Water Chemistry

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CE 326: Principles of Environmental Engineering Water Chemistry Chemistry is one of the important components of water and wastewater treatment system Examples: Removal of pollutants by precipitation from aqueous phase such as phosphorus, hardness, sulfate, etc. Solubilization of chemical during water/wastewater treatment, e.g. lime, alum, etc. pH, and buffering system and carbonate system Alkalinity

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Precipitation and Solubility Precipitation: Dissolved ions can react with each other and form a solid compounds and settle down at the bottom of aqueous phase. This phase-change reaction of dissolved to solid state is called a precipitation. e.g. Ca2+ + CO32-  CaCO3 --- Hardness removal Ca2+ + PO42-  Ca3(PO4)2 --- Phosphorus removal Ba2+ + SO42-  BaSO4 --- Sulfate removal Solubility: Degree of dissolution of compounds in aqueous phase is known as solubility. For compounds/chemicals to be effective in water/wastewater treatment, they should be highly soluble. Some Compounds are more soluble and some are less. For example, Lime, AgCl are rarely soluble compounds whereas NaCl, NaHCO3 are highly soluble.

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Solubility Product The degree of the solubility of any compound can be represented by the following general reaction: AaBb (S)  aAb+ + bBa- Ca3(PO4)2  3Ca2+ + 2PO43- In above equation, the products of activity of both ions is always a constant for a given compound at a given temperature, i.e. [A+]a [B-]b = Ksp [Ca2+]3 [PO43-]2 = Ksp The constant Ksp is known as solubility product constant.

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What Does Solubility Product Mean?? When [A+]a [B-]b < Ksp, the compound continue to solubilize and the solubilization continue to take place until [A+]a [B-]b = Ksp 2. When [A+]a [B-]b > Ksp, the compound will start to precipitate Based on Ksp value, we can actually calculate the solubility of any particular compound e.g. for [Ba2+] [SO42-] = S2= Ksp = 1 x 10-10 Therefore, the solubility of BaSO4 is 1x10-5 M. whereas the solubility of CaF2is 1.96x10-4 M. So calcium fluoride is 20 times more soluble than barium sulfate. For [Ca2+] [F-]2 = (S) (2S)2 = Ksp = 3 x 10-11

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pH, and Buffering and Carbonate System pH is the term used universally to express the intensity of the acid or alkaline condition of aqueous phase. More precisely, acids are defined as those compounds that releases proton (H+) whereas bases are those compounds that accept protons (H+). Thus, pH is the measurement of hydrogen ion activity. Pure water dissociates to yield a conc. of hydrogen ions equal to 10-7 moles/L: H2O  H+ + OH- The water dissociation produces also one mole hydroxyl ion for each mole of hydrogen ion. Therefore, [H+] = [OH-] = 10-7 The equilibrium eq. for pure water is [H+] [OH-]= Kw =10-7x10-7 = 10-14 pH is usually expressed as log 1/[H+] or – log [H+] Thus, pH + POH = 14 with 7 being the neutral value. If we know Hydrogen ion concentration, we can calculate pH. What is the pH for [H+] = 10-4.8 ?

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0 7 14 Acid range Alkaline range Sulfuric acid Hydrochloric acid Nitric acid Sodium hydroxide Sodium bicarbonate Calcium hydroxide

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Buffering and Carbonate System An aqueous solution that resists changes in pH when acid or base is added is known to have sufficient buffering capacity. Atmospheric carbon dioxide produces a natural buffer and is the Most important buffer system in water and wastewater treatment CO2(g)  CO2 (aq) + H2O (1) dissolved CO2 CO2(aq) + H2O  H2CO3 (2) carbonic acid H2CO3  H+ + HCO3- (3) bicarbonate HCO3-  H+ + CO32- (4) carbonate CO2 in solution [CO2(aq)] is in equilibrium with atmospheric CO2(g). The change in any components in eqs. (4), (3), or (2) will cause CO2 either to be released or to dissolve into solution. If acid is added: H+ increases in the system. This derives eq. (4) towards left thereby forming more bicarbonate. The bicarbonate Further combines with H+ and forms carbonic acid [eq. (3)], which dissociates to CO2 and water. The excess carbon dioxide is then released to atmosphere.

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Buffering and Carbonate System CO2(g)  CO2 (aq) + H2O (1) dissolved CO2 CO2(aq) + H2O  H2CO3 (2) carbonic acid H2CO3  H+ + HCO3- (3) bicarbonate HCO3-  H+ + CO32- (4) carbonate If base is added: H+ is consumed increases in the system. This derives above eqs. towards right direction. This therefore converts the gaseous carbon dioxide to carbonate. If CO2 is bubbled into the system: The above reactions shift towards right due to formation of carbonic acids. This is because CO2 will combine with water. The pH will decrease. If CO2 is stripped out from the system: The above reactions shift towards left. CO2 is removed from the solution. The pH will increase

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Buffering and Carbonate System H2CO3  H+ + HCO3- pKa1 = 6.35 at 25o C HCO3-  H+ + CO32- pKa2 = 10.33 at 25oC

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Alkalinity Alkalinity is defined as the sum of all titratable bases down to about pH 4.5. Alkalinity thus measures its capacity to neutralize acid. In most aqueous system, the alkalinity contributions are from carbonate system, and any free H+ or OH-. Mathematically, alkalinity can be determined by: Alkalinity = [HCO3-] + 2[CO32-]+[OH-] –[H+] (5) Alkalinity provides a buffering capacity to aqueous system. The higher the alkalinity is, the higher the buffering capacity against pH changes. In above equation, the alkalinity is expressed as moles/L, however, in environmental engineering, alkalinity is commonly expressed as mg/L as CaCO3. mg/L as CaCO3 = (mg/L of species) {EW of CaCO3/EW of species} Equivalent weight (EW) = Molecular weight/number of charge See Example 4.8, page 207