Ammonia titration curve

Sketch the titration curve for 50.0 mL of 5.00 X 10 M Cd + with 0.010 M EDTA at a pH of 10, and in the presence of an ammonia concentration that is held constant throughout the titration at 0.010 M. This is the same titration for which we previously calculated the titration curve (Table 9.15 and figure 9.27). 

Strong acid with a weak base. The titration of a strong acid with a moderately weak base (K sslO-5) may be illustrated by the neutralisation of dilute sulphuric acid by dilute ammonia solution [curves 1 and 3, Fig. 13.2(a)]. The first branch of the graph reflects the disappearance of the hydrogen ions during the neutralisation, but after the end point has been reached the graph becomes almost horizontal, since the excess aqueous ammonia is not appreciably ionised in the presence of ammonium sulphate. 

Weak acids with weak bases. The titration of a weak acid and a weak base can be readily carried out, and frequently it is preferable to employ this procedure rather than use a strong base. Curve (c) in Fig. 13.2 is the titration curve of 0.003 M acetic acid with 0.0973 M aqueous ammonia solution. The neutralisation curve up to the equivalence point is similar to that obtained with sodium hydroxide solution, since both sodium and ammonium acetates are strong electrolytes after the equivalence point an excess of aqueous ammonia solution has little effect upon the conductance, as its dissociation is depressed by the ammonium salt present in the solution. The advantages over the use of strong alkali are that the end point is easier to detect, and in dilute solution the influence of carbon dioxide may be neglected. 

Acid-Base. The pH of natural waters is determined primarily by the carbonate equilibria. However, organisms may produce amounts of organic matter or ammonia sufficient to influence the pH and buffer capacity of the waters. It would be of interest to determine titration curves of high organic, high color, low alkalinity waters leached from some marshes. It is possible that these waters contain sufficient amounts of organic acids to be significant. 

The results shown in Fig. 15.4 are very interesting like the titration curve in the volumetric analysis, there is a mutation point on the pH curve of the reaction mixture versus the amount of aqua ammonia added. This is opposite to the property of a buffer solution. 

If the strong acid is titrated with a weak base, e.g., an aqueous solution of ammonia, the fii st part of the conductance-titration curve, representing the neutralization of the acid and its replacement by a salt, will be very similar to the first part of Fig. 24, since both salts are strong electrolytes. When. the equivalence-point is passed, however, the conductance will remain almost constant since the free base is a weak electrolyte and consequently has d very small conductance compared with that of the acid or salt. 

Figure 11-5 shows plots of calculated titration curves of 10 M nickel(II), iron(III), and calcium(II) at various pH values in buffers of ammonia-ammonium ion at a total buffer concentration of 0.1 M. Before the end point, the nickel curves at... 

Neutralization of the separator effluent is usually accomplished with sodium hydroxide however, lime or ammonia are occasionally used. A two-stage system using well-mixed tankage of suitable size is recommended. The neutralization invariably occurs on the steep part of the titration curve in these applications. Any inappropriate design, such as an excessively remote pH controllers, will result in a nonfunctional system. As a result, it is highly recommended that an individual experienced in designing pH control systems reviews the final design. 

If the titrant is a weak electrolyte (such as ammonia), the curve is essentially horizontal past the equivalence point, which causes less uncertainty to the extrapolation of a curve. In titration of a weak base, such as acetate ion, with a strong acid, a salt and undissociated acetic acid are formed. After the endpoint is passed, a sharp rise in conductance attends the addition of excess hydronium ions. Salts whose acidic or basic character is too weak to give satisfactory endpoints with indicator are conveniently titrated with the conductometric method. The conductometric titration of a mixture of two acids that differ in degree of dissociation is frequently more accurate than a potentiometric titration. 

Hoijtink et al. [27] also developed an alternative method of generating anionic species, which was improved by Szwarc et al. [28]. The technique involves potentiometric titration of aromatic compounds with a standard solution of Na-biphenylide. The extremely negative reduction potential of biphenyl assures that most of the common aromatics can be reduced to at least their respective radical anions. The values of the thermodynamic reduction potentials are generally obtained from the potentiometric titration curve. As all experiments are usually carried out in ethereal solutions, such as tetrahydrofuran (THF) or dimethoxyethane, problems of follow-up processes are less severe. Later, Gross and Schindewolf [29] reported on the potentiometric titration of aromatics using solvated electrons in liquid ammonia. 

The titration curves for weak bases and strong acids are similar to those for weak acids and strong bases except that they are inverted (recall that strong is added to weak). Figure 19-5 displays the titration curve for 100.0 mL of 0.100 M aqueous ammonia titrated with 0.100 MHCl solution. 

Figure 19-5 The titration curve for 100. mL of 0.100 M aqueous ammonia with 0.100 M HCl. The vertical section of the curve is relatively short because the solution is buffered before the equivalence point. The curve is very similar to that in Figure 19-4, but inverted.

We see that only complexes with formation constants of the order of 106 M-1 or more will lead to titration curves with a sufficiently steep change in pL near the equivalence point (at CM VM / CL VL = 1) to be useful for volumetric analysis. None of the common monodentate ligands, such as the halide anions (Cl-, Br , I-) or the pseudohalides (CN , SCN-, N3 ), form such strong complexes, nor do the carboxylic acid anions (such as acetate) or ammonia (NH3). However, in section 5.2 we will encounter special ligands, the chelates, that do form sufficiently strong 1 1 complexes. 

Because of the volatility of an aqueous ammonia solution, it is more convenient to add hydrochloric acid from a buret to the ammonia solution. Figure 16.5(a) shows the titration curve for this experiment and Figure 16.5(b) shows the titration curve for the case in which a weak base is added from a buret to HCl. 

The titration of a weak base with a strong acid is completely analogous to the above case, but the titration curves are the reverse of those for a weak acid versus a strong base. The titration curve for 100 mL of 0.1 M ammonia titrated with 0.1 M hydrochloric acid is shown in Figure 8.8. The neutralization reaction is... 

Let us now turn to the titration of the base of a weak alkaline pair with the acid of a strong acidic pair using the example of titration of 100 mL of a 0.1 M ammonia solution with the standard solution of hydrochloric acid already used above. At first, the proton potential is —64 kG, which we have already calculated in Sect. 7.3 with the help of Eq. (7.7). The very low proton fill level of 1.2 % in the reservoir NH4 /NH3 is just compensated for by the proton deficiency (which is caused by the OH ions produced by the proton transfer according to NH3 -I- H2O NH + OH ) so that the total fill level in the aqueous solution equals zero. The relatirai in Fig. 7.4, or more exactly, a section of it (Fig. 7.6a), is now what determines the form of the titration curve. 

Fig. 7.6 (a) Fill level of the proton reservoir as a function of the proton potential for an aqueous solution of the acid-base pair NH4 /NH3 (10 mmol in 100 mL solution) at 298 K, (b) Corresponding titration curve of an ammonia solution of equivalent concentration with the acid of a strong acidic pair. 

This example is aimed at illustrating that titration curves may be significantly different for the case that an acid is titrated and the case where the corresponding base is titrated. Fig. 83 displays the titration curves for ammonium ions with sodium hydroxide and ammonia with hydrochloric acid. 

Comparing the two titration curves one can recognize that the steepness of the titration curve at the equivalence point is much larger in case of the titration of ammonia with hydrochloric acid, than for the case of titration of ammonium ions with sodium hydroxide. This means that the random errors will be much smaller when ammonia is titrated with hydrochloric acid, and this titration is strongly to be preferred. 

The titration curve developed in Example 9.1 was hypothetical. In practice, complexometric titrations are conducted at pH values at which the titrant is protonated to some extent. Often, an auxiliary complexing agent such as ammonia is present so that there are various metal-containing species in addition to just the simple hydrated metal ions. Calculations of such titration curves can be readily accomplished within the framework of the material presented in Chapter 5, particularly that dealing with p, the conditional constant. In the treatment that follows, side reactions of the titration ligand and metal ion will be considered using the familiar Ul and... 

Fig. 5 shows the results of both titration experiments. The experimental results are in good agreement with the predictions based upon the equilibrium expressions for Kb the Ka for each indicator, and the mass and charge balances[13]. The data from the acid titration show a sharp equivalence point at approximately 10 m HCl, which suggests that B(OH)4 is still a strong base at 350°C and 0.622 g/mL and capable of neutralizing HCl. This strong acid base titration curve, as was also observed for HCl and KOH, may be contrasted with the weak acid-base behavior observed for the sulfuric acid-ammonia system at 380 C[41]. 

Also shown in Figure 3 is TP hydrolysis of PAM in aqueous solution as measured by the titration method. The data shown were obtained via three separate aging experiments. For comparison with the aqueous hydrolysis curve determined by ammonia analysis, a similarly complex curve has been drawn through the composite data for the three titration experiments. Such a curve appears to give a reasonable fit to the data, but is admittedly arbitrary. Without benefit of the ammonia analyses (curve b), one would probably be inclined to represent the titration data as a smooth curve. A similar comment applies to curve a of Figure 3. 

The calculation that we just did was oversimplified because we neglected any other chemistry of such as formation of MOH, M(OH)2(a ), M(0H )2(5), and M(0H)3. These species decrease the concentration of available and decrease the sharpness of the titration curve. Mg " is normally titrated in ammonia buffer at pH 10 in which Mg(NH3) also is present. The accurate calculation of metal-EDTA titration curves requires full knowledge of the chemistry of the metal with water and any other ligands present in the solution. 

In the mercurimetric titration of thiol compounds in proteins( the flat titration curves became steep after the addition of chloride ions. The titration can be carried out in acetate pH 5-6 buffer with 0-5 M potassium chloride, phosphate pH 7-3 buffer with 0-5 M potassium chloride or borate pH 8-8 buffer with 0-5 M ammonium nitrate. In the presence of ammonia the curves were steep, even in the absence of chlorides. It has been suggested< > that ammonia prevents the bonding of mercury to groups other than the thiol grouping. 

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