Acid-base titrations back titration

Where Is the Equivalence Point In discussing acid-base titrations and com-plexometric titrations, we noted that the equivalence point is almost identical with the inflection point located in the sharply rising part of the titration curve. If you look back at Figures 9.8 and 9.28, you will see that for acid-base and com-plexometric titrations the inflection point is also in the middle of the titration curve s sharp rise (we call this a symmetrical equivalence point). This makes it relatively easy to find the equivalence point when you sketch these titration curves. When the stoichiometry of a redox titration is symmetrical (one mole analyte per mole of titrant), then the equivalence point also is symmetrical. If the stoichiometry is not symmetrical, then the equivalence point will lie closer to the top or bottom of the titration curve s sharp rise. In this case the equivalence point is said to be asymmetrical. Example 9.12 shows how to calculate the equivalence point potential in this situation. 

The determination of neomycin by non-aqueous titration has been described by Penau et all2l. Neomycin base is allowed to react with standardised perchloric acid the excess acid is then back-titrated with potassium hydrogen phtha-late using crystal violet as indicator. To determine the neomycin content of the sulphate salt the same authors precipitated the sulphate with benzidine before reacting the neomycin with perchloric acid. The amount of benzidine required to precipitate the sulphate is calculated from the sulphate content which is itself determined by titration with sodium hydroxide. 

In this experiment the neutralizing power of various antacids will be determined. Antacids contain basic compounds that will neutralize stomach acid (stomach acid is HC1). The amount of base in the antacid tablets will be determined by an acid-base titration. It is a back titration method. This method is used because most antacids produce carbon dioxide gas, which can interfere with the titration. By initially adding an excess of acid, one can drive off the C02 by boiling the solution before titrating the excess acid. There are many brands of commercial antacids with various ingredients. A few of the common ones are listed below ... 

The most important determination is normally the concentration of carbon-magnesium-bonded species in solution. For routine estimation of this concentration for freshly prepared solutions of organomagnesium compounds, an aliquot of the test solution may be added to an excess of standard acid, and then back-titrated with sodium hydroxide. However, this simple determination of total base will give a high estimate of organomagnesium content if products of hydrolysis or oxidation are present. Analytical methods based on the determination of the hydrocarbon formed on hydrolysis of the organomagnesium compound... 

If a gas is soluble in a suitable solvent, and the concentration of the solution can be determined by a simple analytical technique, then accurately measured quantities of the gas can be dispensed by using the appropriate volume of the solution. For example solutions of the hydrogen halides in various solvents can be determined by simple acid/base titration and solutions of chlorine in carbon tetrachloride can be determined by addition of excess potassium iodide and back titration with sodium thiosulphate. Refer to textbooks of inorganic analysis for details of these methods. 

Figure 3-2 Acid-base titration curve for hen lysozyme at 0.1 ionic strength and 25°C. O, initial titration from the pH attained after dialysis , back titration after exposure to pH 1.8 A, back titration after exposure to pH 11.1. The solid curve was constructed on the basis of "intrinsic" pK values based on NMR data. From Kuramitsu and Hamaguchi ...

As an alternative, the acid can be introduced in an amount slightly in excess of that needed to convert the sodium carbonate to carbonic acid. The solution is boiled as before to remove carbon dioxide and cooled the excess acid is then back-titrated with a dilute solution of base. Any indicator suitable for a strong acid/strong base titration is satisfactory. The volume ratio of acid to base must of course be established by an independent titration. 

To avoid these errors one determines the acid groups by back-titrating a small excess of base added to the stirred polysiloxane. 

Figure 8.4 illustrates the colors and transition ranges of some commonly used indicators. The range may be somewhat less in some cases, depending on the colors some colors are easier to see than others. The transition is easier to see if one form of the indicator is colorless. For this reason, phenolphthalein is usually used as an indicator for strong acid-base titrations when applicable (see Figure 8.1, titration of 0.1 M HCl). In dilute solutions, however, phenolphthalein falls outside the steep portion of the titration curve (Figure 8.2), and an indicator such as bro-mothymol blue must be used. A similar situation applies to the titration of NaOH with HCl (Figure 8.3). A more complete list of indicators is given on the inside back cover.

In the conventional Kjedahl method, two standard solutions are required, the acid for collecting the ammonia and the base for back-titration. A modification can be employed that requires only standard acid for direct titration. The ammonia is collected in a solution of boric acid. In the distillation, an equivalent amount of ammonium borate is formed ... 

The CO2 released is captured in a base of a known concentration. The excess base is back titrated by a strong acid in the presence of a universal indicator. The chemical equations used are as follows ... 

The concentration of surface centers is often taken to be equal to the number of hydroxyl groups in a monolayer on the surface [40-42] other values were obtained by acid-base titration [7,43], tritium exchange [44], and equilibration experiments combined with model calculations ("long time titrations") [18,46,47]. Table 3 gives an overview of numbers of centers determined by various researchers. Most values are found between 1.0 x 10 and 3.2 x 10 mol/m only that of Schulthess [7] (1.7 X mol/m") was clearly lower, possibly because of the back-titration technique. So the assumption that the number of centers is equal to the number of surface OH groups seems to be acceptable, though in principle different numbers of... 

The fully automatic [3, 4] quasi-equihbrium acid-base titration was performed under a C02-free atmosphere using NaCl back-... 

Figure 1. Potentiometric titration of cationic caprolactam polymers (initiator CL HCl). 1 — Titration with base, 2 — back titration with acid, E — concentrations of acidic and basic groups, = strong adds, Ech-Hci caprolactam HCl, E A ocyl-amidinium groups, EjfHj — ammonium groups, Ecooh — carboxylic groups, Ecoow — carboxylic groups formed by alkaline hydrolysis of acyUactams.

The fermentation-derived food-grade product is sold in 50, 80, and 88% concentrations the other grades are available in 50 and 88% concentrations. The food-grade product meets the Vood Chemicals Codex III and the pharmaceutical grade meets the FCC and the United States Pharmacopoeia XK specifications (7). Other lactic acid derivatives such as salts and esters are also available in weU-estabhshed product specifications. Standard analytical methods such as titration and Hquid chromatography can be used to determine lactic acid, and other gravimetric and specific tests are used to detect impurities for the product specifications. A standard titration method neutralizes the acid with sodium hydroxide and then back-titrates the acid. An older standard quantitative method for determination of lactic acid was based on oxidation by potassium permanganate to acetaldehyde, which is absorbed in sodium bisulfite and titrated iodometricaHy. 

There is also evidence for stable 3,4-adducts from the X-ray analysis of 2-amino-4-ethoxy-3,4-dihydropteridinium bromide, the nucleophilic addition product of 2-aminopteridine hydrobromide and ethanol (69JCS(B)489). The pH values obtained by potentiometric titration of (16) with acid and back-titration with alkali produces a hysteresis loop, indicating an equilibrium between various molecular species such as the anhydrous neutral form and the predominantly hydrated cation. Table 1 illustrates more aspects of this anomaly. 2-Aminop-teridine, paradoxically, is a stronger base than any of its methyl derivatives each dimethyl derivative is a weaker base than either of its parent monomethyl derivatives. Thus the base strengths decrease in the order in which they are expected to increase, with only the 2-amino-4,6,7-trimethylpteridine out of order, being more basic than the 4,7-dimethyl derivative. 

Dodge has based a process for the determination of benzaldehyde. A strong (2 5 N) alcoholic potash solution is required for the estimation, which is performed. by allowing a mixture of 10 c.c. of this solution with 1 to 2 grams benzaldehyde to stand at the ordinary temperature for twenty-four hours, after which the unabsorbed pota is titrated back with N/2 hydrochloric acid. A blank test is also made, and from the amount of potash entering into reaction, the percentage of aldehyde can be calculated. The process breaks down in the assay of natural oil of bitter almonds, probably due to the presence of benzaldehyde cyanhydrin. 

Two excellent methods (utilising acid-base indicators) are available for standardisation. The first is widely employed, but the second is more convenient, less time-consuming, and equally accurate. A third, back-titration, procedure is also available. 

In the direct method, a solution of the ammonium salt is treated with a solution of a strong base (e.g. sodium hydroxide) and the mixture distilled. Ammonia is quantitatively expelled, and is absorbed in an excess of standard acid. The excess of acid is back-titrated in the presence of methyl red (or methyl orange, methyl orange-indigo carmine, bromophenol blue, or bromocresol green). Each millilitre of 1M monoprotic acid consumed in the reaction is equivalent to 0.017032 g NH3 ... 

In connection with the above, we shall still consider the pH curves of the displacement titrations, because in fact they represent a back-titration of an alkaline reacting salt (e.g., NaA) with a strong acid or of an acidic reacting salt (e.g., MX) with a strong base, so that in Fig. 2.19 the foregoing data of equivalence point and initial point are of direct application. 

This is an example of the method of back titration, in which more acid (HCI) is added than is necessary to stoichiometrically react with the base (Mg(OH)2), in order to be certain that all the base has reacted. One then titrates the excess acid with a standardized base solution (NaOH) and in a series of calculations, determines the amount of unknown base (Mg(OH)2). 

Once the student had determined the exact concentration of the base, the student then proceeded to determine the equivalent mass of an unknown acid. To do this, the student measured out 0.500 grams of an unknown solid acid and titrated it with the standardized base, recording pH with a calibrated pH meter as the base was added. The student added 43.2 mL of the base but went too far past the end point and needed to back-titrate with 5.2 mL of the 0.100 M HC1 to exactly reach the end point. 

This is the basis for a common method for the determination of ammonia in soil.1 Soil is suspended in water and placed in a Kjeldahl flask. The suspension is made basic by the addition of a strong (5-50%) sodium hydroxide solution, and the flask is immediately attached to a steam distillation setup. Steam distillation of the suspension carries the released ammonia to an Erlenmeyer flask, catching the distillate in a standardized acid solution that is subsequently back titrated via acid-base titration. The amount of ammonia in soil can be calculated from the end point of the titration. This procedure is similar to a standard Kjeldahl determination and can be carried out using the same equipment, although no digestion is needed. 

In this chapter, we investigate individual types of reactions that meet all the requirements. We will also discuss back titrations, in which some of the limitations that we may encounter are solved. Our discussions in Chapter 4 involved acid-base reactions. These reactions as well as other applicable reactions will also be discussed here. 

Calcium carbonate is not soluble in water. The addition of the standard acid to a tablet immersed in water causes it to dissolve, but rather slowly. A direct titration is not appropriate because the reaction is not fast, and it would be difficult to detect the end point. A back titration is therefore useful because the standard acid can be added in excess so that the calcium carbonate completely dissolves. The excess acid can then be titrated with a standard base. 

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