Water Analysis Training Manual PDF

FFC Laboratory Machhi Goth Water Analysis Training Manual Prepared By: M Aslam Khan AE-TS (LAB) Approved By: M Shoaib DM-TS (LAB) Rev. 2.0 Sep 2023 WATER ANALYSIS TRAINING MANUAL...

FFC Laboratory Machhi Goth Water Analysis Training Manual Prepared By: M Aslam Khan AE-TS (LAB) Approved By: M Shoaib DM-TS (LAB) Rev. 2.0 Sep 2023 WATER ANALYSIS TRAINING MANUAL TABLE OF CONTENTS WATER HARDNESS (TOTAL, Ca & Mg)....................................................................... 2 WATER ALKALINITY................................................................................................. 4 AMMONIA IN WATER.............................................................................................. 7 CHLORIDES IN WATER............................................................................................... 9 TOTAL IRON......................................................................................................... 11 ZINC BY EDTA..................................................................................................... 13 DETERMINATION OF SILICA..................................................................................... 14 DETERMINATION OF PHOSPHATES............................................................................ 16 OXYGEN SCAVENGER IN WATER............................................................................... 20 DETERMINATION OF UREA...................................................................................... 22 SULFATES (SO4)-2 IN WATER................................................................................... 24 NITRATES DETERMINATION..................................................................................... 26 NITRITES DETERMINATION...................................................................................... 28 FREE CHLORINE +ClO2 IN WATER............................................................................ 30 ALUMINUM IN WATER........................................................................................... 34 FREE MINERAL ACIDITY.......................................................................................... 37 CHEMICAL OXYGEN DEMAND.................................................................................. 39 TURBIDITY IN WATER............................................................................................. 42 TOTAL, DISSOLVED & SUSPENDED SOLIDS.................................................................. 44 OIL IN WATER...................................................................................................... 46 METHANOL IN WATER CONDENSATES....................................................................... 48 DEA IN BENFIELD/ MDEA IN WATER CONDENSATES................................................... 50 ORGANIC MATTER................................................................................................ 52 DISSOLVED OXYGEN IN BOILER FEED WATER.............................................................. 55 Rev. 2.0 Sep 2023 Page 1 of 56 WATER ANALYSIS TRAINING MANUAL WATER HARDNESS (TOTAL, Ca & Mg) WHAT IS WATER HARDNESS? Water which does not give ready and permanent lather with soap is called hard water. Presence of calcium and magnesium salts in the form of bicarbonate, chloride and sulfate in water makes water ‘hard’. Water free from soluble salts of calcium and magnesium is called soft water. It gives lather with soap easily. The property of water, which restricts the lather formation with soap, is called hardness. It is of two types: 1. TEMPORARY HARDNESS: It is due to the presence of magnesium and calcium bicarbonates [Ca(HCO3)2 and Mg(HCO3)2]. 2. PERMANENT HARDNESS: It is due to the presence of soluble salts of magnesium and calcium in the form of chlorides and sulfates in water (CaCl2, CaSO4, MgCl2 and MgSO4). DETERMINATION OF WATER HARDNESS The unit used for expressing the hardness of water is parts per million (ppm). It is the number of parts of calcium carbonate (CaCO3) equivalent hardness present in one million parts of water. TOTAL HADRNESS ERIOCHROME BLACK T/ SOLOCHROME BLACK T is a typical indicator. It contains three ionizable protons, so we will represent it by H3In. This indicator can be used for the titration of Mg2+ with EDTA. A small amount of indicator is added to the sample solution, and it forms a red complex with part of the Mg2+; the color of the uncomplexed indicator is blue. As soon as all the free Mg2+ is titrated, the EDTA displaces the indicator from the magnesium, causing a change in the color from red to blue: Mgln- + H2Y2- → MgY2- + Hln2- + H+ (Red) (Colorless) (Colorless) (Blue). The titration of calcium and magnesium with EDTA is done at pH 10, using an ammonia- ammonium chloride buffer. The pH must not be too high or the metal hydroxide may precipitate, causing the reaction with EDTA to be very slow. The metal-indicator complex must be less stable than the metal-EDTA complex, or else the EDTA will not displace it from the metal. On the other hand, it must not be too weak, or the EDTA will start replacing it at the beginning of the titration, and a diffuse end point will result. In general, the metal-indicator complex should be 10 to 100 times less stable than the metal-titrant complex. The formation constants of the EDTA complexes of calcium and magnesium are too close to differentiate between them in an EDTA titration, so they will titrate together, and the Eriochrome Black T end point can be used as above. This titration is used to determine total hardness of water. Rev. 2.0 Sep 2023 Page 2 of 56 WATER ANALYSIS TRAINING MANUAL CALCIUM HARDNESS CALCON/ SOLOCHROME DARK BLUE: An important application of the indicator is in the complexometric titration of calcium in the presence of magnesium; this must be carried out at a pH of about 12.3 in order to avoid the interference of magnesium. Under these conditions, magnesium is precipitated quantitatively as the hydroxide. The color change is from pink to pure blue. Eriochrome Black T cannot be used to indicate the direct titration of calcium in the absence of magnesium, with EDTA, however, because the indicator forms too weak a complex with calcium to give a sharp end point. Therefore, a small measured amount of Mg2+ is added to the Ca2+ solution; and as soon as the Ca2+ and the small amount of free Mg2+ are titrated. (The Ca2+ titrates first since its EDTA chelate is more stable.) A correction is made for the amount of EDTA used for titration of the Mg2+ by performing a "blank" titration on the same amount of Mg2+ added to the buffer. INTERFERENCES The test method is not suitable for highly colored waters, which obscure the color change of the indicator. As titration is carried out on high pH, i.e., pH=10 & 12, no other metals make stable complex with EDTA and hence do not interfere on this pH range. SUMMARY OF METHOD Total Hardness Sample 100 mL or aliquot up to 100 mL Method Complexometric Titration Reagents TH Buffer 02 mL Indicator Solochrome Black Red Color T (Solid) Titrate Against EDTA 0.01N End Point Blue PH 10 Ca+2 Hardness Sample 100 mL or aliquot up to 100 mL Method Complexometric Titration Reagents Ca. Buffer (NaOH 5% w/v) 02 mL Indicator Calcon (Solid) Red Color Titrate Against EDTA 0.01N End Point Purple PH 12 CALCULATIONS 𝐴 𝑜𝑟 𝐵 × 0.01𝑀 × 𝑀. 𝑊𝑡 𝑜𝑓 𝐶𝑎𝐶𝑂3 × 1000 𝑇𝑜𝑡𝑎𝑙 𝑜𝑟 𝐶𝑎 𝐻𝑎𝑟𝑑𝑛𝑒𝑠𝑠 (𝑎𝑠 𝑝𝑝𝑚 𝐶𝑎𝐶𝑂3 ) = 𝑆𝑎𝑚𝑝𝑙𝑒 𝑇𝑎𝑘𝑒𝑛 (𝑚𝑙) METHOD & REFERENCE As per WI-014-XX-407 WI Reference: ASTM D-1126, Standard Test Method for Hardness in Water Rev. 2.0 Sep 2023 Page 3 of 56 WATER ANALYSIS TRAINING MANUAL WATER ALKALINITY ALKALINITY IS THE WATER’S CAPACITY TO RESIST CHANGE IN pH This capacity is commonly known as "buffering capacity." For example, if you add the same weak acid solution to two vials of water - both with a pH of 7, but one with no buffering power (e.g. zero alkalinity) and the other with buffering power (e.g. an alkalinity of 50 mg/l), - the pH of the zero alkalinity water will immediately drop while the pH of the buffered water will change very little or not at all. The pH of the buffered solution would change when the buffering capacity of the solution is overloaded. Alkalinity refers to the capability of water to neutralize acid. This is really an expression of buffering capacity. A buffer is a solution to which an acid can be added without changing the concentration of available H+ ions (without changing the pH) appreciably. It essentially absorbs the excess H+ ions and protects the water body from fluctuations in pH. In most natural water bodies, the buffering system is carbonate-bicarbonate (H2CO3, HCO3-1 and CO3-2). The main sources for natural alkalinity are rocks, which contain carbonate, bicarbonate, and hydroxide compounds. Borates, silicates, and phosphates also may contribute to alkalinity. Limestone is rich in carbonates, so waters flowing through limestone regions or bedrock- containing carbonates generally have high alkalinity - hence good buffering capacity. Alkalinity is often related to hardness because the main source of alkalinity is usually from carbonate rocks (limestone) which are mostly CaCO3. If CaCO3 actually accounts for most of the alkalinity, hardness in CaCO3 is equal to alkalinity. Since hard water contains metal carbonates (mostly CaCO3), it is high in alkalinity. Conversely, unless carbonate is associated with sodium or potassium which don't contribute to hardness, soft water usually has low alkalinity and little buffering capacity. Therefore, generally, soft water is much more susceptible to fluctuations in pH from acid rains or acid contamination. DETERMINATION OF WATER ALKALINITY The sample is titrated with sulfuric acid (0.02N) to a colorimetric end point corresponding to a specific pH. P (Phenolphthalein) alkalinity is determined by titration to a pH of 8.3, as evidenced by the color change of phenolphthalein indicator, and indicates the total hydroxide and one-half the carbonate present. M (methyl orange) or T (total) alkalinity is determined by titration to a pH between 3.7 and 4.3, and includes all carbonate, bicarbonate and hydroxide. REACTIONS STEP 1 𝑃ℎ𝑒𝑛𝑜𝑙𝑜𝑡ℎ𝑎𝑙𝑒𝑖𝑛 (𝑃 𝑉𝑎𝑙𝑢𝑒) 𝑂𝐻 −1 + 𝐻𝐶𝑂3 −1 ∗ + 𝐶𝑂3 −1 + 𝐻 +1 → 𝐻2 𝑂 + 2 𝐻𝐶𝑂3 −1.. 𝑝𝐻 = 8.3 STEP 2 𝑀𝑒𝑡ℎ𝑦𝑙 (𝑀 𝑉𝑎𝑙𝑢𝑒) 𝐻𝐶𝑂3 −1 ∗ + 𝐻𝐶𝑂3 −1 + 𝐻 +1 → 𝐻2 𝐶𝑂3 ↔ 𝐻2 𝑂 + 𝐶𝑂2 … 𝑝𝐻 = 4.3/3.7 Rev. 2.0 Sep 2023 Page 4 of 56  WATER ANALYSIS TRAINING MANUAL NOTE: In step 2 bicarbonates of step 1(*) and already present in system will be titrated. INTERFERENCES Highly colored or turbid samples may mask the color change at the end point. Use a pH meter for these samples. SUMMARY OF METHOD P-Value Sample 100 mL or aliquot up to 100 mL Method Acid Base Titration Reagents Indicator Phenolphthalein Pink Color Titrate Against H2SO4 (0.02N) End Point Colorless PH 8.3 At end point * M-Value Sample Continue with * Method Acid Base Titration Reagents Indicator Mixed Indicator Greenish Color Titrate Against H2SO4 (0.02N) End Point Yellowish Brown PH 3.7 At end point INTERPRETATION OF RESULTS There are 5 possible cases of P & M values and from them we can predict the presence or absence of any specie in sample. P=0 P=M 2P>M 2P 12.0 Rev. 2.0 Sep 2023 Page 7 of 56 WATER ANALYSIS TRAINING MANUAL CALCULATIONS 𝑁𝐻3 (𝑝𝑝𝑚 ) = 𝐴𝑏𝑠 (𝑆𝑎𝑚𝑝𝑙𝑒 − 𝐵𝑙𝑎𝑛𝑘) × 𝑆𝑙𝑜𝑝𝑒 × 𝐷𝑖𝑙𝑢𝑡𝑖𝑜𝑛 𝐹𝑎𝑐𝑡𝑜𝑟 (𝑖𝑓 𝑎𝑛𝑦) METHOD & REFERENCE As per WI-003-XX-407 WI Reference: ASTM D-1426, Standard Test Method for Ammonia Nitrogen in Water, Method – A (Direct Nesslerization) Rev. 2.0 Sep 2023 Page 8 of 56 WATER ANALYSIS TRAINING MANUAL CHLORIDES IN WATER THEORY Chlorides in water are estimated by precipitation titration. The indicator theory for these titrations is different from that for acid-base indicators. The properties of the indicators do not necessarily depend on the concentration of some ion in solution, that is, on Cl- or Ag+. The indicator type forms a colored compound with the titrant when the titrant is in excess. The chloride is titrated with standard silver nitrate solution. A soluble chromate salt is added as the indicator. This produces a yellow solution. When the precipitation of the chloride is complete, the first excess of Ag+ reacts with the indicator to precipitate red silver chromate: 𝐶𝑟𝑂4−2 + 2𝐴𝑔+ → 𝐴𝑔2 𝐶𝑟𝑂4 (𝑌𝑒𝑙𝑙𝑜𝑤) (𝑅𝑒𝑑) The concentration of the indicator is important. The Ag2Cr04 should just start precipitating at the equivalence point, where we have a saturated solution of AgCl. From Ksp, the concentration of Ag+ at the equivalence point is 10-5 M. (It is less than this before the equivalence point.) Therefore, Ag2Cr04 should precipitate just when [Ag+] = 10-5 M. The solubility product of Ag2Cr04 is [1.1 X 10-12], by inserting the Ag+ concentration in the Ksp equation for Ag2Cr04, we calculate that, for this to occur, [CrO4-2] should be 0.011 M: [𝐴𝑔]2 [𝐶𝑟𝑂4 ] = 1.1 × 10−12 𝑀 [10−5 ]2 [𝐶𝑟𝑂4 ] = 1.1 × 10−12 𝑀 [𝑪𝒓𝑶𝟒 ] = 𝟏. 𝟏 × 𝟏𝟎−𝟐 𝑴 If the concentration is greater than this, Ag2Cr04 will begin to precipitate when [Ag+] is less than 10-5 M (before the equivalence point). If it is less than 0.011 M, then the [Ag+] will have to exceed 10-5 M (beyond the equivalence point) before precipitation of Ag2Cr04 begins. In actual practice, the indicator concentration is kept at 0.002 to 0.005 M. If it is much higher than this, the intense yellow color of the chromate ion obscures the red Ag2Cr04 precipitate color, and an excess of Ag+ is required to produce enough precipitate to be seen. An indicator blank should always be run and subtracted from the titration to correct for errors. THE pH The titration is performed in slightly alkaline solution, at a pH of about 8. If the solution is too acid (pH < 6), then part of the indicator is present as HCr04 -, and more Ag+ will be required to form the Ag2Cr04 precipitate. Above pH 8, silver hydroxide may be precipitated (at pH> 10) to give a pH about 8. The titration is useful for determining chloride in neutral or un-buffered solutions, such as drinking water. INTERFERENCES The anions and cations generally found in water offer no interference. Zinc, lead, nickel, and ferrous and chromous ions affect solution and end-point colors, but do not reduce the Rev. 2.0 Sep 2023 Page 9 of 56 WATER ANALYSIS TRAINING MANUAL accuracy of the titration when present in concentrations up to 100 mg/L. Copper is tolerable up to 50 mg/L. Ferric ion above 10 mg/L must be reduced before titration, and sulfite ion must be oxidized. Bromide and fluoride will be partially titrated with the chloride. Quaternary ammonium salts also interfere if present in significant amounts (1 to 2 mg/L). Deep color may also interfere. SUMMARY OF METHOD - Cl 100 mL or aliquot up to 100 mL Method Precipitation Titration Reagents Indicator K2CrO4 (5% w/v) Yellow Color Titrate against AgNO3 (0.1N) End Point Red PH Neutral- Slight Alkaline (8) CALCULATIONS 𝐴 × 0.1𝑁 × 𝐸𝑞. 𝑊𝑡. 𝑜𝑓 𝐶𝑙 (35.45) × 1000 𝐶𝑙 − (𝑎𝑠 𝑝𝑝𝑚) = 𝑆𝑎𝑚𝑝𝑙𝑒 𝑇𝑎𝑘𝑒𝑛 (𝑚𝐿) METHOD & REFERENCE As per WI-010-XX-407 WI Reference: ASTM D-512, Standard Test Method for Chloride Ion in Water Rev. 2.0 Sep 2023 Page 10 of 56 WATER ANALYSIS TRAINING MANUAL TOTAL IRON 1,1O-PHENANTHROLINE METHOD: Iron (II) reacts with 1,10-phenanthroline to form an orange-red complex [(C12H8N2)Fe]+2. The color intensity is independent of the acidity in the pH range 2-9, and is stable for long periods. Iron (III) may be reduced with hydroxylammonium chloride. REACTIONS 𝐹𝑒 +3 + 2𝑁𝐻4 𝑂𝐻 → 𝐹𝑒 +2 + 2𝑁2 + 4𝐻2 𝑂 𝑝𝐻 3−6 𝐹𝑒 +2 + 3 → Red 2,2’-BIPYRIDYL METHOD 2,2’-Bipyridyl reacts with Fe+2 the same way as O-Phenanthrolein. It has also two lone pairs of electrons on Nitrogen and a bidentate ligand. The reaction is also more like earlier method. 2,2’-Bipyridyl THE IRON BUFFER The complex is stable at pH about 4.7 but the color intensity is independent of the pH range over 2-9 and is stable for long period. For the specifically iron complex with O-phenanthrolein, pH must be adjusted to slightly acidic 3-6 by the acetate buffer. Acidic media also inhibits the precipitation of Iron as hydroxide. That is why samples are acidified before analysis to dissolve the iron and to lower the pH. INTERFERENCES The metal ions other than ferrous iron, which can form a complex with O-phenanthroline, are manganese, cadmium, copper, zinc, cobalt, nickel, chromium, and ruthenium. The complexation/extraction is carried out at pH 4.0 to 4.5 in the presence of excess O-phenanthroline to achieve maximum color development with Fe(II) and also to eliminate interferences of competing ions. In acid solution, all competing ions form colorless complexes except ruthenium and cobalt, which are yellow, and none except the colorless copper complex are extractable into an organic solvent. In a natural water sample buffered at pH 4, cuprous copper is the only metal ion that could potentially affect the measurement of ferrous iron; both species compete for the Rev. 2.0 Sep 2023 Page 11 of 56 WATER ANALYSIS TRAINING MANUAL complexing agent. However, excess O-phenanthroline is present to complex both the ferrous iron and cuprous copper in the sample. SUMMARY OF METHOD Total Iron Sample Aliquot in 100 mL flask Method Colorimetric Reagents HCl (1:1) 1 mL Hydroxylamine HCl (10% Aq.) 2 mL Boil & Cool Iron Buffer (pH 4.7) 5 mL O-Phenanthroline (0.1% in 10% 4 mL Ethanol) NH4OH (1:1) 3 mL Make up 100 mL Color Wait 5 Min @ 510nm in 10/ Against RB Development 50mm pH 3-6 (5.4) CALCULATIONS 𝐴𝑏𝑠 × 𝑆𝑙𝑜𝑝𝑒 × 100 𝐼𝑟𝑜𝑛 𝑖𝑛 𝐵𝑒𝑛𝑓𝑖𝑒𝑙𝑑 (𝑎𝑠 𝑝𝑝𝑚) = 𝑆𝑎𝑚𝑝𝑙𝑒 𝑇𝑎𝑘𝑒𝑛 (𝑔) METHOD & REFERENCE As per WI-016-XX-407 WI Reference: ASTM D-1068, Standard Test Methods for Iron in Water, Method – A Rev. 2.0 Sep 2023 Page 12 of 56 WATER ANALYSIS TRAINING MANUAL ZINC BY EDTA Zinc can be determined in water samples in the presence of Mg with EBT indicator. In order to determine zinc is first masked by adding sodium cyanide and tetracyano sodium zincate Na2[Zn(CN)4] is formed. After all free Mg+2 is titrated zinc is damasked by formaldehyde acetic acid at lower pH (4-5) and titrated against EDTA. For direct determination of zinc in water sample, Xylenol indicator is used. It retains acid base properties in cresol red and displays metal indicator properties even in acidic solution (3- 5pH). Acidic solutions of indicator are lemon yellow and of metal complexes are intense red. Direct titration with EDTA and color change is sharp. HMTA BUFFER HMTA (Hexamethylenetetraamine) acts as buffer and maintain the pH at 5-5.5, which is optimum for metal indicator complex to show color in this range. XYLENOL ORANGE Metal indicator having red color complex with zinc at pH 5 and with EDTA complex color is red. INTERFERENCES Iron (III) can cause interference in determination of zinc in water samples. Interference of trace level analysis is not tolerate able hence a reducing agent (Hydroxylamine HCl) is added to reduce the Iron (III) into Iron (II) which don’t cause interference. SUMMARY OF METHOD Total Iron Sample Aliquot in 100 mL flask Method Complexoetric Titration Reagents Indicator Xylenol Orange (0.2% Aq.) Red Color Titrate against EDTA (0.0025M) End Point Yellow PH Slight Acidic (5-6) CALCULATIONS 𝐴 × 0.0025𝑀 × 𝐸𝑞. 𝑊𝑡 𝑜𝑓 𝑍𝑛(65.4) × 1000 𝑍𝑛 (𝑝𝑝𝑚) = 𝑆𝑎𝑚𝑝𝑙𝑒 𝑇𝑎𝑘𝑒𝑛 (𝑚𝑙) METHOD & REFERENCE As per WI-031-XX-407 WI Reference: Vogel’s Text Book of Quantitative Chemical Analysis, 5th Ed. Chapter X: Metal Ion Indicators Rev. 2.0 Sep 2023 Page 13 of 56 WATER ANALYSIS TRAINING MANUAL DETERMINATION OF SILICA SILICA ORIGIN Silicon comprises about 28 % of the lithosphere and is, next to oxygen, the most abundant element. It is found as the oxide in crystalline forms, as in quartz; combined with other oxides and metals in a variety of silicates; and in amorphous forms. Silica is only slightly soluble in water. The presence of most silica in natural waters comes from the gradual degradation of silica-containing minerals. The type and composition of the silica-containing minerals in contact with the water and the pH of the water are the primary factors controlling both the solubility and the form of silica in the resulting solution. Silica may exist in suspended particles, as a colloid, or in solution. It may be monomeric or polymeric. In solution, it can exist as silicic acid or silicate ion, depending upon pH. METHOD OF DETERMINATION This colorimetric test method is based on the reaction of the soluble silica with molybdate ion to form a greenish-yellow complex, which in turn is converted to a blue complex by reduction with reducing agent (Ascorbic Acid). Small quantities of dissolved silicic acid react with the solution of Molybdate in an acid medium to give an intense yellow colored complex due to the probability of complex Molybdosilicic acid H4[SiMo12O40]. This can be applied to colorimetric determination of silicates at 400nm. It is usually better to reduce the metal complex to blue; a solution of mixture of 1-amino-2-naphthol-4-sulphonic acid and sodium hydrogen sulfite is a satisfactory reducing agent. Solution of Ascorbic Acid can also be employed as reducing agent for the complex. This test method is useful for concentrations as low as 20 μg/L. NOTE: Batch to batch variations in ammonium molybdate have been found to affect results at low concentrations (below 0.1 mg/L). High blanks, nonlinear calibration curves, and poor reproducibility have been observed with some batches of this compound. When working with low concentrations of silica, a batch of ammonium molybdate known to produce reasonable blanks, linearity, and reproducibility should be set aside for this purpose. WHY WE MEASURE SILICA? Silica concentration is an important consideration in some industrial installations such as steam generation and cooling water systems. Under certain conditions, silica forms troublesome silica and silicate scales, particularly on high-pressure steam turbine blades. In cooling water systems, silica forms deposits when solubility limits are exceeded. In contrast, silica may be added as a treatment chemical in some systems, for example, in corrosion control. Silica removal is commonly accomplished by ion exchange, distillation, reverse osmosis, or by precipitation, usually with magnesium compounds in a hot or cold lime softening process. Colloidal silica is removed in clarifier and major portion of soluble silica is removed in strong cationic resin V 908. INTERFERENCES Phosphates, Arsenate and Germinates give similar color & either, they must be removes or their interferences must be eliminate by the addition of suitable reagent, Arsenic and Germinates can be removed by evaporation with HCl. Phosphates are masked by the addition of Rev. 2.0 Sep 2023 Page 14 of 56 WATER ANALYSIS TRAINING MANUAL Oxalic Acid. Colorimetric method that determines molybdate reactive silica. It is applicable to most waters, but some waters may require filtration and dilution to remove interferences from color and turbidity. A high dissolved salts concentration, such as in seawater or brine samples, can affect color development. This can be compensated for by preparing standards in a matrix similar to that of samples or by using a standard additions technique. Strong oxidizing and reducing agents that may be found in some industrial wastewaters may interfere in the reduction step of the reaction. Such wastewaters may also contain organic compounds that may interfere in the color formation. SUMMARY OF METHOD Silica Sample 100 mL or aliquot up to 100 mL in plastic beaker Reagents HCl (1:4) 01 mL Amm. Molybdate (7.5% 02 mL Wait 5 Min Aq.) Oxalic Acid (10% Aq.) 02 mL Wait 5 Min Ascorbic Acid (1% Aq.) 02 mL Color Development Wait 5 Min @ 650/ 810nm in 10/ Against RB 50mm pH 1.1-1.3 CALCULATIONS 𝑆𝑖𝑂2 (𝑝𝑝𝑏/ 𝑝𝑝𝑚 ) = 𝐴𝑏𝑠 × 𝑆𝑙𝑜𝑝𝑒 × 𝐷𝑖𝑙𝑢𝑡𝑖𝑜𝑛 𝐹𝑎𝑐𝑡𝑜𝑟 (𝑖𝑓 𝑎𝑛𝑦) METHOD & REFERENCE As per WI-026-XX-407 WI Reference: ASTM D-859, Standard Test Method for Silica in Water Rev. 2.0 Sep 2023 Page 15 of 56 WATER ANALYSIS TRAINING MANUAL DETERMINATION OF PHOSPHATES TYPES OF PHOSPHATES Phosphorus occurs in natural waters and in wastewaters almost solely as phosphates. These are classified as orthophosphates, condensed phosphates (pyro-, meta-, and other polyphosphates), and organically bound phosphates. They occur in solution, in particles or detritus, or in the bodies of aquatic organisms. These forms of phosphate arise from a variety of sources. Small amounts of orthophosphate or certain condensed phosphates are added to some water supplies during treatment. Phosphates are used extensively in the treatment of boiler waters. Orthophosphates applied to agricultural or residential cultivated land as fertilizers are carried into surface waters with storm runoff and to a lesser extent with melting snow. Organic phosphates are formed primarily by biological processes. They are contributed to sewage by body wastes and food residues, and also may be formed from orthophosphates in biological treatment processes or by receiving water biota. Phosphorus analyses embody two general procedural steps: (a) conversion of the phosphorus form of interest to dissolved orthophosphate, and (b) colorimetric determination of dissolved orthophosphate. Phosphates that respond to colorimetric tests without preliminary hydrolysis or oxidative digestion of the sample are termed ‘‘reactive phosphorus.’’ While reactive phosphorus is largely a measure of orthophosphate, a small fraction of any condensed phosphate present usually is hydrolyzed unavoidably in the procedure. Reactive phosphorus occurs in both dissolved and suspended forms. Acid hydrolysis at boiling-water temperature converts dissolved and particulate condensed phosphates to dissolved orthophosphate. The hydrolysis unavoidably releases some phosphate from organic compounds, but this may be reduced to a minimum by judicious selection of acid strength and hydrolysis time and temperature. The term ‘‘acid-hydrolysable phosphorus’’ is preferred over ‘‘condensed phosphate’’ for this fraction. The phosphate fractions that are converted to orthophosphate only by oxidation destruction of the organic matter present are considered ‘‘organic’’ or ‘‘organically bound’’ phosphorus. The severity of the oxidation required for this conversion depends on the form—and to some extent on the amount—of the organic phosphorus present. Like reactive phosphorus and acid-hydrolysable phosphorus, organic phosphorus occurs both in the dissolved and suspended fractions. The total phosphorus as well as the dissolved and suspended phosphorus fractions each may be divided analytically into the three chemical types that have been described: reactive, acid- hydrolysable, and organic phosphorus. Rev. 2.0 Sep 2023 Page 16 of 56 WATER ANALYSIS TRAINING MANUAL PHOSPHATE DOSING IN BOILERS Phosphate is added in form of Na3PO4 in boilers feed drum. There are two reasons for this treatment. 1. To maintain alkaline pH (9.5-10) to avoid acid corrosion. 2. Ortho acts as anodic inhibitor while poly (organic) after combining with calcium acts as cathodic inhibitor. Na3 𝑃𝑂4 + 3𝐶𝑎𝑆𝑂4 → 3𝑁𝑎𝑆𝑂4 + 𝐶𝑎3 (𝑃𝑂4 )2 ↓ Na3 𝑃𝑂4 + 3𝑀𝑔𝑆𝑂4 → 3𝑁𝑎𝑆𝑂4 + 𝑀𝑔3 (𝑃𝑂4 )2 ↓ Calcium and Magnesium Phosphate is removed in boiler blow down operation as soft scale. Otherwise, they can combine with silica and form silicates (hard scale) in steam lines, which has the insulation effect also along with scale deposit. SAMPLING & STORAGE If dissolved phosphorus forms are to be differentiated, filter sample immediately after collection. Preserve by freezing at or below - 10°C. In some cases, 40 mg HgCl2/L may be added to the samples, especially when they are to be stored for long periods before analysis. CAUTION: HgCl2 is a hazardous substance; take appropriate precautions in disposal; use of HgCl2 is not encouraged. Do not add either acid or CHCl3 as a preservative when phosphorus forms are to be determined. If total phosphorus alone is to be determined, add H2SO4 or HCl to pH - 2 and cool to 4°C, or freeze without any additions. Do not store samples containing low concentrations of phosphorus in plastic bottles unless kept in a frozen state because phosphates may be adsorbed onto the walls of plastic bottles. Rinse all glass containers with hot dilute HCl, then rinse several times in reagent water. Never use commercial detergents containing phosphate for cleaning glassware used in phosphate analysis. More strenuous cleaning techniques may be used. METHODS OF DETERMINATION High Range In a dilute orthophosphate solution, ammonium molybdate reacts under acid conditions to form a heteropoly acid, molybdophosphoric acid. In the presence of vanadium, yellow vanadomolybdophosphoric acid is formed. The intensity of the yellow color is proportional to phosphate concentration. The minimum detectable concentration is 200 μg P/L in 10mm spectrophotometer cell. Low Range Molybdophosphoric acid is formed and reduced by stannous chloride to intensely colored molybdenum blue. This method is more sensitive than Method C and makes feasible measurements down to 7 μg P/L by use of increased light path length. Below 100 μg P/L an extraction step may increase reliability and lessen interference. The minimum detectable concentration is about 3 μg P/L. The sensitivity at 0.3010 absorbance is about 10 μg P/L for an absorbance change of 0.009. Rev. 2.0 Sep 2023 Page 17 of 56 WATER ANALYSIS TRAINING MANUAL INTERFERENCES Positive interference is caused by silica and arsenate only if the sample is heated. Negative interferences are caused by arsenate, fluoride, thorium, bismuth, sulfide, thiosulfate, thiocyanate, or excess molybdate. Blue color is caused by ferrous iron but this does not affect results if ferrous iron concentration is less than 100 mg/L. Sulfide interference may be removed by oxidation with bromine water. Ions that do not interfere in concentrations up to 1000 mg/L are Al3+, Fe3+, Mg2+, Ca2+, Ba2+, Sr2+, Li+, Na+, K+, NH4+, Cd2+, Mn2+, Pb2+, Hg+, Hg2+, Sn2+, Cu2+, Ni2+, Ag+, U4+, Zr4+, AsO3-, Br-, CO32-, ClO4-, CN-, IO3-, SiO44- NO3-, NO2-, SO42-, SO32-, pyrophosphate, molybdate, tetraborate, selenate, benzoate, citrate, oxalate, lactate, tartrate, formate, and salicylate. If HNO3 is used in the test, Cl- interferes at 75 mg/L. SUMMARY OF METHOD O-(PO4)-3 Sample 25 mL in 50 mL V. Flask × 2 (S & Sb) For Sample In case of C.W. also prepare Sample Blank For Sample Blank (Sb) Vanadate-Molybdate 10 mL HCl (1:3) 10 mL Reagent Make up volume Color Development Wait 10 Min @ 400 in 10 mm Against RB pH 1.0 – 2.0 Total-(PO4)-3 50 mL in 250 mL Titration Flask K2S2O8 0.5 g H2SO4 (30%) 1 mL Water 50 mL Place of hot plate and boil till 10 mL solution is left Cool and transfer the contents in 50 mL V. flask Phenolphthalein Few drops NaOH (10%) ~2 mL Neutralize with H2SO4 (1N) till colorless Transfer to 50 mL V. flask and Make up V. Split into 25 mL (× 2) 50 mL V. Flasks (S & Sb) For S For Sb HCl (1:3) 10 mL Vanadate-Molybdate 10 mL Reagent Color Development Wait 10 Min @ 400 in 10 mm Against RB pH 1.0 – 2.0 CALCULATIONS 𝑂𝑟𝑡ℎ𝑜 𝑃𝑂4 (𝑝𝑝𝑚) = 𝐴𝑏𝑠 (𝑆 − 𝑆𝑏) × 𝑆𝑙𝑜𝑝𝑒 × 𝐷𝑖𝑙𝑢𝑡𝑖𝑜𝑛 𝐹𝑎𝑐𝑡𝑜𝑟 (𝑖𝑓 𝑎𝑛𝑦) 𝑇𝑜𝑡𝑎𝑙𝑃𝑂4 (𝑝𝑝𝑚) = 𝐴𝑏𝑠 (𝑆 − 𝑆𝑏) × 𝑆𝑙𝑜𝑝𝑒 × 𝐷𝑖𝑙𝑢𝑡𝑖𝑜𝑛 𝐹𝑎𝑐𝑡𝑜𝑟 (𝑖𝑓 𝑎𝑛𝑦) 𝑂𝑟𝑔𝑎𝑛𝑖𝑐 𝑃𝑂4 (𝑝𝑝𝑚 ) = (𝑇𝑜𝑡𝑎𝑙 𝑃𝑂4 ) − (𝑂𝑟𝑡ℎ𝑜 𝑃𝑂4 ) METHOD & REFERENCE As per WI-025-XX-407 WI Reference: Standard Methods of Water & Waste Water; Phosphorous 4500-P C (High Range) Rev. 2.0 Sep 2023 Page 18 of 56 WATER ANALYSIS TRAINING MANUAL Standard Methods of Water & Waste Water; Phosphorous 4500-P D (Low range) Rev. 2.0 Sep 2023 Page 19 of 56 WATER ANALYSIS TRAINING MANUAL OXYGEN SCAVENGER IN WATER OXYGEN SCAVENGER DOSING Hydrazine is historically the more commonly used oxygen scavenger for higher pressure boilers. However, due to its toxicity and related legislation, its use has been replaced with alternatives. Elimin-Ox (Carbohydrazide) is a volatile oxygen scavenger which reacts relatively slow with oxygen. The reaction product of carbohydrazide is carbon dioxide, a gas that dissolves in water as carbonic acid, lowering the pH value of the steam and condensate. (𝑁𝐻2 𝑁𝐻)2 𝐶𝑂 + 2𝑂2 → 𝐶𝑂2 + 2𝑁2 + 3𝐻2 𝑂 Carbonic acid can be corrosive to the steam and condensate part of the boiler and addition of condensate treatment chemicals or pH adjustment with ammonia is required. Some studies on high pressure steam system show the amount of formed carbon dioxide is limited and does not affect the cycle chemistry. At higher temperatures carbohydrazide decomposes to hydrazine as a result carbohydrazide cannot be used for systems that are applied for food. The formed hydrazine reacts with oxygen so there is no effect on oxygen scavenging properties when carbohydrazide degrades. 𝐻𝑖𝑔ℎ𝑒𝑟 𝑇𝑒𝑚𝑝𝑒𝑟𝑎𝑡𝑢𝑟𝑒𝑠 (>180℃) (𝑁𝐻2 𝑁𝐻)2 𝐶𝑂 + 𝐻2 𝑂 → 𝑁2 𝐻4 + 𝐶𝑂2 𝑁2 𝐻4 + 𝑂2 → 2𝐻2 𝑂 + 𝑁2 Carbohydrazide has passivating properties due to the reaction with the steel surface to promote the formation of magnetite in the boiler system. Carbohydrazide can be used in all types of pressure boilers. NOTE: When in an aqueous solution, hydrazine will oxidize to nitrogen and water in the presence of air over a relatively short period. Sampling is recommended in acidic solution to prevent oxidation of hydrazine in alkaline solution with atmospheric O2. METHOD OF DETERMINATION For Carbohydrazide Carbohydrazide is determined as DEHA (Diethylhydroxylamine), because of method availability. (DEHA) or other oxygen scavengers present in the sample react with ferric iron in DEHA Reagent 2 Solution to produce ferrous ion in an amount equivalent to the DEHA concentration. This solution then reacts with DEHA 1 Reagent, which forms a purple color with ferrous iron proportional to the concentration of oxygen scavenger. Test results are measured at 562 nm. This method reacts with all oxygen scavengers and does not differentiate samples containing more than one type of oxygen scavenger. Substances which reduce ferric iron will interfere. Substances which complex iron strongly may also interfere. Rev. 2.0 Sep 2023 Page 20 of 56 WATER ANALYSIS TRAINING MANUAL METHOD & REFERENCE As per WI-015-XX-407 WI Reference: HACH® DR3900 Iron Reduction Method for Oxygen Scavengers Method 8140 Rev. 2.0 Sep 2023 Page 21 of 56 WATER ANALYSIS TRAINING MANUAL DETERMINATION OF UREA UREA IN HIGH RANGE Urea in wastewater is determined by colorimetric method of determination with the reagent of p-DMAB same like Hydrazine. The color of the resulting complex is same but maximum absorbance of urea-p-DMAB complex occurs at 525nm. The sample blank must be carried out along the estimation because wastewaters are highly turbid. This method is employed for determination of urea in high range (0-100 ppm). pH dependence is important like in hydrazine. Complex is stable in high acidic media, pH