18-Env-B9 Environmental Chemistry and Microbiology · May 2013
Nivaar worked solution (AI-drafted; not reviewed by a licensed engineer)
National Exams — May 2013 — 04-Env-B9, Environmental Chemistry/Microbiology. 3 hours duration; closed-book exam (approved Casio or Sharp calculator only). The paper has two sections — Section 1: Chemistry (11 questions, 50 marks) and Section 2: Microbiology (14 questions, 50 marks) — twenty-five questions constitute the complete exam and all are answered below. Total examination mark 100.
Reference texts. Davis & Cornwell, Introduction to Environmental Engineering (6th ed.) (water chemistry, disinfection, water/wastewater microbiology, indicator organisms); Metcalf & Eddy (Tchobanoglous, Stensel, Tsuchihashi & Burton), Wastewater Engineering: Treatment and Resource Recovery (5th ed.) (chemical unit processes, chemical phosphorus precipitation, biomass stoichiometry, activated-sludge microbiology); Guidelines for Canadian Drinking Water Quality (Health Canada); MWH's Water Treatment: Principles and Design (3rd ed.) (chlorine disinfection, contact-tank sizing).
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.
3.1 Ionization. The process by which a neutral molecule (or ionic compound) splits into, or generates, charged ions when dissolved — e.g. acids and bases ionizing in water, $\text{HCl}\to\text{H}^++\text{Cl}^-$ — producing a solution capable of conducting electricity.
3.2 Solubility product ($K_{sp}$). The equilibrium constant for the dissolution of a sparingly soluble ionic solid into its constituent ions, $\text{A}_a\text{B}_b(s)\rightleftharpoons a\text{A}^{n+}+b\text{B}^{m-}$, given by $K_{sp}=[\text{A}^{n+}]^a[\text{B}^{m-}]^b$ at saturation; used to predict precipitation (e.g. water hardness, metal-hydroxide removal).
3.3 Specific heat. The quantity of heat energy required to raise the temperature of a unit mass of a substance by one degree (J/(kg·K) or cal/(g·°C)); water's high specific heat (4.186 kJ/(kg·K)) is why large water bodies moderate local climate.
3.4 Specific gravity. The dimensionless ratio of a substance's density to the density of a reference substance (water at 4°C for liquids/solids, air for gases) at a stated temperature; used throughout water/wastewater engineering to characterize sludges, chemicals (e.g. liquid alum SG 1.2), and grit.
3.5 Generalized gas law. The combined ideal-gas relation $PV=nRT$, linking pressure, volume, moles and absolute temperature through the universal gas constant $R$; reduces to Boyle's, Charles's, and Avogadro's Laws as special cases and underlies gas-transfer and aeration calculations.
3.6 Charles's Law. At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature, $V/T=\text{constant}$, i.e. $V_1/T_1=V_2/T_2$.
3.7 Equivalent weight. The mass of a substance (in grams) that supplies or reacts with one mole of reference charge/H⁺ ions/electrons — equal to molecular weight divided by the number of replaceable H⁺, OH⁻, or the ionic charge/valence involved in the reaction; used to relate mass concentrations to normality.
3.8 Solute. The substance present in the smaller amount that is dissolved in a solvent to form a solution (e.g. dissolved salts in water).
3.9 Normality. A concentration unit expressing the number of gram-equivalent weights of solute per litre of solution ($N=$ equivalents/L $=M\times$ valence); used in titrimetric methods such as EDTA hardness or acid–base titrations because it accounts directly for reacting capacity.
3.10 Molecular weight. The sum of the atomic weights of all atoms in a molecule's chemical formula, expressed in g/mol (e.g. water, $\text{H}_2\text{O}=2(1)+16=18\ \text{g/mol}$); the basis for converting between mass and moles in every stoichiometric calculation.
