Alkalinity titration endpoint

The alkalinity is measured by titration with a standardized solution of a strong acid such as H2SO4, HNO3, or HCl. The titration curve is identical, but inverse to the acidity titration curve in Fig. 5.6. The caustic alkalinity titration endpoint near pH 11 measures free OH from strong bases. At this endpoint HCO3 = OH . At the carbonate alkalinity endpoint (pH = 8.3) conditions are identical to those defined above for C02-acidity. 

The important total alkalinity titration endpoint near pH 4.5 is defined exactly by... 

We can now compute the pH of the alkalinity titration endpoint for different total carbonate concentrations. Some results are as follows... 

Primarily the sum of carbonate, bicarbonate and hydrate ions in water, but phosphate, silicate etc. may also contribute partially to alkalinity. Normally expressed as ppm (mg/1) CaC03. Phenolphthalein alkalinity (P Aik.) is that portion of alkalinity titrated with acid to pH 8.2 end-point, while total alkalinity (T Aik. or M Aik.) is that titrated with methyl orange indicator to pH 4.2 endpoint. 

For analysis in solutions, the most frequently used CL reaction is alkaline oxidation of luminol and lucigenin in the presence of hydrogen peroxide as oxidant, although sodium hypochlorite, sodium perborate, or potassium ferricyanide may also be used. CL reactions involving alkaline oxidation have been used to indicate acid-base, precipitation, redox, or complexometric titration endpoints either by the appearance or the quenching of CL when an excess of titrant is present [114, 134], An example of these mechanisms is shown in Figure 14. 

Fig. 15.4. Use of reaction modeling to derive a fluid s carbonate concentration from its titration alkalinity, as applied to an analysis of Mono Lake water. When the correct HCO3 total concentration (in this case, 25 100 mg kg-1) is set, the final pH matches the titration endpoint.

Acidity and alkalinity titrations determine the total capacity of natural waters to consume strong bases or acids as measured to specified pH values defined by the endpoints of titrations. Of more interest for many purposes is the ability of a water or water-rock system to resist pH change when mixed with a more acid or alkaline water or rock. This system property is called its buffer capacity. Buffer capacity is important in aqueous/environmental studies for reasons that include ... 

Titrations are carried out to end-points , defined by inflection points in titration curves. These theoretical end points are often difficult to determine in practice, so fixed endpoints are now universally used, defined by the change in color of an indicator substance, or by a pH reading. For alkalinity titrations, a pH of about 4.3 is used, as indicated by an electrode or by an indicator such as methyl orange. In other words, any solution having a pH greater than 4.3 is titrated with a standard acid, and the quantity required to lower the pH to 4.3 is reported in milliequivalents per liter of solution (meqL-1).5 Alternatively, this quantity of milliequivalents may be converted into an equivalent quantity of calcite (CaCOs), or of HCCfj". 

The pH at the true endpoint of the total alkalinity titration should be that of a solution of H2CO3 and H2O. Since we will assume that the alkalinity titration is a closed system (no atmosphere), the solution at the endpoint of the titration should be an H2CO3 solution with a Ct.coj ihat equals the Cy.coa Ih solution being titrated. We refer to the pH of such a solution as pHcoa solution a CT.coa-molar solution of CO2. 

Figure 4-19 shows the relationships between the acidity and alkalinity titration curves and the pC-pH diagram. It shows that 1 equivalent fraction of strong acid, f = 1), must be added to lower the pH from pHcoj - to pHncOj-/ the carbonate alkalinity endpoint. An additional equivalent fraction must be added to lower the pH from pHncoa- to pHcoj. the total... 

Fig. 4-20. The pH values for alkalinity and acidity titration endpoints in a closed system as a function of the initial total carbonate concentration.

The variation of pHco-, with Ct.coj bas led to alkalinity titration methods (employed in Europe) in which the majority of the acid is added, the solution is boiled to expel CO2, and then, after cooling, the small amount of remaining acid is added. The pH of the endpoint is then at about pH 5.1 and is independent of alkalinity, since it is governed solely by the small amount of CO2 generated from the last small acid addition plus the CO2 in the atmosphere. If a sample is warmed and bubbled with nitrogen gas (free of CO2) during titration, the equivalence point would be the pH, where [HsCO ] = 0 and [H+] = [OH ] which at 25°C is pH 7. 

Wet-Chemical Determinations. Both water-soluble and prepared insoluble samples must be treated to ensure that all the chromium is present as Cr(VI). For water-soluble Cr(III) compounds, the oxidation is easily accompHshed using dilute sodium hydroxide, dilute hydrogen peroxide, and heat. Any excess peroxide can be destroyed by adding a catalyst and boiling the alkaline solution for a short time (101). Appropriate ahquot portions of the samples are acidified and chromium is found by titration either using a standard ferrous solution or a standard thiosulfate solution after addition of potassium iodide to generate an iodine equivalent. The ferrous endpoint is found either potentiometricaHy or by visual indicators, such as ferroin, a complex of iron(II) and o-phenanthroline, and the thiosulfate endpoint is ascertained using starch as an indicator. 

When performing the two-phase titration with benzethonium chloride in alkaline medium and in the presence of phenolphthalein as indicator, not only the monosulfonates but also the di- and polysulfonates are determined [19]. The content of di- and polysulfonates is the difference between two titration results giving the amounts of total sulfonates and monosulfonates. Identifying the endpoint requires some experience. 

Alkalinity measurement is also required for the determination of active matter by difference and equivalent weight calculations. It can be determined as two of the following compounds sodium bicarbonate, sodium carbonate, or sodium hydroxide. The sample is titrated to a phenolphthalein endpoint to determine the sodium hydroxide/sodium carbonate content. An added measure of acid converts any bicarbonate to carbon dioxide, which is subsequently removed from the solution. Back-titration of the excess acid gives a measure of the amount of bicarbonate and/or carbonate present. 

An analytical solution for molecules with alkaline functionality is acid/base titration. In this technique, the polymer is dissolved, but not precipitated prior to analysis. In this way, the additive, even if polymer-bound, is still in solution and titratable. This principle has also been applied for the determination of 0.01 % stearic acid and sodium stearate in SBR solutions. The polymer was diluted with toluene/absolute ethanol mixed solvent and stearic acid was determined by titration with 0.1 M ethanolic NaOH solution to the m-cresol purple endpoint similarly, sodium stearate was titrated with 0.05 M ethanolic HC1 solution [83]. Also long-chain acid lubricants (e.g. stearic acid) in acrylic polyesters were quantitatively determined by titration of the extract. 

Most water analysis results are rather easily interpreted. However, two simple and useful tests need explanation. These are the P and M alkalinity. The water is titrated with N/30 HC1 to the phenolphthalein endpoint at pH 8.3. This is called the P alkalinity. Similar titration to the methyl orange end point at pH 4.3 is called the M alkalinity. They are reported as ppm CaC03. 

A more rigorous method for interpreting an alkalinity measurement is to use a reaction model to reproduce the titration. The technique is to calculate the effects of adding acid to the original solution, assuming various carbonate contents. When we produce a model that reaches the endpoint pH after adding the acid, we have found the correct carbonate concentration. 

Mg in fertilizers is based on such proceedings thereof has been applied on multiple occasions. In milk fermentation, where the samples were dried, calcined in a furnace at 600 °C, the ash was dissolved in 0.03 M HCl, the solution was centrifuged and the supernatant was thus analyzed . The complexometric method for determination of Ca(II) and Mg(II) can be carried out in a single titration with EDTA in alkaline solution, using a Ca-ISE for potentiometric determination of two endpoints. This is accomplished on digitally plotting pCa values measured by the ISE as a function of the volume V of titrant added to the aliquot of analyte the first and second inflection points of the curve mark the Ca(II) and Mg(n) equivalences, respectively. ... 

Alkalinity is a measure of the acid-neutralising capacity of water and is usually determined by titration against sulphuric acid to the endpoint of the acid—base reaction. In groundwaters, the carbonate species predominate and an endpoint of... 

Reducing Substances Transfer 1.0 g of sample into a 250-mL conical flask, dissolve it in 20 mL of water, and add 25 mL of alkaline cupric citrate TS. Cover the flask, boil the contents gently for 5 min, accurately timed, and cool rapidly to room temperature. Add 25 mL of 0.6 A acetic acid, 10.0 mL of 0.1 A iodine, and 10 mL of 3 A hydrochloric acid, and titrate with 0.1 A sodium thiosulfate, adding 3 mL of starch TS as the endpoint is approached. Perform a blank determination (see General Provisions), make any necessary correction, and note the difference in volumes of 0.1 A sodium... 

If the sample contains significant acidity or alkalinity, correct the results as follows Dissolve approximately 15 g of sample, accurately weighed, in 40 mL of undistilled pyridine, and add 60 mL of water and 0.5 mL of Phenolphthalein Indicator. If the solution is colorless, titrate with 0.1 N sodium hydroxide to a light pink endpoint, recording the volume required, in milliliters, as v. If the solution is pink, titrate with 0.1 N hydrochloric acid to the disappearance of the pink color, recording the volume required, in milliliters, as v. Calculate the acidity correction factor, A, by the formula... 

To correct for the alkalinity in the sample, place 150 mL of anhydrous methanol into a 500-mL conical flask, add about 50 g of sample, accurately weighed, and swirl to effect solution. Add 1 mL of Indicator, and titrate with 0.1A hydrochloric acid to a yellow endpoint, recording the volume required, in milliliters, as C. Calculate the percent of propylene oxide in thesample taken by the formula... 

Place 25 mL of the decolorized filtrate obtained from the Sodium Chloride test (above) into a 125 mL Erlenmeyer flask, and add 1 drop of a 0.5% phenolphthalein solution in 50% ethanol. Add 0.05 N sodium hydroxide, dropwise, until the solution is alkaline to pH paper, and then add 0.002 N hydrochloric acid until the indicator is decolorized. Add 25 mL of ethanol and about 0.2 g of tetrahydroquinone sulfate indicator. Titrate with 0.03 N barium chloride solution to a red endpoint. Make a blank determination. 

Soap in oil (ppm) Measures alkalinity (as sodium oleate) of oil sample by titration to bromophenol blue endpoint with 0.01 N HCI Cd 17-95... 

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