GATE Environmental Engineering Environmental Chemistry Syllabus

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.

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 $t$.

List key physical/chemical water quality parameters a student must know.

pH, turbidity, color, temperature, total dissolved solids (TDS), total suspended solids (TSS), alkalinity, hardness, dissolved oxygen (DO), BOD, COD, conductivity, chloride, nitrate, sulfate, fluoride, and total/fecal coliform (microbiological).

Define total hardness and how it is classified.

Hardness is the concentration of multivalent cations, chiefly $\ce{Ca^2+}$ and $\ce{Mg^2+}$, expressed as $\ce{CaCO3}$ equivalent (mg/L). Classification: soft $<75$, moderately hard $75\text{-}150$, hard $150\text{-}300$, very hard $>300$ mg/L as $\ce{CaCO3}$.

Differentiate carbonate (temporary) and non-carbonate (permanent) hardness.

Carbonate (temporary) hardness is due to bicarbonates/carbonates of $\ce{Ca}$ and $\ce{Mg}$ and is removed by boiling. Non-carbonate (permanent) hardness is due to sulfates, chlorides and nitrates of $\ce{Ca}$/$\ce{Mg}$ and is not removed by boiling (needs softening, e.g. lime-soda).

Define pH and pOH and state their relationship at $25^\circ\text{C}$.

$$\text{pH} = -\log[\ce{H+}], \quad \text{pOH} = -\log[\ce{OH-}].$$ At $25^\circ\text{C}$, $K_w = [\ce{H+}][\ce{OH-}] = 10^{-14}$, so $\text{pH} + \text{pOH} = 14$.

Distinguish strong and weak acids using the acid dissociation constant $K_a$.

A strong acid dissociates essentially completely ($K_a$ very large, e.g. $\ce{HCl}$). A weak acid only partially dissociates ($K_a$ small, e.g. $\ce{CH3COOH}$, $K_a \approx 1.8\times10^{-5}$). Lower $K_a$ (higher $pK_a$) means weaker acid.

Write the Henderson-Hasselbalch equation for a buffer.

$$\text{pH} = pK_a + \log\frac{[\text{A}^-]}{[\text{HA}]}$$ where $[\text{A}^-]$ is the conjugate base concentration and $[\text{HA}]$ the weak acid concentration.

What is a buffer solution and how is it typically composed?

A buffer resists pH change on addition of small amounts of acid or base. It is composed of a weak acid and its conjugate base (e.g. $\ce{CH3COOH/CH3COO-}$) or a weak base and its conjugate acid (e.g. $\ce{NH3/NH4+}$).

When is a buffer's capacity maximum, and what is the useful buffering range?

Buffer capacity is maximum when $[\text{HA}] = [\text{A}^-]$, i.e. $\text{pH} = pK_a$. The useful buffering range is approximately $\text{pH} = pK_a \pm 1$.

Write the carbonate system equilibria that control natural water pH.

$$\ce{CO2(aq) + H2O <=> H2CO3 <=> H+ + HCO3- <=> 2H+ + CO3^2-}$$ governed by $K_{a1} \approx 10^{-6.3}$ and $K_{a2} \approx 10^{-10.3}$ at $25^\circ\text{C}$.

Which carbonate species dominates at different pH ranges in natural water?

Below pH $\approx 6.3$: $\ce{H2CO3^*}$ (dissolved $\ce{CO2}$) dominates. Between pH $\approx 6.3$ and $10.3$: $\ce{HCO3-}$ dominates. Above pH $\approx 10.3$: $\ce{CO3^2-}$ dominates.

Define alkalinity and write its expression in terms of carbonate species.

Alkalinity is the acid-neutralizing capacity of water. $$\text{Alk} = [\ce{HCO3-}] + 2[\ce{CO3^2-}] + [\ce{OH-}] - [\ce{H+}]$$ usually expressed as mg/L of $\ce{CaCO3}$.

State Henry's law for the solubility of a gas in water.

$$C = K_H \, p_{\text{gas}}$$ The concentration of a dissolved gas is proportional to its partial pressure above the liquid; $K_H$ is the Henry's law constant. Solubility increases with partial pressure and decreases with temperature.

How do temperature and salinity affect the solubility of oxygen in water?

DO solubility decreases as temperature increases (gas solubility is exothermic) and decreases as salinity/dissolved-solids increase. Hence cold, fresh water holds more dissolved oxygen than warm, saline water.

Define a redox reaction and identify oxidation vs reduction.

A redox (oxidation-reduction) reaction involves transfer of electrons. Oxidation is loss of electrons (increase in oxidation state); reduction is gain of electrons (decrease in oxidation state). The oxidizing agent is reduced and the reducing agent is oxidized.

What is complexation (chelation) and give an example relevant to water chemistry.

Complexation is the formation of a coordinate bond between a central metal ion and ligands that donate electron pairs. Example: $\ce{EDTA}$ chelates $\ce{Ca^2+}$ and $\ce{Mg^2+}$, the basis of the EDTA titration for hardness.

Define solubility product $K_{sp}$ for a sparingly soluble salt $\ce{A_xB_y}$.

For $\ce{A_xB_y <=> x A^{m+} + y B^{n-}}$, $$K_{sp} = [\ce{A^{m+}}]^{x}[\ce{B^{n-}}]^{y}.$$ Precipitation occurs when the ion product exceeds $K_{sp}$; undersaturation when it is below.

What is the speciation of a contaminant and why does it matter?

Speciation is the distribution of an element among its chemical forms (oxidation states, complexes, free vs bound). It matters because toxicity, mobility, and bioavailability depend on form — e.g. $\ce{Cr^{6+}}$ (toxic, mobile) vs $\ce{Cr^{3+}}$ (less toxic), or $\ce{As^{3+}}$ vs $\ce{As^{5+}}$.

Give examples of inorganic and organic contaminants in water.

Inorganic: heavy metals ($\ce{Pb}$, $\ce{As}$, $\ce{Cd}$, $\ce{Cr}$, $\ce{Hg}$), nitrate, fluoride, cyanide. Organic: pesticides, phenols, polycyclic aromatic hydrocarbons (PAHs), trihalomethanes (THMs), volatile organic compounds (VOCs), and synthetic detergents.

What does soil/water organic matter consist of, and what are humic substances?

Organic matter is decomposing plant/animal residues plus living biomass. The stable fraction is humus, comprising humic substances: humic acid (soluble in base, not acid), fulvic acid (soluble at all pH), and humin (insoluble at all pH).

Describe the key steps of the nitrogen cycle: nitrification and denitrification.

Nitrification (aerobic, autotrophic): $\ce{NH4+ -> NO2- -> NO3-}$ by Nitrosomonas then Nitrobacter. Denitrification (anoxic): $\ce{NO3- -> NO2- -> N2}$ gas. Also fixation ($\ce{N2 -> NH3}$) and ammonification (organic N $\to \ce{NH4+}$).

What are the inorganic nitrogen species measured in water, and which form indicates recent vs old pollution?

Species: ammonia/ammonium ($\ce{NH3/NH4+}$), nitrite ($\ce{NO2-}$), nitrate ($\ce{NO3-}$), plus organic nitrogen. High ammonia/organic N indicates recent pollution; high nitrate indicates older, fully oxidized pollution. (Total Kjeldahl N = organic N + ammonia N.)

Why is phosphorus important in water quality and in what forms does it occur?

Phosphorus is a limiting nutrient causing eutrophication (algal blooms) when in excess. Forms: orthophosphate ($\ce{PO4^3-}$, $\ce{HPO4^2-}$, $\ce{H2PO4-}$), polyphosphates (condensed), and organic phosphorus. Orthophosphate is the most bioavailable form.

What is eutrophication and which nutrient typically limits it in freshwater?

Eutrophication is nutrient over-enrichment of a water body leading to excessive algal/plant growth, oxygen depletion on decay, and ecosystem degradation. Phosphorus is usually the limiting nutrient in freshwater (nitrogen often limits in marine systems).

What roles does potassium ($\ce{K+}$) play in soil and water, and how is it held in soil?

Potassium is an essential plant macronutrient (osmoregulation, enzyme activation). In soil it exists as solution $\ce{K+}$, exchangeable $\ce{K+}$ (held on exchange sites), fixed/non-exchangeable K (in clay interlayers), and mineral K. It is highly mobile and held by cation exchange.

Define Cation Exchange Capacity (CEC) and its units.

CEC is the total quantity of exchangeable cations ($\ce{Ca^2+}$, $\ce{Mg^2+}$, $\ce{K+}$, $\ce{Na+}$, $\ce{H+}$, $\ce{Al^3+}$) that a soil/colloid can adsorb and exchange, expressed in $\text{meq}/100\text{ g}$ (or $\text{cmol}_c/\text{kg}$) of soil.

What soil properties increase Cation Exchange Capacity?

Higher clay content, type of clay mineral (montmorillonite > illite > kaolinite), higher organic matter (humus) content, and higher pH (more negative charge from deprotonation) all increase CEC.

What is base saturation in the context of CEC?

Base saturation is the percentage of the CEC occupied by basic (non-acidic) cations $\ce{Ca^2+}$, $\ce{Mg^2+}$, $\ce{K+}$, $\ce{Na+}$, as opposed to acidic cations $\ce{H+}$ and $\ce{Al^3+}$: $$\% \text{ BS} = \frac{\text{base cations}}{\text{CEC}} \times 100.$$

Define the equivalence point of an acid-base titration and how it differs from the end point.

The equivalence point is where moles of acid equal moles of base (stoichiometric neutralization). The end point is the experimentally observed point where the indicator changes color; a well-chosen indicator makes the end point coincide closely with the equivalence point.

What is the common ion effect and its consequence for solubility?

The common ion effect is the suppression of dissociation (or solubility) of a weak electrolyte or sparingly soluble salt when a soluble salt sharing a common ion is added. It shifts equilibrium per Le Chatelier, decreasing solubility and lowering ion concentration of the dissolving species.