16-Civ-B5 Water Supply and Wastewater Treatment · December 2018
Nivaar worked solution (AI-drafted; not reviewed by a licensed engineer)
Paper format. National Examination, December 2018 — 16-Civ-B5 Water Supply and Wastewater Treatment. Three hours, closed book, one aid sheet written on both sides, and only an approved Casio or Sharp calculator. Question 1 is compulsory and candidates attempt any three of Questions 2–5; every question carries 25 marks. Marks are shown at the end of each question and the paper explicitly invites candidates to state any assumptions they make. All five questions are worked below, because the set is a study resource rather than a three-hour sitting.
Reference texts.
Question text not reproduced: the examination questions are © Engineers and Geoscientists BC. Open the official past paper (linked at the top of this page) to read the question, then follow the worked solution below.
Each of the five items is a measured parameter that a treatment engineer uses to decide something specific: whether a source can be used at all, what unit process must be added, or how an existing process must be operated. The answers below therefore follow the same pattern — what the parameter is, what governs it chemically, why it matters to public health or to the process, and the number that makes the point.
Sulfate ($\text{SO}_4^{2-}$) is the fully oxidised form of sulphur and is chemically stable and unreactive in an aerobic water. It is therefore not a health hazard at ordinary concentrations, and Health Canada sets an aesthetic objective of about 500 mg/L rather than a maximum acceptable concentration: above roughly 500 mg/L the water tastes bitter and acts as a laxative on people not acclimatised to it, and magnesium sulphate is the compound responsible. Sulfate is also the ion that scales reverse-osmosis membranes as calcium sulphate and that consumes lime uselessly in a softening plant, so it is a process parameter as much as a quality parameter.
The engineering significance of sulfate appears once the water turns anaerobic. Sulphate-reducing bacteria use $\text{SO}_4^{2-}$ as a terminal electron acceptor in the slime layer of a septic sewer, producing hydrogen sulphide; on a mass basis $96.1\ \text{g}$ of sulfate can yield $34.1\ \text{g}$ of $\text{H}_2\text{S}$, i.e. $0.355\ \text{g}$ per gram of sulfate, so a water at the 500 mg/L aesthetic objective carries the potential for roughly 177 mg/L of sulphide. The sulphide partitions to the sewer atmosphere, is re-oxidised to sulphuric acid on the moist crown by Thiobacillus, and destroys unlined concrete pipe — the classic crown-corrosion failure. A high-sulphate water supply is therefore a sewer-materials decision, not only a taste question.
Nitrate ($\text{NO}_3^-$) is the opposite case: it is aesthetically undetectable but has a genuine acute health effect. The Health Canada maximum acceptable concentration is 45 mg/L as nitrate, which is $45 \times 14.007/62.004 = 10.2\ \text{mg/L}$ expressed as nitrogen, and the accompanying nitrite MAC of 3 mg/L as $\text{NO}_2^-$ is 0.91 mg/L as N. Always state the basis — "as N" or "as $\text{NO}_3$" — because the two differ by a factor of 4.4 and mixing them up is the most common error in this topic. The hazard is methaemoglobinaemia: in the near-neutral gut of an infant under six months, nitrate is reduced to nitrite, which oxidises haemoglobin iron from Fe(II) to Fe(III) so that it can no longer carry oxygen. Because the mechanism is acute and the population at risk is infants, nitrate is not a parameter that can be averaged over a year.
Nitrate also matters as a tracer and as a nutrient. It is conservative and highly mobile in groundwater, so it is the standard indicator of agricultural or septic-field impact on a well field; and conventional treatment does not remove it — coagulation, filtration and disinfection leave it untouched, so control means ion exchange, reverse osmosis, biological denitrification or blending with a clean source. In a receiving water it is the nitrogen species that drives eutrophication in nitrogen-limited estuaries.
Total ammonia nitrogen (TAN) is the sum of the two forms in which ammonia exists in water, the ammonium ion and un-ionised ammonia, $\text{TAN} = \text{NH}_4^+\text{-N} + \text{NH}_3\text{-N}$, and it is what the standard laboratory method reports. Free (un-ionised) ammonia is the neutral $\text{NH}_3$ molecule alone. The split between them is set by an acid–base equilibrium and therefore by pH and temperature:
$$f_{\text{NH}_3} = \frac{1}{1 + 10^{\,\text{p}K_a - \text{pH}}}, \qquad \text{p}K_a = 0.09018 + \frac{2729.92}{T\,(\text{K})}$$At $25\,{}^{\circ}\text{C}$ this gives $\text{p}K_a = 9.246$, so the un-ionised fraction is 0.56 per cent at pH 7, 5.37 per cent at pH 8 and 36.2 per cent at pH 9 — very nearly a factor of ten for each pH unit. That sensitivity is the whole point of the distinction. It is the free ammonia that is toxic, because the neutral molecule crosses the gill membrane of a fish whereas the charged ammonium ion largely does not, and the CCME long-term guideline for the protection of aquatic life is therefore written on un-ionised ammonia, 0.019 mg/L as $\text{NH}_3$ ($0.0156\ \text{mg/L}$ as N). A secondary effluent carrying 25 mg/L of TAN discharged at pH 8 contains $25 \times 0.0537 = 1.34\ \text{mg/L}$ of free ammonia-N, about 86 times the guideline; the same effluent at pH 7 would be nine times less toxic. A permit written on TAN alone is thus not protective unless the receiving-water pH is also known.
Inside the plant, TAN is what sizes the aeration system and the alkalinity supply. Complete nitrification, $\text{NH}_4^+ + 2\,\text{O}_2 \rightarrow \text{NO}_3^- + 2\,\text{H}^+ + \text{H}_2\text{O}$, requires $2 \times 31.998/14.007 = 4.57\ \text{g}$ of oxygen and destroys $2 \times 50.04/14.007 = 7.14\ \text{g}$ of alkalinity as $\text{CaCO}_3$ for every gram of nitrogen oxidised. Free ammonia also matters to disinfection: it reacts with chlorine to form chloramines, so a plant that does not nitrify cannot achieve free-chlorine contact without breakpoint chlorination, and free ammonia above roughly 10 mg/L as N inhibits the nitrite oxidisers themselves and causes nitrite lock.
Alkalinity is the acid-neutralising capacity of the water — the quantity of strong acid required to titrate it to a defined end-point, conventionally reported in mg/L as $\text{CaCO}_3$ where $1\ \text{meq/L} = 50.04\ \text{mg/L}$. In municipal wastewater it is carried almost entirely by bicarbonate, with carbonate and hydroxide appearing only in industrial or lime-dosed streams. Alkalinity is not itself a pollutant; its significance is that it is the buffer that keeps the biological reactors at a workable pH, and it is consumed by the very reactions the plant is designed to run.
Given. A plant treats a sewage with an influent alkalinity of 250 mg/L as $\text{CaCO}_3$ and a TKN of 40 mg/L as N, of which 32 mg/L is actually oxidised (the balance is assimilated into cell mass). Find. Whether the buffer survives nitrification, and what an anoxic zone would recover.
Nitrification destroys $7.14\ \text{mg}$ of alkalinity per mg of N oxidised, so $32 \times 7.145 = 228.6\ \text{mg/L}$ as $\text{CaCO}_3$ disappears and only $250 - 228.6 = 21.4\ \text{mg/L}$ remains. That is well below the 50 to 70 mg/L needed to hold the mixed liquor above pH 6.8, and below pH 6.8 the nitrifiers slow sharply, so the reactor enters a destructive loop in which failing nitrification is diagnosed as a nitrogen problem when it is really a pH problem. Denitrification returns one equivalent of alkalinity per mole of nitrate reduced, $50.04/14.007 = 3.57\ \text{mg}$ as $\text{CaCO}_3$ per mg of N, so denitrifying 20 mg/L of nitrate-N recovers $20 \times 3.572 = 71.4\ \text{mg/L}$ and restores the residual to about 93 mg/L.
This is the quantitative argument for placing an anoxic zone ahead of the aerobic zone in a nitrifying plant: the pre-anoxic configuration recovers half the destroyed alkalinity for free and avoids a continuous lime or soda-ash chemical cost. Alkalinity is also monitored as an early-warning parameter in anaerobic digesters, where the volatile acid to alkalinity ratio rising above about 0.3 signals souring days before the gas production falls.
Chemical oxygen demand is the oxygen equivalent of the organic matter in a sample that can be oxidised by a strong chemical oxidant — in Standard Methods 5220, boiling acidic dichromate with a silver catalyst for two hours. Because the oxidant is aggressive and non-selective, COD captures essentially all organic carbon, biodegradable or not, and it is reported in mg/L of $\text{O}_2$. Its two operational virtues are speed and reproducibility: a result in two to three hours against five days for BOD, and no dependence on the health of a seed culture. For process control — which must respond within a shift — COD is the parameter that is actually usable.
The theoretical oxygen demand of a known compound follows directly from its oxidation half-reaction. For glucose, $\text{C}_6\text{H}_{12}\text{O}_6 + 6\,\text{O}_2 \rightarrow 6\,\text{CO}_2 + 6\,\text{H}_2\text{O}$, so $\text{ThOD} = 6 \times 31.998/180.16 = 1.066\ \text{g O}_2$ per gram of glucose, and a well-run COD test recovers better than 95 per cent of that. More usefully, the electron-equivalent bookkeeping behind COD is universal: one electron equivalent carries $32/4 = 8\ \text{g}$ of oxygen demand, which is why COD, BOD and methane potential can all be converted into one another.
The significance of COD lies mostly in its ratio to BOD. For a raw municipal sewage with $\text{COD} = 500$ and $\text{BOD}_5 = 220\ \text{mg/L}$, the ratio is 2.27; converting the five-day value to ultimate demand at the usual $\text{BOD}_5/\text{BOD}_u = 0.68$ gives $\text{BOD}_u = 323\ \text{mg/L}$, so the biodegradable fraction is $323/500 = 0.65$ and roughly 177 mg/L of the COD is inert. That inert fraction passes through the biology unchanged and sets the floor on the effluent COD no matter how the plant is operated. A COD/BOD ratio near 2 signals a readily treatable domestic waste; a ratio above about 4 signals a refractory or toxic industrial component, and is the usual trigger for source control, chemical pre-oxidation, or a decision not to accept the discharge at all.
Iron and manganese are nuisance metals rather than toxins at the concentrations found in Canadian groundwater. They are mobilised in the reduced forms $\text{Fe}^{2+}$ and $\text{Mn}^{2+}$, which are soluble and colourless, so a raw water drawn from an anoxic aquifer looks perfectly clear at the wellhead. On contact with air or with residual chlorine in the distribution system they oxidise to $\text{Fe(OH)}_3$ and $\text{MnO}_2$, which are insoluble and intensely coloured — red-brown and black respectively — and the customer receives the complaint, not the plant. Health Canada therefore sets aesthetic objectives of 0.3 mg/L for iron and 0.02 mg/L for manganese, with a health-based MAC of 0.12 mg/L for manganese; note that the manganese aesthetic objective is six times more stringent than its own health limit, because staining appears long before any toxicological concern.
The consequences are laundry and fixture staining, a metallic taste, and the growth of iron and manganese bacteria such as Gallionella and Crenothrix in the mains. Those organisms build tubercles that reduce the effective diameter and increase pumping head, harbour a biofilm that exerts a chlorine demand, and slough periodically to produce a "dirty water" event with no change at the treatment plant at all — which is why the parameter is a distribution-system issue as much as a treatment one.
Treatment is oxidation followed by solid–liquid separation, and the oxidant demand follows straight from stoichiometry. Per gram of metal, oxidation requires $0.143\ \text{g O}_2$ or $0.635\ \text{g Cl}_2$ or $0.94\ \text{g KMnO}_4$ for iron, and $0.291\ \text{g O}_2$ or $1.29\ \text{g Cl}_2$ or $1.92\ \text{g KMnO}_4$ for manganese; a groundwater with 2.5 mg/L Fe and 0.6 mg/L Mn therefore exerts a chlorine demand of $2.5 \times 0.635 + 0.6 \times 1.291 = 2.36\ \text{mg/L}$ before any disinfection credit is earned. The practical difficulty is that the two metals do not behave alike: iron oxidises by aeration alone within minutes at pH above about 7, whereas the oxidation of manganese by oxygen is impractically slow below pH 9.5, so manganese normally needs a stronger oxidant (permanganate, chlorine dioxide, ozone) or a catalytic manganese-dioxide-coated filter medium. Answering "aerate and filter" for both metals is the standard mistake.
