GATE Environmental Engineering Environmental Chemistry Syllabus

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 Environmental Chemistry in GATE Environmental Engineering, not a summary of it.

1. 1

   Fundamentals of Environmental Chemistry

   6 topics

   • Covalent and Ionic Bonding
   • Chemical Equations, Concentration and Activity
   • Structure and Chemistry of Organic Molecules
   • Radioactivity of Elements
   • Chemical Equilibria
   • Thermodynamics and Kinetics of Chemical Reactions
2. 2

   Principles of Water Chemistry

   7 topics

   • Water Quality Parameters and their Measurement
   • Acid-Base Equilibria
   • Buffer Solution
   • Carbonate System
   • Solubility of Gases in Water
   • Complexation, Precipitation, and Redox Reactions
   • Inorganic and Organic Contaminants in Water and their Speciation
3. 3

   Soil Chemistry

   7 topics

   • Organic Matter
   • Nitrogen
   • Phosphorous
   • Potassium
   • Cation Exchange Capacity
   • Base Saturation
   • Sodium Absorption Ratio
4. 4

   Atmospheric Chemistry

   5 topics

   • Composition of the Atmosphere
   • Reactivity of Trace Substances in the Atmosphere
   • Urban Atmosphere—Smog and Particulate Pollution
   • Chemistry of Ozone Formation
   • Chemistry of Stratosphere

Environmental Chemistry flashcards for GATE Environmental Engineering

22 of 51 cards from the Environmental Chemistry deck — real questions with worked answers.

1. What is the fundamental difference between an ionic bond and a covalent bond?

   An ionic bond forms by the complete transfer of electrons from one atom to another, creating oppositely charged ions held by electrostatic attraction (e.g. $\ce{NaCl}$). A covalent bond forms by the sharing of electron pairs between atoms (e.g. $\ce{Cl2}$). Ionic bonds occur between metals and non-metals (large electronegativity difference); covalent bonds occur between non-metals (small difference).

2. How does electronegativity difference ($\Delta\chi$) classify a bond as ionic, polar covalent, or nonpolar covalent?

   Roughly: $\Delta\chi > 1.7$ gives a predominantly ionic bond; $0.4 \leq \Delta\chi \leq 1.7$ gives a polar covalent bond; $\Delta\chi < 0.4$ gives a nonpolar covalent bond.

3. State the formula for percent ionic character of a bond from dipole moment.

   $$\% \text{ ionic character} = \frac{\mu_{\text{observed}}}{\mu_{\text{ionic}}} \times 100$$ where $\mu_{\text{ionic}} = q \cdot d$ is the dipole moment assuming complete charge transfer.

4. Define molarity, molality and normality and give their units.

   Molarity $M = \frac{\text{moles of solute}}{\text{litres of solution}}$ (mol/L). Molality $m = \frac{\text{moles of solute}}{\text{kg of solvent}}$ (mol/kg). Normality $N = \frac{\text{gram equivalents of solute}}{\text{litres of solution}}$ (eq/L). Relation: $N = M \times n$ where $n$ is valence/equivalence factor.

5. What is chemical activity and how does it relate to concentration?

   Activity $a$ is the effective (thermodynamic) concentration of a species: $a = \gamma \, C$, where $\gamma$ is the activity coefficient and $C$ is the molar concentration. For ideal/dilute solutions $\gamma \to 1$ and $a \approx C$; as ionic strength rises, $\gamma < 1$.

6. State the Debye-Hückel limiting law for the activity coefficient of an ion and define ionic strength.

   $$\log \gamma_i = -A z_i^{2} \sqrt{I}$$ where $A \approx 0.509$ (water, $25^\circ\text{C}$), $z_i$ is ion charge, and ionic strength $$I = \frac{1}{2}\sum_i C_i z_i^{2}.$$

7. What does it mean to balance a chemical equation, and what conservation principles must hold?

   Balancing means choosing stoichiometric coefficients so that the number of atoms of each element and the total charge are equal on both sides, satisfying conservation of mass and conservation of charge.

8. Classify organic molecules by their main functional groups: alcohol, aldehyde, ketone, carboxylic acid.

   Alcohol: $\ce{-OH}$ ($\ce{R-OH}$). Aldehyde: $\ce{-CHO}$ (terminal carbonyl). Ketone: $\ce{C=O}$ between two carbons ($\ce{R-CO-R'}$). Carboxylic acid: $\ce{-COOH}$.

9. What is the difference between saturated and unsaturated hydrocarbons? Give the general formulas.

   Saturated hydrocarbons (alkanes) contain only single C-C bonds, general formula $\ce{C_nH_{2n+2}}$. Unsaturated hydrocarbons contain double bonds (alkenes, $\ce{C_nH_{2n}}$) or triple bonds (alkynes, $\ce{C_nH_{2n-2}}$).

10. What is hybridization, and what are the geometries of $sp$, $sp^{2}$ and $sp^{3}$ carbon?

    Hybridization is the mixing of atomic orbitals to form equivalent bonding orbitals. $sp$: linear, $180^\circ$ (e.g. alkynes). $sp^{2}$: trigonal planar, $120^\circ$ (e.g. alkenes). $sp^{3}$: tetrahedral, $109.5^\circ$ (e.g. alkanes).

11. Define radioactivity and name the three classical types of radioactive decay with their particle/charge.

    Radioactivity is the spontaneous emission of particles/energy from an unstable nucleus. Alpha ($\alpha$): $\ce{^4_2He}$ nucleus, $+2$ charge. Beta ($\beta^-$): high-energy electron, $-1$ charge. Gamma ($\gamma$): high-energy electromagnetic photon, no charge/mass.

12. State the radioactive decay law and define the decay constant.

    $$N = N_0 e^{-\lambda t}$$ where $N$ is the number of undecayed nuclei at time $t$, $N_0$ the initial number, and $\lambda$ the decay constant (probability of decay per unit time). Activity $A = \lambda N$.

13. How is half-life related to the decay constant?

    $$t_{1/2} = \frac{\ln 2}{\lambda} = \frac{0.693}{\lambda}.$$ After each half-life the quantity of radioactive material halves.

14. State the law of chemical equilibrium and write the equilibrium constant for $\ce{aA + bB <=> cC + dD}$.

    At equilibrium the ratio of product to reactant activities (each raised to its stoichiometric coefficient) is constant: $$K = \frac{[\ce{C}]^{c}[\ce{D}]^{d}}{[\ce{A}]^{a}[\ce{B}]^{b}}.$$

15. State Le Chatelier's principle.

    If a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the equilibrium shifts in the direction that partially counteracts (relieves) the imposed change.

16. Relate the standard Gibbs free energy change to the equilibrium constant.

    $$\Delta G^\circ = -RT \ln K.$$ A negative $\Delta G^\circ$ gives $K > 1$ (products favored); positive $\Delta G^\circ$ gives $K < 1$ (reactants favored).

17. State the relationship $\Delta G = \Delta H - T\Delta S$ and the condition for spontaneity.

    $$\Delta G = \Delta H - T\Delta S.$$ A process is spontaneous when $\Delta G < 0$, at equilibrium when $\Delta G = 0$, and non-spontaneous when $\Delta G > 0$ (at constant $T$, $P$).

18. State the Arrhenius equation for the temperature dependence of a rate constant.

    $$k = A \, e^{-E_a / RT}$$ where $A$ is the pre-exponential (frequency) factor, $E_a$ the activation energy, $R$ the gas constant and $T$ the absolute temperature.

19. Write the integrated rate laws and half-lives for zero-, first- and second-order reactions.

    Zero order: $[A] = [A]_0 - kt$, $t_{1/2} = \frac{[A]_0}{2k}$. First order: $\ln[A] = \ln[A]_0 - kt$, $t_{1/2} = \frac{0.693}{k}$. Second order: $\frac{1}{[A]} = \frac{1}{[A]_0} + kt$, $t_{1/2} = \frac{1}{k[A]_0}$.

20. Define BOD (Biochemical Oxygen Demand) and what it measures.

    BOD is the amount of dissolved oxygen consumed by microorganisms while biologically oxidizing organic matter in water over a specified period (typically 5 days at $20^\circ\text{C}$, giving $\text{BOD}_5$). It indicates biodegradable organic pollution, expressed in mg/L.

21. Define COD and how it compares with BOD.

    COD (Chemical Oxygen Demand) is the oxygen equivalent of organic matter oxidizable by a strong chemical oxidant (e.g. $\ce{K2Cr2O7}$). COD $\geq$ BOD always, because COD oxidizes both biodegradable and non-biodegradable matter; the ratio $\frac{\text{BOD}}{\text{COD}}$ indicates biodegradability.

22. State the first-order model for BOD exertion (the BOD curve).

    $$\text{BOD}_t = L_0 \left(1 - 10^{-k t}\right) \quad\text{or}\quad L_0\left(1 - e^{-k_1 t}\right)$$ where $L_0$ is the ultimate BOD and $k$ (or $k_1$) is the deoxygenation rate constant; $\text{BOD}_t$ is BOD exerted up to time $t$.

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