Titration alkalinity

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 will next compute the pH of the endpoint, which depends on Cr- Because the endpoint pH is generally below 5, carbonate and hydroxyl terms may be neglected so that H = HCO3. Substituting the expression given in Eq. (5.23) for HCO3 leads to 

For pH values below 8 we can ignore 1 in the expression for oth- Then, dividing by H )/K2 gives 

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

TITLE Titrate MW-36 with HCl to get alkalinity SOLUTION 1 MW-36 pH 7.4 unit mg/1 temp 15.0 Na 61. 

In this section we illustrate the computed alkalinity titration discussed in Chapter 3. We use phreeqc to perform the calculations, using sample MW-36 from the Bear Creek Uranium site (see Chapter 6). The idea is 

to use the alkalinity reported in the analysis to determine the total carbonate content of the solution, mCo3,total, and 

The end-point at pH 4.3 occurs at about 0.0031 moles HC1, giving an alkalinity of (following the example calculation on p. 63) 0.0031 x 100.08 x 0.5 = 0.155gL 1, or 155mgL-1. This is almost exactly the analyzed alkalinity of 153 mgL-1, meaning that carbonate anions dominate in this solution. However, this does not give us the total carbonate content. 

We don t present the results here, but in fact if we titrate the same solution (MW-36) with NaOH to pH 8.3, we find that about 2.5 x 10-4 moles NaOH are required, almost exactly the m co content found by the calculation, as it should be. This confirms the point made in Chapter 3 that the sum of the alkalinity and the acidity titration is equal to the total carbonate content (in the absence of competing acids and bases, and assuming jnHw4.3 and 8.3 end-points). 

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The alkalinity is determined by titration of the sample with a standard acid (sulfuric or hydrochloric) to a definite pH. If the initial sample pH is >8.3, the titration curve has two inflection points reflecting the conversion of carbonate ion to bicarbonate ion and finally to carbonic acid (H2CO2). A sample with an initial pH <8.3 only exhibits one inflection point corresponding to conversion of bicarbonate to carbonic acid. Since most natural-water alkalinity is governed by the carbonate—bicarbonate ion equiUbria, the alkalinity titration is often used to estimate their concentrations. 

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. 

Field measurements included pH, Eh, EC, temperature, depth of water table and depth of sample collection. Separate, field preserved sub-samples were collected for cation (ICP-MS/OES), anion (IC), alkalinity (titration), DOC, P04 and Au and PGE analysis (using carbon sorption). 

As mentioned above, several other weak bases can consume acid during the alkalinity titration of seawater. In order of decreasing ability to react with H", they include ... 

During the standard technique used to measure T.A.. the addition of acid is halted at a pH that is too high to drive the titration of to H3PO4. Therefore, during an alkalinity titration, any HPO ... 

Ribonuclease contains no tryptophan. The absorption near 280 nm is almost entirely resulting from the 6 tyrosine residues. The ionization of tyrosine produces a marked shift to longer wavelengths in the absorption spectrum. The ionization can be monitored near 295 nm. Shugar (293) was the first to point out the abnormal behavior of 3 of the tyrosine residues on alkaline titration. Three titrate normally with apparent pK values near 10, but three do not titrate until much more alkaline pH values have been reached and irreversible alkaline de-naturation has set in. Some typical spectra and difference spectra are... 

This conclusion is borne out by other evidence. Fluorimetric alkaline titration gives a simple curve with a pX of 10.3 similar to the first part of the usual spectral titration. The phenolate ion is nonfluorescent. By itself the observed effect could represent quenching by nonfluorescent free tyrosines of fluorescent buried ones. This is particularly unlikely in view of the action of dioxane. The fluorescence of simple tyrosine peptides increases linearly with volume per cent of dioxane in the solvent. 

Titrimetric measurements alkalinity titration for carbonate and bicarbonate ion determinations, argentometric and potentiometric titrations for determining chloride, and iodo-metric titration for sulfite, chlorine, and dissolved oxygen. 

NH4 is distilled after alkalinization. Titration with standardized 0.01 M H2S04 and a mixed indicator (methylene blue and methylene red) NHJ is distilled from water after alkalinization. Ammonium reacts with Nessler s reagent (I2Hg—2IK) to form a yellow-brown colored complex (410—425 nm)... 

An alternative, simpler, procedure for improving the inflexion in the neutralization of an amino-acid is to add formaldehyde to the solution although this does not affect the acid-titration curve, the one for alkaline titration is changed, as seen in Fig. 107. The effect of the formaldehyde is to increase the strength of the ammonium ion acid which is being titrated, and so the pH inflexion at the equivalence-point becomes much more obvious. This is the basis of the formol titration of amino-acids discovered by Sorensen (1907) approximately 10 per cent of formaldehyde is added to the solution which is then titrated with standard alkali using phenolphthalein as indicator. In the presence of thii concentration of formaldehyde the pH-neutralization curve has a sharp inflexion in the region of pH 9, and so a satisfactory end-point is possible with the aforementioned indicator. 

The natural purine bases adenine and guanine do not show appreciable fluorescence under neutral conditions at room temperature. It was also found that purine itself shows only very low fluorescence intensity as a neutral molecule, but this increases markedly on either acid or alkaline titration. For the purine monocation, emission at 400 nm with a quantum yield (4>f) of 0.008 was observed for the anion emission at 370 nm with of 0.045 was found. Appreciable fluorescence at room temperature can be observed for the adenine cation (Table 16), while, similarly, guanine in neutral solution is also nonfluorescent but emission becomes appreciable on protonation. However, with an increase of the pH beyond 11, decreased fluorescence is observed. A case where the neutral molecule is more fluorescent than either the cation or the anion is presented by purin-2-amine or purine-2,6-diamine as well as their nucleosides (Table 16, 5-7). 

Alkalinity titrations of the 15 sequential batch stabilized/solidified waste extraets and of the acetic acid digestates were performed to determine the buffering capacity of the fixed waste and the rate of leaching of alkalinity from the waste. Fig. 1 shows the pH and alkalinity releases associated with one series of extractions. These analyses show a total buffering capacity of the stabilized waste ranging from 17.6 to 19.9 meg/g (dry weight) of fixed waste, with an average of 18.3 meq/g. 

Gran analyses were performed on the titration curves to determine the end points of alkalinity titrations and to calculate the equilibrium constants of the buffering species. The titration curves of the first 3 extracts showed 2... 

The protonation of organic acid anions during alkalinity titration can cause serious errors in alkalinity determination and the use of the data for carbonate speciation and equilibrium chemical modeling. Errors in alkalinity depend on the same factors as for anion-cation balance and may vary from negligible values to as much as 1270% error as documented by Beck et (1974). Accurate values for alkalinity and carbonate speciation can be determined by direct measurement of CO2 evolved from acidified water samples. 

The development of a mathematical model for the carbonate alkalinity titration with a standard acid with correction for volume change, can begin with the solution charge-balance equation, which is... 

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 ... 

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