Quantitative analysis of concentration of magnesium ion in tap water by using EDTA?
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If you asked me this same question about 5 yrs ago, I couldve answered without blinking. Now I would have to go back to my notes from when I actually had done the experiment to be of any help. Good Luck
130307
as far as my knowledge goes SDS is a deterget it helps in distroying the proteins in the sample,EDTA is a chelator which chelates the ions lik Ca+ and Mg[req by enzym to break DNA],and chloroform being organic solvent keeps the DNA away frm the water and prevts it frm getting dissolved
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http://www.itt.edu/staff/boyter/links/index.html#CHEMISTRY_ANALYTICAL
280765
Look at the links at\n\nhttp://www.itt.edu/staff/boyter/links/index.html#CHEMISTRY_ANALYTICAL
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Na2EDTA is a complex molecule. Its name stands for ethylenediaminetetraacetic acid – disodium salt. The formula is Na2C10H14N2O8. Even though its name and structure are complex, you only need to know that: EDTA reacts with Ca+2 and Mg+2 in a one to one ratio as follows:\n:Ca+2 + EDTA-2 -----> CaEDTA\n\nMg+2 + EDTA-2 -----> MgEDTA\nThe formula mass of Na2EDTA dihydrate is 372.24 g/mole.Na2EDTA comes as a dihydrate(absorbs water from the atmosphere) which is efflorescent (literally means "bursting into flower" but in this case just means that it forms bubbles), you cannot reliably weigh out the desired number of grams. The salt may be fully, or only partially, hydrated at the time you weigh it. The result is that you cannot be certain that you have prepared the proper strength reagent of EDTA..therefore: you need to standardize the solution  The results of the titration is reported as mg of CaCO3/liter.\nTo convert Calcium plus Magnesium (meq/L) Hardness as CaCO3 (mg/L) 50\nOf course since Fe and Mg and Ca and other metals are also present in water it is better (more accurate) to titrate for each metal species individually and calculate the hardness using the results of the individual results. EDTA hardness is sometimes referred to as "Total hardness"and is reported as " Hardness as mg/l CaCO3."
1000766
The low technology approach is to use a densitometer.  This can be as simple as a hydrometer as used in the brewing industry to check for specific gravity (density).  From a measured density and a known volume and if you know the densities of both liquid components of a binary mixture you can work out the relative proportions of each using simultaneous equations and hence the concentration of one in the other.
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From this patent\n\nwww.freepatentsonline.com/6962719.html\n\n\n\nThe total extract described in EXAMPLE 1 was loaded in a dialysis bag with molecular weight cut off 500 (Spectra/Por®, Cat. No. 131 057). Both ends of the bag were sealed with clamps. The bag was placed in a bucket contained deionized water, in which the ratio of the supernatant and deionized water was 1:10. The total extract was dialyzed at 4° C. with stirring thrice each for 20 hours. The solution remaining in dialysis bag was collected and concentrated under reduced pressure to give the extract.
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Water from the tap has several ions in it.  Some are added by people (like fluorine) and some arent (like sodium).  Also, water from the tap has several impurities.  \n\nWhen water is distilled, it is purified.  This water is deionized (some people call it DI water).  The water is free of impurities.  Without impurities, chemical "tests" can be run very efficiently and effectively with water.
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The basic check is: measure the mass, divide by the volume of water displaced, check that the density is correct. After that you should still do a chemical analysis to make sure there aren't other heavy metals, like uranium or lead mixed in.
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There are a number of methods. One way is to crash out the cadmium as an insoluble salt and than weigh that salt, to back-figure the cadmium. The problem there is that you would have interferences from many other metals.\n\nSpectrometrically, you could accomplish it using an ICP or simple atomic absorption. The later might require the use of standard addition, particularly if your matrix is complex. You figure out the absorbance after adding known amounts of cadmium.
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Most common method would be an oxidative digestion (most likely using ultraviolet radiation) in order to eliminate the organic matrix, and subsequent simultaneous ICP-OES analytics. Numerous other detection methods are available (electrochemistry, mass spectroscopy, atom absorption spectroscopy, even classical wet chemistry might give satisfactory results if done with extreme care and insight).\n\nThe heavy metals would be the usual suspects also monitored in drinking water, with special emphasis on any material used or present as impurity in the packaging (tin, lead, cadmium, zinc...)
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You could determine the copper content of brass (an alloy of copper and zinc) in one of three general ways:\ngravimetric,\nvolumetric,\ncolorimetric.\n\nYou first need to dissolve a known mass of brass completely in nitric acid, neutralise, then transfer to a volumetric flask and make up to volume. \n\nIf you then choose to do it gravimetrically, you should then prepare a stable copper compound of high molecular mass OR you could electroplate out the copper and mass it OR you could simply add excess zinc metal, dissolve any left over zinc with say hydrochloric acid or sulfuric acid, dry and mass the copper. \n\nThere are a number of ways that you could titrate the copper: you could add excess potassium iodide which reacts quantitatively with copper to produce copper(I) iodide and iodine. The resulting iodine can then be titrated with standardised thiosulfate using starch as an indicator. You could also do a complexiometric titration with EDTA.\n\nIf you decide on a colorimetric method, you need to make a standard solution of copper (or brass of known copper content - better) and compare the absorbance of your brass solution with the standard. \n\nYou could do two different determinations and compare results.\n\nThere are many other possibilities, but i do not know what a board project is, or your level of expertise and what equipment you have available. If you need more detail, please re-post and i or someone will be able to help you further.
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Anhydrous magnesium sulfate is used to dry the organic solvent phase isolated after an extraction of your solute (compound) from an aqueous phase.  After isolating the organic phase, there is a certain amount of water present (solubility factor, suspended microdroplets of water, how careful you were at separating the layers).  While it is relatively easy to remove the solvent (ie: ether, ethyl acetate, hexane) by evaporation, it is more difficult to remove traces of water.  Often water becomes hydrogen bonded to your solute resulting in a sticky mess that will affect your weighing measurements.  Water is especially hard to get rid of if your compound is a liquid.  Weighing errors will lead to inaccurate concentration measurements.  Inaccurate concentration measurements lead to less than an "A" grade on your experiment.\n\nThe magnesium sulfate hydrates rapidly (forming both magnesium sulfate monohydrate and magnesium sulfate heptahydrate) and is then easily separated from the organic phase containing your dissolved solute by simple filtration.  Voila, the water is now gone, you have a "dry" organic phase which, when evaporated, will give you a dry compound.
833753
sodium, potassium, chloride, trace amounts of other metal ions.
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Peter B has it right -- atomic absorption will do elemental analysis.\n\nAtomic emission spectroscopy is another tack, though a bit more challenging.
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Some info at http://www.aquapurefilters.com/contaminates/165/toc.html and at http://www.stateoftheart.it/toc-2.htm
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Spectrophotometric methods, or those based on the absorption of light by an analytical species, are some of the simplest and most common methods used in quantitative analysis. The amount of light which is absorbed by the sample (A) is correlated with the concentration (c) of the analytical species present using Beer's Law (A = e bc; where e is the molar absorptivity for the analytical species and b is the pathlength of the sample cell). \n\nIn this experiment, the analytical species, iron, will be studied using a colorimetric method in which a reagent is added to transform iron into a highly absorbing complex. Iron is often found in nature in reasonably high concentrations and may enter the hydrosphere through the weathering of iron salts and minerals. Both iron (II) and iron (III) ions are commonly found dissolved in water either in a colloidal form or as inorganic or organic complexes. There are also multiple anthropogenic sources of iron in the environment such as canneries, tanneries, textile mills, and metal cleaning operations. \n\nA relatively simple but sensitive colorimetric procedure for the determination of iron involves first reducing all iron present in the sample to iron (II) ion using hydroxylamine hydrochloride and then complexing with 1,10-phenanthroline in a solution buffered at low pH:\n\nPart 1: Preparation of Iron Standard Solutions\n\nWeigh 351 mg of high quality ferrous ammonium sulfate hexahydrate (FeSO4(NH4)2SO4 Ÿ 6H2O) and quantitatively transfer to a 500 mL volumetric flask. Add 50 mL of distilled water followed by 1 mL of concentrated sulfuric acid. Dilute to the mark with distilled water and mix thoroughly. Using this solution, prepare five standard solutions of iron (II) ion (about 10 - 15 mL should be sufficient) ranging in concentration from 0.5 to 5 ppm.\n\nPart 2: Preparing a Beer's Law Plot\n\nTransfer a 5 mL aliquot of the lowest concentration iron standard solution (probably 0.5 ppm) to a 125 mL Erlenmeyer flask and test the pH with test paper. If the pH is greater than 4.5, add enough 0.2 M sulfuric acid to bring the pH to about 3.5, counting the number of drops that you have added. \n\nAgain, counting the drops that you add, add sodium citrate (259 g/L) buffer to bring the pH to about 4.5. Pipet 1 mL of 10% hydroxylamine hydrochloride and 3 mL of 0.3% 1,10-phenanthroline into the sample and mix the solution well. Allow about 5 min for development of color.\n\nUse the same number of drops of sulfuric acid and sodium citrate for your remaining standard solutions. For each solution, add 1 mL of 10% hydroxylamine hydrochloride and 3 mL of 0.3% 1,10 phenanthroline and mix the solutions well. Allow about 5 min for development of color for each solution.\n\nScan the absorbance of each of the solutions from 300 to 600 nm using the Shimadzu ultraviolet-visible spectrometer, noting the absorbance value of each solution at 510 nm. For each sample, use water as your reference blank.\n\nPart 3: Analysis of samples\n\nWater samples should be measured for each member of your group in addition to determining the iron concentration in an unknown water sample provided by the Instructor. Treat the samples in the same way as the standards: using a 5 mL sample, add sulfuric acid initially, if necessary, to adjust the pH, followed by citrate buffer, reducing agent (1 mL of 10% hydroxylamine hydrochloride), and the indicator (3 mL of 0.3% 1,10 phenanthroline). Mix the solutions well and allow about 5 minutes for development of the complex. \n\nScan the absorbance of each of your sample solutions from 300 to 600 nm using the Shimadzu ultraviolet-visible spectrometer, noting the absorbance value of each solution at 510 nm. Using your Beer's Law plots prepared in Part 2 and the measured absorbances of your samples, determine the iron concentrations in each of your water samples.