   Inorganic and Organic Components of Soils

   1 topic

   • Biogeochemical Cycles – Nitrogen, Carbon, Phosphorus and Sulphur
5. 5

   Toxic Chemicals

   4 topics

   • Pesticides and Their Classification and Effects
   • Biochemical Aspects of Heavy Metals (Hg, Cd, Pb, Cr) and Metalloids (As, Se)
   • CO, O3, PAN, VOC and POP
   • Carcinogens in the Air
6. 6

   Principles of Analytical Methods

   14 topics

   • Titrimetry
   • Gravimetry
   • Bomb Calorimetry
   • Chromatography
   • Flame Photometry
   • Spectrophotometry
   • Electrophoresis
   • XRF
   • XRD
   • NMR
   • FTIR
   • GC-MS
   • SEM
   • TEM

Unit-II: Environmental Chemistry flashcards for UGC NET Environmental Science

20 of 54 cards from the Unit-II: Environmental Chemistry deck — real questions with worked answers.

1. In the periodic classification of elements, on what basis are elements arranged in the modern periodic table, and what is the modern periodic law?

   Elements are arranged in order of increasing atomic number (number of protons). Modern periodic law: the physical and chemical properties of elements are periodic functions of their atomic numbers.

2. What is stoichiometry, and what fundamental law is it based on?

   Stoichiometry is the quantitative study of reactants and products in a chemical reaction. It is based on the law of conservation of mass, where moles of reactants and products are related by the balanced equation's coefficients.

3. Define Gibbs free energy (G) and give the equation linking it to enthalpy and entropy.

   Gibbs free energy is the maximum non-expansion work obtainable from a system at constant temperature and pressure. G = H − TS, and for a process ΔG = ΔH − TΔS.

4. What does the sign of ΔG indicate about the spontaneity of a reaction?

   ΔG < 0: spontaneous (exergonic); ΔG > 0: non-spontaneous (endergonic); ΔG = 0: system at equilibrium.

5. Define chemical potential (μ) of a species.

   Chemical potential is the partial molar Gibbs free energy — the change in Gibbs energy of a system when one mole of a component is added at constant temperature, pressure, and amounts of other components: μ_i = (∂G/∂n_i)_T,P,nj.

6. How is chemical potential related to activity (or concentration) of a species?

   μ = μ° + RT ln a, where μ° is the standard chemical potential, a is the activity (≈ concentration for dilute solutions), R is the gas constant, and T is temperature.

7. State the rate law expression and define order and molecularity in chemical kinetics.

   Rate = k[A]^m[B]^n. Order is the sum of exponents (m+n), determined experimentally. Molecularity is the number of reacting species in an elementary step, always a whole number ≥ 1.

8. Give the integrated rate equation and half-life expression for a first-order reaction.

   Integrated: ln[A] = ln[A]₀ − kt, or k = (2.303/t) log([A]₀/[A]). Half-life: t½ = 0.693/k, which is independent of initial concentration.

9. What is the Arrhenius equation and what does it describe?

   k = A·e^(−Ea/RT), where k is the rate constant, A is the pre-exponential (frequency) factor, Ea is activation energy, R the gas constant, and T temperature. It describes the temperature dependence of reaction rate.

10. Write the equilibrium constant expression (Kc) for the general reaction aA + bB ⇌ cC + dD.

    Kc = ([C]^c[D]^d) / ([A]^a[B]^b), using equilibrium molar concentrations. (Pure solids and liquids are omitted.)

11. State Le Chatelier's principle and how a chemical equilibrium responds to disturbance.

    If a system at equilibrium is disturbed by a change in concentration, temperature, or pressure, the equilibrium shifts in the direction that counteracts the change to re-establish equilibrium.

12. What law governs the solubility of gases in water, and how does temperature affect gas solubility?

    Henry's law: the solubility of a gas is proportional to its partial pressure above the liquid (C = kH·P). Gas solubility decreases as temperature increases.

13. Describe the carbonate system equilibria that buffer natural waters.

    CO₂(g) ⇌ CO₂(aq); CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ ⇌ 2H⁺ + CO₃²⁻. This CO₂–bicarbonate–carbonate system controls pH and alkalinity of natural waters.

14. Differentiate saturated and unsaturated hydrocarbons with examples.

    Saturated hydrocarbons (alkanes, e.g. CH₄, C₂H₆) have only single C–C bonds. Unsaturated hydrocarbons contain double bonds (alkenes, e.g. C₂H₄) or triple bonds (alkynes, e.g. C₂H₂) and are more reactive (addition reactions).

15. What are radioisotopes, and name one common environmental tracer/dating radioisotope with its use.

    Radioisotopes are unstable isotopes that emit radiation (α, β, γ) during decay. Example: Carbon-14 (half-life ~5730 yr) used for radiocarbon dating; ³H, ³²P, ¹³⁷Cs are used as environmental tracers.

16. Distinguish particles, ions, and radicals in the atmosphere.

    Particles are solid/liquid aerosols suspended in air. Ions are charged atoms/molecules (e.g. NO₃⁻, NH₄⁺). Radicals are highly reactive species with unpaired electrons (e.g. OH•, NO₃•, HO₂•) that drive atmospheric reactions.

17. What is chemical speciation and why is it environmentally important?

    Chemical speciation is the distribution of an element among its different chemical forms/species (e.g. oxidation states, complexes). It is important because toxicity, mobility, and bioavailability depend on the species, not just total concentration (e.g. Cr(III) vs Cr(VI)).

18. How do primary and secondary particulate matters differ in their formation?

    Primary particulates are emitted directly into the atmosphere (e.g. dust, soot). Secondary particulates form in the atmosphere via gas-to-particle conversion through chemical reactions (e.g. sulfate from SO₂, nitrate from NOx, secondary organic aerosols).

19. Distinguish thermochemical and photochemical reactions in the atmosphere.

    Thermochemical reactions are driven by heat/thermal energy. Photochemical reactions are initiated by absorption of light (photons), e.g. NO₂ + hν → NO + O, which begins photochemical smog and ozone formation.

20. Explain the Chapman cycle of stratospheric ozone formation and destruction.

    Formation: O₂ + hν → 2O; O + O₂ + M → O₃. Destruction: O₃ + hν → O₂ + O; O + O₃ → 2O₂. This natural cycle maintains the stratospheric ozone layer that absorbs UV radiation.

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