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Aqualab Part 1

This study of water, a very ordinary product, based on physical and chemical concepts and carried out as part of RECYCLABS, has made it possible to integrate many of the concepts covered in teaching without using expensive and sophisticated methods. Simple manipulations help us to understand the reality of our environment and how it is changing.<br>What's more, the results obtained are reliable and reproducible, proving that a basic but well-managed practical environment provides access to scientific reasoning.

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Aqualab Part 1

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  1. Discover WATER with AQUALAB ! Part 1 Water is at the heart of sustainable development. It supports all aspects of life on Earth, and access to safe, clean water is a fundamental human right. By Jean-Jacques BOURGOIS www.recyclabs.org 1

  2. Study of the main physical and chemical parameters of natural waters Introduction The educational aim is to study the characteristics of different types of water in order to : • Explain the role of a substance that is ubiquitous in ecosystems. • Show that such a commonplace product can in fact be complex. • Realise that water is a very simple teaching aid for "rediscovering" certain concepts in chemistry and biology. • Take a brief look at technological concepts, for example in aquaculture/aquaria, drinking water quality control and water treatment engineering. • Determine the relationships that exist between the units of measurement used in these fields, which are not harmonised. We will be studying the main characteristics of the following waters : • Mineral water • River water • Tap water • Sea water 2

  3. Study of the main physical and chemical parameters of natural waters Table of contents Section Slides 1. Special case of pure water 4 - 5 Important: 2. pH & Alkalinity 6 - 18 • This presentation is an English version of a reduced content published in French on www.recyclabs.org . • All references are available in that site. The connection between both versions shouldn’t give much trouble. • The experimental protocols pertaining to each section are available on a separate presentation. 3. Water hardness : Ca & Mg 19 - 26 4. Sulfates : frequent compounds 27 - 30 5. Eutrophication : Nitrates 31 - 34 6. Eutrophication : Phosphates 35 7. Seawater salinity and sea salt 36 - 40 General conclusion 41 3

  4. 1.The special case of 'pure' water The purity of water, or more precisely its mineral content, can be measured by its electrical conductivity. Pure water is not a good conductor of electricity. Ordinary distilled water, in equilibrium with carbon dioxide (CO2) from the air, already has a measurable conductivity. Because the electric current is carried by the ions in the solution, the conductivity increases as the ionic concentration increases (Kohlrausch's law). Here are some examples of the conductivity of water from different sources: • Water for pharmaceutical use: ≤ 1.3 μS.cm-1 • Distilled water: 0.5 - 2 μS.cm-1 • Demineralised water: 2 - 10 μS.cm-1 • Tap water: 500 - 800 μS.cm-1 • Drinking (bottled) water: < 1055 μS.cm-1 • Seawater: 40 - 60 mS.cm-1 4

  5. 1. Does pure water really exist? Chemically pure water exists in theory, but in practice it is very difficult to isolate. In fact, if (liquid) water is pure, the only chemical species present in the aqueous medium are H2O, the oxonium cation H3O+ (or the proton H+) and the hydroxide anion OH- (noted more precisely as HO-). These two ions result from the autoprotolysis of water: 2 H2O(l)⇋ H3O+(aq) + OH-(aq) However, this is almost never the case, as water is a very powerful solvent and always contains traces of other ions in solution (electrolytes), including those formed by the solvation of atmospheric CO2. 5

  6. 2. pH & Alkalinity a) Alkalinity refers to water's ability to neutralise acids, due to dissolved bases. It is therefore a measure of the buffering effect of the water in question. The alkalinity of most waters is produced mainly by dissolved carbonate (CO32-), hydrogen carbonate (HCO3-) and hydroxide (OH-) anions. The alkalinity of a water sample is determined by titration with a strong acid (H3O+) to obtain defined pH values. Water parameters: 1) Acid neutralizing capacity to pH 8.2 (KS8.2) : The corresponding acid consumption expressed in mmol.L-1 of H3O+ up to pH 8.2. 2) Acid neutralizing capacity to pH 4.3 (KS4.3): The corresponding acid consumption expressed in mmol.L-1 of H3O+ up to pH 4.3. Important! With the exception of seawater (and some industrial waters), the pH of natural waters is generally below 8.2 ; KS8.2 is therefore zero, and only KS4.3 is taken into account. b) The French degree (°f) chemically expresses the "capacity for acids". - Simple Alkalimetric Titration (TA): is the volume of strong acid (expressed in mL) at 20 mmol.L-1 in H3O+ ions required to titrate 100 mL of water in the presence of phenolphthalein. - Complete Alkalimetric Titration (TAC): is the volume of strong acid (expressed in mL) at 20 mmol.L-1 in H3O+ ions required to titrate 100 mL of water in the presence of methyl orange (or bromocresol green). Remark : TA and TAC are also evaluated by titrimetry . TA : equivalent volume of titrant at pH 8.2. TAC : equivalent volume of titrant at pH 4.3. Basic definition 6

  7. 2. pH & Alkalinity Table 1. Turning points of the pH indicators used to determine °f. pH turning point (Sensitive shade) pH indicator Acid shade Alkaline shade Can be used to assess: TA: KS8,2 Phenolphtalein Colorless 8,2 Rose 9,9 Violet [CO32-] + [OH-] TAC : Methyl Orange (MO) Red 3,1 Orange 4,3 Yellow [CO32-] + [OH-] + [HCO3-] KS4,3 3,6 Green 5,4 » » 4,3 Bromocresol Green (BG) Yellow Blue 7

  8. 2. pH & Alkalinity c) There are other definitions of these parameters depending on the country, which often leads to confusion. In Europe, for example, there are German degrees (°d) for "carbonate hardness" or "Karbonat Härte (KH)", English degrees (°e). In the USA, there are other units! Problem: for German, English and American degrees, the units refer to calcium (Ca)* compounds: CaO: 1°d (KH) = 10 mg.L-1 CaO (German degree) CaCO3: 1°e = 1 grain of calcium carbonate per GB gallon (65 mg / 4.55 L) or 14.25 mg.L-1 of CaCO3 (English degree) : In the USA: parts per million (ppm) is usually defined as 1 mg.L-1 CaCO3 , it is equivalent to mg.L-1, without chemical compound specified, (American degree). *Although alkalinity does not refer to calcium! For conversion factors, see explanations below. 8

  9. 2. pH & Alkalinity • What is the meaning of the chemical equivalence point ? At the equivalence point, the oxonium cations (H3O+) present in the strong acid solution (the titrant, HCl 20 mmol.L-1) and the antagonistic anions present in the analyte (water sample) have been introduced in stoichiometric proportions. Assuming that the antagonistic anions are mainly OH-, HCO3- and CO32-, we obtain, using the formula C1Vw= C2Veq with : C1 = sum of hydroxide and carbonate concentrations (unknown) Vw = volume of water to be analyzed (100 mL) C2 = H3O+ concentration of titrant (20 mmol.L-1) Veq = volume used for titration up to the equivalence point 9

  10. 2. pH & Alkalinity Table 2. The stoichiometric proportions are as follows : OH- + H3O+ → 2H2O x moles x moles CO32- + 2H3O+ → CO2(aq) + 3H2O y moles 2y moles HCO3- + H3O+ → CO2(aq) + 2H2O z moles z moles x mol = [OH-] Vw y mol = [CO32-] Vw z mol = [HCO3-] Vw x + 2y + z = Veq x [H3O+] = ([OH-] + 2[CO32-] + [HCO3-]) Vw Vw x ([OH-] + 2[CO32-] + [HCO3-]) = Véq x [H3O+] Volumes and concentrations are normally expressed in liters and mol.L-1, but if expressed in mL and millimol.L-1, the relationship remains unchanged. 10

  11. 2. pH & Alkalinity $!" $#× H"O%, * Veq #!+ HCO" != #& '&&×V() = ' It comes : OH!+ 2 CO" or : OH!+ 2 CO" #!+ HCO" != If we go back to the definition of the French degree, the Veq gives the TAC (!f), so we have : 5 OH!+ 10 CO" #!+ 5 HCO" != TAC (°f) Since KS8.2 (TA) is zero (in most cases), carbonate ions are absent and the hydroxide ion concentration is negligible because: pOH = 14 - pH = 14 - 4.3 = 9.7 à[OH-] = 10-9.7 mol.L-1. We can therefore make the following approximation : 5 HCO" Now, the capacity for acids up to pH 4.3 (KS4.3) is the corresponding acid consumption in mmol.L-1 of H3O+. At pH 4.3, we're close to the equivalence point when titrating HCO3- ions in the water sample, so we can assess the [HCO3-] present in the sample, since this ion reacts stoichiometrically with H3O+ : HCO3- + H3O+àH2CO3 + H2O and [HCO3-] = [H3O+] = KS4,3 expressed in mmol.L-1 and 5 KS4,3 = TAC (!f) or KS4,3 = 0,2 TAC (!f) != TAC (°f) 11

  12. 2. pH & Alkalinity In drinking water quality, mg.L-1 is used to quantify chemical species. Since the molar mass of HCO3- is 61 g.mol-1, we have : [HCO3-] in mg.L-1 = [HCO3-] in mmol.L-1 x 61. 1 unit of TAC expressed in °f corresponds to 0.2 x 61 = 12.2 mg.L-1 of HCO3-. Special case: aquarium and fish farming In the aquarium and fish farming sectors, hardness is expressed in German degrees (KH (d°) strip tests). This can lead to confusion, as the definitions do not refer to the same chemical species. To remain "chemically correct", units should be standardized. Method : 1) First of all, the values given in °d are converted to mmol.L-1, then back to mg.L-1. 2) The basic definition of °d (KH) is : 1°d = 10 mg.L-1calcium oxide (CaO), but KH, also expressed in °d, refers to the hydrogen carbonate ion HCO3-, so we need to find a correspondence. Expressed in mmol.L-1, 1°d = 0.178 mmol.L-1 CaO (molar mass: 56.1 g.mol-1) It takes 2 HCO3- ions to combine with calcium to form Ca(HCO3)2. We therefore consider that, equivalently : 1°d KH = 0,178 x 2 = 0,356 mmol.L-1 of HCO3–, or 21,78 mg.L-1 of HCO3– 12

  13. 2. pH & Alkalinity 3) Conversion of °d KH in °f TAC : 1 mg.L-1 HCO3– = 0,046 °d KH = 0,016 mmol.L-1 of HCO3– = 0,080 °f TAC, because 5[HCO3–] = TAC (°f) with [HCO3–] expressed in mmol.L-1 . Thus 1°d KH = 0,080 / 0,046 °f TAC 1°d KH = 1,74 °f TAC 1°f TAC = 0,57 °d KH • Case of the English degree 1°e = 1 grain of calcium carbonate per GB gallon (65 mg / 4.55 L) or 14.25 mg.L-1 of CaCO3. Since the molar mass of CaCO3 = 100 g.mole-1, the same reasoning applies as for d° KH: 1 mmol.L-1 CaCO3 = 100 mg.L-1 CaCO3⇒ 1 mg.L-1 CaCO3 = 0.01 mmol.L-1 CaCO3. Alkalinity refers to the HCO3- ion, and it takes 2 HCO3- ions to combine with Ca to form Ca(HCO3)2. We therefore consider that, equivalently: 0.01 mmol.L-1 CaCO3⇔ 0.02 mmol.L-1 Ca(HCO3)2 or 1.22 mg.L-1 HCO3-. 1°e = 14.25 x 1.22 = 17.38 mg.L-1 of HCO3-. 13

  14. 2. pH & Alkalinity • Case of the American degree (same reasoning as for the English degree) Alkalinity is a parameter calculated in mg.L-1 CaCO3 using the equation: With : • Talk: Total alkalinity expressed in mg.L-1 (ppm) CaCO3 • Veq: total titrant volume required to reach bicarbonate (HCO3-) equivalence point (mL) • NH2SO4 : normality of sulfuric acid (eq.L-1) • Vw : sample volume titrated (mL) • MW CaCO3: molecular weight of calcium carbonate (100 g.mol-1) Remark : H2SO4 can be replaced by any strong acid (e.g. HCL). 14

  15. 2. pH & Alkalinity • American degree : example Assume a 200 mL sample of water takes 8.7 mL of 0,02 N (= equivalents.L-1) H2SO4 to reach a pH of 4.3. Calculate alkalinity of the sample as mg.L-1 CaCO3. 1) Talk (expressed in milliequivalents.L-1) = (8.7 x 0.02 x 1000)/200 = 0.87 meq.L-1 in fact 0.87 mmoles.L-1 of HCO3- ! 2) Talk (expressed in mg.L-1 CaCO3) = 0.87 x 50 = 43.5 mg.L-1 (ppm) as CaCO3. Conclusion: 1 mmole.L-1 HCO3- (KS4,3) Û 1 American degree Û Û50 ppm CaCO3 Û 0.02 KS4.3 15

  16. 2. pH & Alkalinity Table 3. Conversion factors based on equivalences in HCO3- anions Looking for ➡ ➡ HCO3- (KS4.3) 1 American degree (ppm CaCO3) 1 mmol.l-1 1 mg.L-1 HCO3- °f °d °e Given value ⬇ 1 mmol.L-1 HCO3- (KS4.3) 1 1 mg.L-1 HCO3- 0,016 1 °f 0,2 1 °d 0,356 ⬇ Conversion factor 5 2.8 0,082 0,046 1 0,57 1,74 61 1 12,2 21,78 3.51 0,058 0,69 1,25 50 0.82 10 Exercise Exercise 1 1 °e 0,285 17,38 1,46 0,8 1 1 American degree (ppm CaCO3) Exercise Exercise 0,02 1.22 0,1 1 16

  17. 2. pH & Alkalinity Theoretically, equivalence corresponds to the inflection point of the titration curve. Mathematically, decreasing function as in this case, it is the minimum of the first derivative of the curve , which corresponds to the minimum of the slope of the tangent at a point on the curve. In Fig. 1, the equivalence point is at pH 4.41 for a titrant volume of 15.93 mL (the point on the titration curve where the slope (m) is minimum: -2.22). In practice, this point is not always exactly located at pH 4.3 (here 4.32), but this does not have a major impact on the difference volume (here 0.03 mL) and subsequent calculations. the point Fig 1. Example of a titration curve of 100 mL tap water with HCl 20 mmol.L-1. for a in titrant 17

  18. 2. pH & Alkalinity Table 4. The following table shows the Veq values (mL) obtained for the different waters using color indicators and pH titrimetry according to the definitions (slide 8). Methyl Orange Bromocresol Green Water pH Average Tap 17,3 17,5 16,2 17,0 Exercise: According to the values in the column « Average», calculate de relevant degrees listed in table 3. Evian (mineral) Muga (river) 28,1 27,8 28,9 28,3 21,5 20,2 21,8 21,2 Sea 15,0 14,5 12,7 14,1 18

  19. 3. Water Hardness : Calcium and Magnesium Definitions Water hardness or « Titre Hydrotimétrique (TH in French) » is an indicator of the mineralization of water due to the alkaline-earth ions calcium and magnesium (Ca2+ and Mg2+). It is customary to relate water hardness to calcium only, i.e. to express Mg2+ content as Ca2+ content. As in the case of alkalinity, there are several definitions and units of hardness, depending on the country concerned. Barium and strontium are also alkaline earth metals, but as they are present in trace amounts in natural freshwaters, their content is not taken into account in hardness calculations. There are three types of hardness: •total hardness = hydrotimetric titre (TH) •permanent hardness is that part of the hardness ions for which there is an equivalent quantity of chloride, sulfate, nitrate, phosphate anions... It is determined on water that has undergone prolonged boiling, which has the effect of precipitating Ca and Mg carbonates (but not other Ca and Mg salts). •temporary hardness = "carbonate hardness" = total alkalimetric titre (TAC). This is the part of the hardness ions for which there is an equivalent quantity of carbonic acid anions (CO32- and HCO3-). Total hardness = permanent hardness + temporary harness 19

  20. 3. Water Hardness : Calcium and Magnesium Remark : in general, seawater hardness is not measured, as calcium and magnesium contents are constant: 413 mg.L-1 and 1290 mg.L-1 respectively. Basic definition and units : The « Titre Hydrotimétrique » = French degree (TH °f) corresponds to a Ca2+ + Mg2+ concentration equal to 10-1 mmol.L-1. From this definition, as in the case of alkalinity, various conversions can be made: 1) All units are calculated in relation to calcium (Ca2+)or its following combinations: oxide (CaO) or carbonate (CaCO3). 2) Use either mmol.L-1 or mg.L -1 (ppm) 3) Depending on the country, either chemical species and unit are considered: °d: German degree for GH (Gesamthärte) °e: English degree (°Clark) American degree (ppm) CaO mmol.L-1 °f (TH) - - mg.L-1 - °d (GH) °e (Clark) Ca2+ Table 5 summarizes these particularities: CaCO3 - mg.L-1 (ppm) 20

  21. 3. Water Hardness : Calcium and Magnesium Table 6 : equivalence between molecular masses and mg.L-1 g.mol-1 mmol.L-1 mg.L-1 40,1 Ca2+ 40,1 1 56,1 CaO 56,1 1 CaCO3 100,1 1 100,1 mg.L-1 Basic definition Equivalence in CaCO3 10,01 of CaCO3 17,84 of CaCO3 14,25 of CaCO3 1 of CaCO3 4) The definitions become as follows, expressed in mg.L-1 : see Table 7. 1 °f 4,01 of Ca2+ 1 °d 10 of CaO Table 7 1 °e ppm 21

  22. 3. Water Hardness : Calcium and Magnesium Table 8 : Conversion factors between degrees 1 American degree (ppm CaCO3) Looking for ➡ ➡ °f °d °e Given value ⬇ ⬇ Conversion factor 1 °f 1 °d 1 °e 1 0,56 1 0,8 0,71 1,25 1 10 1,78 1,42 17,85 14,25 1 American degree (ppm CaCO3) 0,1 0,056 0,07 1 22

  23. 3. Water Hardness : Calcium and Magnesium Note Hardness is not generally measured in seawater, as calcium and magnesium contents are constant: 413 mg.L-1 and 1,290 mg.L-1 respectively. In marine aquaria and aquaculture, they are measured separately using "kits". Seawater is considered to be very hard due to various dissolved salts. Typically seawater's hardness is in the area of 6,570 ppm (6.57 grams per litre). What do these values mean in reality? Among natural waters, including tap and drinking water, a distinction is made according to e.g. the French degree (TH°f): Range of values for Total Hardness: TH (°f) 0 - 7 7 -15 15 - 30 30 - 40 > 40 Water Very soft Soft Moderately hard Hard Very hard 23 Table 9

  24. 3. Water Hardness : Calcium and Magnesium 1) Principle of the dosage Water hardness evaluation is based on a complexometric titration. The titration is carried out with a solution of the di-sodium salt of ethylene diamine tetraacetic acid (Na2EDTA), symbolized by Na2H2Y. The [ H2Y ] 2- anion is a complex ion which, together with a number of cations, produces stable but colorless compounds. An endpoint reaction indicator must therefore be used to highlight the end of the Na2EDTA assay: Eriochrome Black T (EBT). This indicator will be used to measure Ca2+ and Mg2+ ions. The complexation reactions are as follows Ca2+ + Na2H2Y ⇋ [CaH2Y] + 2Na+ Mg2+ + Na2H2Y ⇋ [MgH2Y] + 2Na+ For these reactions to be used for the determination of these ions (they are dosed together), it is necessary to proceed at a pH close to 10, where the Na2H2Y form is in the majority. The EBT colour changes from purplish pink (presence of Ca2+ and Mg2+)to blue: Ca2+ and Mg2+ are totally complexed by Na2EDTA. 24

  25. 3. Hardness : Calcium and Magnesium 2) Practical method to directly measure the water hardness in French degree TH°f. (For details see « Experimental protocols »). Titrant solution: Na2EDTA 0,01 mol.L-1 Solution to be titrated (sample) : 100 mL water Calculation of total hardness Let : CE Na2EDTA concentration = 0.01 mol.L-1 VE the volume of Na2EDTA consumed during titration (in mL) Vw water sample volume = 100 mL Cw the sum of Ca2+ and Mg2+ concentrations in mol.L-1 in the sample We have : CE × VE = Vw × Cw , hence Cw= (CE x VE)/Vw from which Cw = 10-4 .VE mol.L-1 = 10-1 .VE mmol.L-1 In this case, VE directly gives the TH °f (without unit). 25

  26. 3. Water Hardness : Calcium and Magnesium Exercise : The following table shows the VE values obtained for different waters. Calculate the total hardness degrees according to Table 7 : Water VE (mL) Tap 22,0 Mineral 32,0 River 31,0 Table 10 26

  27. 4. Sulfates : frequent compounds The sulfate (SO42-)anion content is often indicated in the composition of drinking and tap waters (European standard: < 250 mg.L-1), but some mineral waters contain more than 1,000 mg.L-1! These waters have laxative properties! In seawater, it is the most abundant anion (2,700 mg.L-1) after chlorides. Sulfur (S) contributes to protein metabolism, as it is a component of the sulfur-containing amino acids methionine and cysteine. Amino acids are the precursors of proteins. Principle of the dosage The SO42-anions react with the barium Ba2+ cations to form a white, insoluble (pKs = 9.96), quantifiable precipitate of barium sulfate BaSO4. In practice, a solution of barium chloride BaCl2of known concentration is reacted with a solution of SO42-ions of unknown concentration. BaCl2(aq)+ SO42-(aq)→ BaSO4(s)+ 2Cl–(aq) As soon as the BaCl2solution comes into contact with the solution containing sulfate ions, a white precipitate of barium sulfate (BaSO4) appears. Equivalence is reached when n(Ba2+) = n(SO42-). Since the precipitate forms as soon as the two solutions are mixed, and there is no colored indicator to show an excess of Ba2+ ions, another technique is required. Two methods are commonly used based on this principle: turbidimetry and conductimetric titration. Turbidimetry has been used here as conductimetry often gives doubtful results. 27

  28. 4. Sulfates : frequent compounds Turbidity measurement Turbidimetry is the measurement of the degree of cloudiness of a suspension. It can be determined using a conventional colorimeter or spectrophotometer, which measures the decrease, due to absorbance, in the intensity of a light beam of known wavelength passing through the suspension. If the particle volume is kept constant, a law identical to Lambert-Beer's law can be applied with: A : absorbance of the suspension (dimensionless quantity) Io : light intensity of the incident beam I : light intensity of the transmitted beam ε : a constant related to the suspension of particles L : optical path (1 cm-thick cuvettes are generally used) C : "concentration" of suspended particles, i.e. turbidity Absorbance is directly proportional to turbidity, at least for low-concentration suspensions. Important: Make sure that the suspension remains homogeneous during measurements. This is achieved by keeping the BaSO4particle size constant and avoiding settling. A mixture of glycerine, alcohol and soluble salts (high ionic strength) is used to stabilize the particles and prevent rapid sedimentation. 28

  29. 4. Sulfates : frequent compounds Note : The absorbance values obtained are highly dependent on the preparation of the solutions. Absorbance can be measured at any wavelength (here 567 nm). In fact, absorbance is not measured due to a colored solution (electronic transition), but to a loss of light intensity due to turbidity. Measurements are taken using a PASCO spectrophotometer. The data are acquired and analyzed using "Spectrometry " software. Here's the calibration line. The linear fit gives its equation (A = aC + b) and its correlation coefficient. In the case shown opposite, the unknown concentration has an absorbance of 0.09, corresponding to a value of 24.3 mg.L-1 (Evian water). 29

  30. 4. Sulfates : frequent compounds Results : Table 11 shows the results obtained for 3 waters: • mineral water (Evian), enabling us not only to carry out a dosage but also to evaluate the method (turbidimetry) • tap water • river water (Muga, Spain) [SO42-] mg. L-1 Water Maesured value Value on the label Evian 24,3 15 Tap 37 - Muga 80,5 - Table 11 Discussion The only water available for comparison is Evian. Under the conditions of the experiment, there is a 62% difference between the announced value and the observed value. However, with a correlation coefficient of 0.977, we can reasonably assert that the method used is acceptable. It should be noted that to obtain a result close to reality, it would be necessary to repeat the manipulation at least five times and statistically analyze the values obtained. The sulfate content of tap water is well within European standards. The high sulfate concentration in the Muga is probably due to wastewater discharges, as many consumer products contain sulfates. Nevertheless, the measured value is not toxic to the aquatic environment. 30

  31. 5. Eutrophication : Nitrates Eutrophication (from ancient Greek: εὖ / eû, "good", and τροφή / trophḗ, "food") is the process by which nutrients accumulate in an environment or habitat (terrestrial or aquatic). The causes are multiple and can give rise to complex interactions between different factors. The nutrients concerned are mainly nitrogen (mainly from agricultural nitrates and wastewater, and secondarily from automobile pollution) and phosphorus (mainly from agricultural phosphates and wastewater). Eutrophication can be exacerbated by sunshine or water temperature (which is tending to increase with global warming). Source: http://05lovesgeography.blogspot.com/2011/02/eutrophication.html 31

  32. 5. Eutrophication : Nitrates Nitrates (compounds of the anion NO3-) are highly soluble in water, so it's normal to find them in moderate quantities everywhere, in fresh and marine waters, on the surface and in groundwater. In small doses, they are an extremely important nutrient for plants, algae and certain photosynthetic bacteria, and constitute one of the mineral sources that enable them to produce organic matter. The nitrate ion is therefore indispensable in ecosystems. But today, due to the large quantities of nitrates added by agriculture (fertilizers) and urban or industrial waste, the naturally nitrate-poor environments to which most ecosystems are adapted are shrinking considerably. Even allowing for natural denitrification phenomena, low-nitrate waters are increasingly rare, due to runoff and sometimes groundwater recharge. In the countries of the European Union, the presence of this compound in ecosystems is regulated by the "Nitrates Directive ". The nitrate limit set by the WHO is 50 mg.L-1 for water to be considered potable. 32

  33. 5. Eutrophication : Nitrates • Principle of the dosage In anhydrous media, the action of sulfuric acid on sodium salicylate, together with nitrates, produces a mixture of sodium ortho- and para-nitrosalicylate. In a basic medium, the nitrosalicylate anion is released, and its stable yellow coloration enables colorimetric determination at 415 nm wavelength. The addition of Na, K tartrate to the soda prevents the precipitation of calcium and magnesium salts. This colorimetric nitrate ion assay is applicable for NO3- ion concentrations of between 0.15 and 15 mg.L-1. Reactions: the mechanism of the main reaction is typically an aromatic electrophilic substitution (SE2 - nitration) , the simplified balance of which is: Determination is performed spectrophotometrically (PASCO equipment) using a 6-point calibration curve. For details see « Part 2. Experimental protocols». 33

  34. 5. Eutrophication : Nitrates Results 3 waters were analyzed: Evian, tap, and Muga (river). Only mineral water (Evian) was used for comparison. At the sampling site, the Muga water was found to contain NaCl, which prevented the measurement from being carried out. According to the method used, NaCl is part of the dry residue and reacts strongly with H2SO4 . Considering tap water, nitrate content is very low and cannot be detected by the method used here. Here's the result obtained for Evian water : The equation of the regression line is : A = 0.0467C + 0.0401. Since the absorbance of the unknown concentration (Evian) is 0.24, the concentration (C) of NO3- in this water is 4.28 mg.L-1. The label indicates a concentration of 3.8 mg.L-1, giving a difference of 0.48 mg.L-1 (12.6%) with the measured value, which is perfectly acceptable under RECYCLABS conditions. 34

  35. 6. Eutrophication : Phosphates Without phosphorus, life is impossible. A fundamental element of living organisms, phosphorus is essential to both natural ecosystems and agricultural production. However, human activities (agriculture, wastewater, urban expansion, industry) are profoundly altering the phosphorus cycle. As a result, aquatic ecosystems are disrupted, algae proliferate and then decompose, consuming the oxygen needed by many species: this is known as eutrophication. Moreover, phosphorus is not a renewable resource. On a global scale, estimates of future consumption predict that phosphate deposits will be depleted within one to two centuries. Controlling the flow of phosphorus into the environment is therefore vital to restoring degraded environments and securing human nutrition. In the European Union, the use of phosphates in consumer products (textile detergents and dishwashers) is regulated. Phosphate content is not indicated on labels affixed to mineral water bottles. For tap water, the standard in Brussels (Belgium) is based on total phosphorus, which must not exceed 2.185 mg.L-1. Preliminary tests have shown that the determination of orthophosphates (H2PO4- , HPO42- and PO43- ions) at very low concentrations (in the mg.L-1 range) does not give reliable results for RECYCLABS. To obtain correct values, it is necessary to use specific, expensive reagents. However, it is possible to measure inorganic phosphorus in products containing high proportions (e.g. fertilizers) of phosphates using a simple method known as thermometric titration. 35

  36. 7. Seawater salinity and sea salt Introduction • In 1776, the father of "modern" chemistry, Antoine Lavoisier, studied the composition of seawater (SW). Later, in 1819, the Swiss chemist Alexandre Marcet enunciated a principle that was verified in 1884 by the German chemist William Dittmar. This principle, known as "Dittmar's law", states "... the overall salt content varies from one sample (of seawater) to another, yet all samples contain the same ingredients and in identical proportions to one another...". This law of constant proportions is still used today, but the definitions of the various parameters characterizing seawater, including salinity, have evolved considerably. UNESCO considered the problem and developed a thermodynamic approach encompassing several physical parameters such as pressure, temperature, density, etc. This resulted in the creation of the "Law of Constant Proportions". This resulted in the development of a "Seawater Equation of State" called EOS-80. In 2010, it updated all these data and adopted a new version of this equation, renamed "Thermodynamic Equation Of Seawater" (TEOS-10). This equation, based on numerical algorithms, is solved by computer and is mainly used in oceanography. Definitions • Salinity is the quantity expressed in grams of inorganic ions dissolved in one kg of SW. In oceanography, several aspects of this parameter are considered (each with a different definition) derived from TEOS-10, and their study is outside the scope of RECYCLABS. The average salinity of SW in the world ocean is S = 35, which implies that SW contains 3.5% 'dissolved salts' and 96.5% water (by mass!). 36

  37. 7. Seawater salinity and sea salt Measurement methods • Ideally, a direct method should be used, i.e. complete evaporation of the water and gravimetric measurement of the residual solids. However, this technique is not valid as most salts are hygroscopic and some of them decompose into volatile products on heating, giving an erroneous result. Several indirect techniques have been developed to overcome this problem: 1. The volumetric method by Sørensen/Mohr-Knudsen titration. 2. Hydrometer: the density is a function of salinity . 3. Refractometry: the refractive index is a function of salinity. 4. Conductimetry.: application of Kohlrausch's law. The first 3 methods are relatively simple and give good, reproducible results. Conductimetry is more accurate but requires good calibration of the conductimeters. In oceanography, fairly sophisticated equipment is used to compare the conductivity of a SW sample with a standardised conductivity standard; these more professional instruments are known as salinometers. We will cover all 4 methods without using sophisticated and expensive equipment. 37

  38. 7. Seawater salinity and sea salt The samples • Samples of "natural" SW were taken in the Bay of Roses (Catalonia, Spain). The results are compared with synthetic SW used in marine aquariums and with food SW in order to obtain points of comparison. Experimental set-up : see Experimental Protocols • Results (Tables 12 & 13) • The following tables show the different values obtained experimentally by the methods considered and for the three SW samples of different origins. The exact determination of the salinity of an SW sample raises the following question: what is the 'right' method? As explained above, professional oceanography uses salinometers based on measuring the conductivity ratio between that of a sample and that of a standard SW . These devices can give a result with an accuracy of up to ± 2.10-3 salinity units. Here, it is best to compare the various results obtained with a theoretical value calculated from TEOS-10. Salinity is expressed here in g.kg-1. The samples come from the following sources: - Western Mediterranean Sea coast ("natural"), Empuriabrava, Bay of Roses (Spain). - Agua de Mar Mediterranea ("food"). Origin: Western Mediterranean Sea. - Salt for aquarium « hw Marinemix » 33,5 g.L-1 (" artificial"). 38

  39. 7. Seawater salinity and sea salt Density/Salinity Table 12 Refractometry (20°C) Hydrometer (20°C) Digital Hand-held Density 1.027 1.028 1.019 Salinity 37.93 39.24 27.43 Density 1.028 1.028 1.021 Salinity 37 37 27 Density 1.028 1.028 1.020 Salinity 38 38 27 Natural SW Food SW Artificial S W Table 13 Salinity Mohr (Dittmar) Conductivity/Salinity Salinity TEOS-10 K2CrO4 titration Conductometric titration μS.cm-1 Salinity Density Conductivity Natural SW 37.57 33.75 51,045 37.58 37.95 37.6 Food SW 37.88 34 51,940 38.33 39.25 38.35 Artificial SW 31.74 24.1 39,448 28.17 27.4 28.15 39

  40. 7. Seawater salinity and sea salt Discussion of the results • Of the four techniques used to study salinity, only determination by conductimetric titration gave anomalous values: 10% lower for natural and food SW and up to 24% lower for artificial SW than titration with a precipitation indicator. Since conductimetric titration is not subject to operator errors in judging the equivalence point, since it appears objectively on the graph σ = f (VAg+ added), we can wonder about the reason for this anomaly. It could be due to the extreme dilution of the samples during titration (175 x), which would lead to an underestimation of the equivalence point; this remains to be verified. Apart from this observation, the other methods give very similar and reproducible results. Titration with K2CrO4 as equivalence point indicator gives results that are nearly identical to the physical methods (density, refractometry and conductimetry). It can be seen that food SW has a very slightly higher salinity than natural SW, which can be explained by the fact that the place where natural SW is taken is close to an area where softer water can mix. It should be noted that the average salinity of the western Mediterranean Sea is 37.5 g.kg-1. The product intended for marine aquariums (solid salts) has a density of 1.020 for a concentration of 33.33 g which has been verified. If the salts used were not "wet" (hygroscopic) we would obtain a salinity of 32.68 g.kg-1. However, we obtain a value of ≈ 28, which clearly indicates that the salts absorbed a certain amount of atmospheric moisture before reconstitution. 40

  41. General conclusion • This study of water, a very ordinary product, based on physical and chemical concepts, has made it possible to integrate many of the concepts covered in teaching without using expensive and sophisticated methods. Simple experiments allow students to appreciate the reality of our environment and how it is changing. • Morover, the results obtained are reliable and reproducible, proving that a basic but well-managed practical environment provides access to scientific reasoning. Ideally, the results obtained should be subjected to a statistical study (error calculations, multiplication of essays, etc.); this essentially mathematical activity is left to users, to avoid weighing down this content. • This first part is an initial approach of experimental subjects; subsequent sections will focus more on the biochemical and physiological aspects of living organisms. 41

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