Sodium hydroxide weak acid-strong base titration

A typical weak acid-strong base titration is that of acetic acid with sodium hydroxide. The net ionic equation for the reaction is... 

For a weak acid/strong base titration, such as that between 0.1 moldm ethanoic acid and 0.1 moldm sodium hydroxide, the indicator needs to change between pH values 6 and 10. Phenolphthalein is a suitable indicator, but methyl orange is not. 

Weak acid with a strong base. In the titration of a weak acid with a strong base, the shape of the curve will depend upon the concentration and the dissociation constant Ka of the acid. Thus in the neutralisation of acetic acid (Ka— 1.8 x 10-5) with sodium hydroxide solution, the salt (sodium acetate) which is formed during the first part of the titration tends to repress the ionisation of the acetic acid still present so that its conductance decreases. The rising salt concentration will, however, tend to produce an increase in conductance. In consequence of these opposing influences the titration curves may have minima, the position of which will depend upon the concentration and upon the strength of the weak acid. As the titration proceeds, a somewhat indefinite break will occur at the end point, and the graph will become linear after all the acid has been neutralised. Some curves for acetic acid-sodium hydroxide titrations are shown in Fig. 13.2(h) clearly it is not possible to fix an accurate end point. 

In many titrations, one solution—either the analyte or the titrant—contains a weak acid or base and the other solution contains a strong base or acid. For example, if we want to know the concentration of formic acid, the weak acid found in ant venom (1), we can titrate it with sodium hydroxide, a strong base. Alternatively, to find the concentration of ammonia, a weak base, in a soil sample, titrate it with hydrochloric acid, a strong acid. Weak acids are not normally titrated with weak bases, because the stoichiometric point is too difficult to locate. 

It is possible by suitable selection of indicators to titrate separately a strong acid and a weak acid or a strong base and a weak base in a mixture of the two. Let us consider, for example, a solution of sodium hydroxide and ammonium hydroxide. If strong acid is added until the pH is 11.1, which is that of 0.1 N ammonium hydroxide solution, the strong base will be within 1% of neutralization (Fig. 20-3). Hence by using alizarine yellow [pK 11) as indicator the concentration of strong base can be determined, and then by a second titration with methyl orange the concentration of ammonium hydroxide can be found. 

The titration of a weak acid (or base) is illustrated in figure C. 14(b). The reagent used is a weak base (or acid). In other methods the end-point of such a titration would not be sharp on account of hydrolysis, but this effect does not detract from the accuracy of a conductance titration. A semi-strong acid (pK 2-4) does not give a very satisfactory end-point with either strong or weak base in this case some ammonia is added to the acid in the cell, and titration is then carried out with sodium hydroxide. This completes the neutralisation of the acid, after which the ammonium ion present is replaced by the sodium ion, with a fall in conductance. Finally the conductance rises as... 

The most common strong base for titrating acidic analytes in aqueous solutions is NaOH. Sodium hydroxide is available both as a solid and as an approximately 50% w/v solution. Solutions of NaOH may be standardized against any of the primary weak acid standards listed in Table 9.7. The standardization of NaOH, however, is complicated by potential contamination from the following reaction between CO2 and OH . 

Discussion. The hydroxides of sodium, potassium, and barium are generally employed for the preparation of solutions of standard alkalis they are water-soluble strong bases. Solutions made from aqueous ammonia are undesirable, because they tend to lose ammonia, especially if the concentration exceeds 0.5M moreover, it is a weak base, and difficulties arise in titrations with weak acids (compare Section 10.15). Sodium hydroxide is most commonly used because of its cheapness. None of these solid hydroxides can be obtained pure, so that a standard solution cannot be prepared by dissolving a known weight in a definite volume of water. Both sodium hydroxide and potassium hydroxide are extremely hygroscopic a certain amount of alkali carbonate and water are always present. Exact results cannot be obtained in the presence of carbonate with some indicators, and it is therefore necessary to discuss methods for the preparation of carbonate-free alkali solutions. For many purposes sodium hydroxide (which contains 1-2 per cent of sodium carbonate) is sufficiently pure. 

The theory of titrations between weak acids and strong bases is dealt with in Section 10.13, and is usually applicable to both monoprotic and polyprotic acids (Section 10.16). But for determinations carried out in aqueous solutions it is not normally possible to differentiate easily between the end points for the individual carboxylic acid groups in diprotic acids, such as succinic acid, as the dissociation constants are too close together. In these cases the end points for titrations with sodium hydroxide correspond to neutralisation of all the acidic groups. As some organic acids can be obtained in very high states of purity, sufficiently sharp end points can be obtained to justify their use as standards, e.g. benzoic acid and succinic acid (Section 10.28). The titration procedure described in this section can be used to determine the relative molecular mass (R.M.M.) of a pure carboxylic acid (if the number of acidic groups is known) or the purity of an acid of known R.M.M. 

It may be noted that very weak acids, such as boric acid and phenol, which cannot be titrated potentiometrically in aqueous solution, can be titrated conductimetrically with relative ease. Mixtures of certain acids can be titrated more accurately by conductimetric than by potentiometric (pH) methods. Thus mixtures of hydrochloric acid (or any other strong acid) and acetic (ethanoic) acid (or any other weak acid of comparable strength) can be titrated with a weak base (e.g. aqueous ammonia) or with a strong base (e.g. sodium hydroxide) reasonably satisfactory end points are obtained. 

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. 

To select an indicator for an acid-base titration it is necessary to know the pH of the end point before using equation (5.5) or standard indicator tables. The end point pH may be calculated using equations (3.27), (3.29) or (3.30). Alternatively, an experimentally determined titration curve may be used (see next section). As an example, consider the titration of acetic acid (0.1 mol dm 3), a weak acid, with sodium hydroxide (0.1 mol dm-3), a strong base. At the end point, a solution of sodium acetate (0.05 mol dm 3) is obtained. Equation (3.28) then yields... 

One of the techniques used to characterize an unknown base or acid is titration. Titration is when a strong base (e.g., sodium hydroxide, NaOH - Na+ + OH ) is added to a weak acid (e.g., HA in water). For example,... 

It is thus possible to calculate the whole of the pH-neutralization curve of a weak acid by a strong base equations (45) and (47) are used for the beginning and end, respectively, and equation (43), without the activity corrections, for the intermediate points. The pH values obtained in this manner for the titration of 100 cc. of 0.1 n acetic acid, for which ka is taken on 1.75 X 10 with 0.1 n sodium hydroxide are quoted in Table LXXI. 

It is of interest to compare the titration curves for the weak acid with a strong base with the titration curve for a strong acid with a strong base. In Illustration 15.1-2 and here we have used equal amounts of acidic solutions of equal concentrations, and titrated both with the same sodium hydroxide solution. However, we see that the initial parts of the titration curve look somewhat different while the parts beyond neutrality are identical. In particular, the titration curve for the strong acid starts at a much lower pH than the titration curve for the weak acid. 

You might think that all titrations must have an equivalence point at pH 7 because that is the point at which concentrations of hydrogen ions and hydroxide ions are equal and the solution is neutral. This is not the case, however. Some titrations have equivalence points at pH values less than 7, and some have equivalence points at pH values greater than 7. These differences occur because of reactions between the newly formed salts and water, as you will read later. Figure 18.22b shows that the equivalence point for the titration of methanoic acid (a weak acid) with sodium hydroxide (a strong base) lies between pH 8 and pH 9. 

In titrating a weak acid with a strong base) or a weak base with a strong acid) greater care is needed in the selection of an indicator. Let us consider the titration of 0.2 N acetic acid, a moderately weak acid with = L80 X 10 , with 0.2 N sodium hydroxide. When an amount of the alkali equivalent to that of the acid has been added, the resultant solution is the same as would be obtained by dissolving 0.1 mole of the... 

The very weak value explains why it is not possible to titrate in water an amino acid by a strong base nor is it possible to titrate, in the same conditions, protonated aliphatic amines. This is the same problem as that encountered for the titration of the third acidity of phosphoric acid by sodium hydroxide (see Chap. 9). 

A strong acid is released during the reaction. An important point to note is that the formed oxime does not exhibit any basic character also, the initial solution was neutral. This is easily explained by the fact that hydroxylamine is a strong base and, hence, its conjugate acid a very weak one. It is interesting also to note that the reaction medium is nonaqueous for the most part, since it only contains 10% water (in mass). Water is added essentially in order to dissolve the hydroxylamine hydrochloride. The medium also contains pyridine in order to displace the oxime formation equilibrium toward the right by formation of the pyridinium ion. The hydrochloric acid released is titrated with a methanolic sodium hydroxide solution. The indicator chosen is bromophenol blue, whose color-change interval is located on the acidic side. Actually, it is the pyridinium cation that is titrated with the methanolic sodium hydroxide solution. The pyridinium cation results from the protonation of the pyridine by the released hydrochloric acid. 

The addition of pyridine is doubly advantageous. On the one hand, it displaces the equilibrium of the reaction toward the right. Behaving in such a manner, it precludes the precipitation of silver oxide. On the other hand, the displaced protons become fixed on the nitrogen of pyridine, which is added in excess. The pyridinium ion, which is hence formed mole to mole with respect to phenytoin, can, as a result, be titrated with sodium hydroxide in the presence of phenolphthalein (titration of a weak acid with a strong base). 

I would now like to consider the titration of acidic compounds in nonaqueous solutions. If you wish to titrate an acid in nonaqueous solution, you should choose a solvent that is not acidic and a titrant that is as strong a base as possible. The paper that really aroused people s imagination and created a lot of interest was the one published by Moss, Elliot, and Hall in 1948, in which they introduced ethylenediamine as a solvent. This compound certainly doesn t have any acidic properties and these authors showed that you can titrate phenol, which is normally too weak to titrate as an acid. In recent years, however, the trend has been away from the use of strongly basic solvents because they have a leveling effect on many bases and they are somewhat unpleasant to handle. Solvents now in use are pyridine, which is an inert solvent and a very weak base, acetonitrile, and acetone. Acetone and certain other ketones are surprisingly good. Recently we have done some work with tertiary butyl alcohol, an excellent solvent for certain cases. Sodium or potassium hydroxide can be used as tltrants, but these are not particularly... 

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