Titration of the sample

Before stoppering, I mL of the zinc acetate solution is added to the sample and the bottle is shaken vigourously. The zinc sulphide is allowed to settle. 

From the supernatant solution, 50 mL (V in Section 6.2.5) are pipetted carefully into the titration beaker. Care must be taken to avoid stirring up the precipitated zinc sulphide. Then 1 mL of the iodide solution and exactly 10 mL of the iodate standard solution are pipetted into the titration beaker followed by 5 mL of the diluted sulphuric add. The sample is shaken gently and the titration beaker is covered with a watch glass or plastic film and set aside for approximately 5 min. The surplus of iodine not consumed by the thiosulphate initially present in the sample is then titrated with the stardardized thiosulphate solution. The endpoint is determined as described in Section 6.2.4.L 

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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. 

Acidity. Acidity is the base-neutrali2ing capacity of a sample of water. It is determined by titration of the sample with standard base to pH 8.3... 

The end point of the reaction is conveniently determined electrometrically using the dead-stop end point procedure. If a small e.m.f. is applied across two platinum electrodes immersed in the reaction mixture a current will flow as long as free iodine is present, to remove hydrogen and depolarise the cathode. When the last trace of iodine has reacted the current will decrease to zero or very close to zero. Conversely, the technique may be combined with a direct titration of the sample with the Karl Fischer reagent here the current in the electrode circuit suddenly increases at the first appearance of unused iodine in the solution. 

Using the pKa and the estimated So, the DTT procedure simulates the entire titration curve before the assay commences. Figure 6.7 shows such a titration curve of propoxyphene. The simulated curve serves as a template for the instrument to collect individual pH measurements in the course of the titration. The pH domain containing precipitation is apparent from the simulation (filled points in Fig. 6.7). Titration of the sample suspension is done in the direction of dissolution (high to low pH in Fig. 6.7), eventually well past the point of complete dissolution (pH <7.3 in Fig. 6.7). The rate of dissolution of the solid, described by the classical Noyes-Whitney expression [37], depends on a number of factors, which the instrument takes into account. For example, the instrument slows down the rate of pH data taking as the point of complete dissolution approaches, where the time needed to dissolve additional solid substantially increases (between pH 9 and 7.3 in Fig. 6.7). Only after the precipitate completely dissolves, does the instalment collect the remainder of the data rapidly (unfilled circles in Fig. 6.7). Typically, 3-10 h is required for the entire equilibrium solubility data taking. The more insoluble the... 

The catalyzed decomposition of H202 in aqueous solution is followed by removing equal volume samples at various times and titrating them with KMn04 to determine the undecomposed H202- The results are tabulated. Confirm that the reaction is first order and find the cc of KMn04 required for the titration of the sample removed at t =0. 

For higher accuracy in the low polymer concentration range, two different methods were used. In the case of PAA, potentiometric titrations of solutions of PAA were performed with 0.01 N NaOH using a Brinkman model, Westbury, NY, automated titrator. Blank tests indicated no interfering species. Known amounts of PAA were used to prepare a calibration curve immediately after titration of the samples containing unknown amounts of polymer. The starting point of the titration was pH 4.0, and the end point was reached near pH 8. Total volumes of 75 or 100 cc were used for the titrations, and the ionic strength was controlled at 0.01 M NaCl. 

Alkalinity has to be measured in the field because C02 is often pressurized in groundwater (due to addition of C02 from soil and other underground sources). Upon exposure to the atmosphere some C02 may leave the water, causing part of the HCO3 to break down. For these reasons it is highly recommended that alkalinity be determined in the field. Field-measured alkalinity values are needed for water-rock saturation calculations. Various setups are available for alkalinity measurements in the field by titration of the sample with an acid and pH coloring indicator. 

Spectrophotometric titration [29-32] consists in repeated measurement of an absorbance which changes in the course of titration of the sample solution. The use of this method depends on the existence of a linear relationship between the absorbance measured and the concentration of the absorbing substance in the solution being titrated. The course of the titration is represented graphically by two intersecting straight lines. To find the titration end-point it is necessary to determine the absorbance at two points before and two points behind it. The graphs are drawn in the system of A (absorbance) versus v (volume of titrant solution). To increase the accuracy of determination, corrections are made for the dilution caused by the addition of the titrant solution. 

Titration of the sample. The pooled serum or plasma sample will have been prepared by touching the end of a stining rod to 5 ome Antifoam A and rotating it in the pooled sample. This wiU prevent excess foaming when the sample is swirled. Place 0.100 mL of serum or plasma in a 25-mL Erlenmeyer... 

These conclusions are fully supported by the results in which the oxygen storage capacities (OSC) was measured for the samples. The values of OSC were determined by titration of the samples at 450°C with consequent pulses of CO which were introduced into the dry nitrogen flow passing through the sample bed in a quartz reactor till the formation of CO2 ceases. The results obtained in this series of determinations are given in Table 4. 

The neutralization or acid value (AV) of a lipid sample gives an indication of the unbound or free FA content. This is achieved by direct titration of the sample in an appropriate solvent with alkali (e.g., ISO 660 1996). The AV is defined as the number of milligrams of potassium hydroxide required to neutralize a 1-g sample. Although refined oils are largely devoid of free FA, considerable amounts may be present in crude oils. Their presence may be an index of oil purity. The degree of edibility of a fat is generally considered to be inversely proportional to the total amount of free FA. 

According to the chloride concentration, titrate with 0.05 m or 0.02 m silver nitrate solution until the colour changes from greenish-yellow to reddish-brown. In order to assist determination of the end point of the titration, the analyst may add a drop of sodium chloride solution after completion of the titration of the sample. If the yellow-green colour reappears the end point of the titration was reached. 

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