Unit 9 Neutralization Titrations-Ii

Estimations Based On Kinetic and Acid-Base UNIT 9 NEUTRALIZATION TITRATIONS-II Equilibria Studies Structure 9.1 Introduction Objectives 9.2 Non-aqueous Titrations 9.3 Role of Solvents in Acid-Base Reactions 9.4 Solvent Systems 9.5 Importance of Dielectric Constant 9.6 Hammett’s Acidity Functions 9.7 Titrants and End Point Detection 9.8 Some Applications 9.9 Summary 9.10 Terminal Questions 9.11 Answers 9.1 INTRODUCTION In the last unit you have learnt about the neutralization titrations in aqueous medium. I hope you know that water is poorly dissociated. But that does not prevent water to dissolve many of the electrolytes. Quantitative methods of analysis have been developed by these sort of rapid ionic reactions. It is economical and easier to handle aqueous solutions leading to the wide use of aqueous solution for analysis. But in many instances it is seen that non-aqueous ionizing solvents are advantageous in case of acidimetry and alkalimetry. This is specially true for cases where compounds cannot be titrated in an aqueous medium. In this unit I am going to introduce you to the non-aqueous titrations, the purpose of a particular solvent in specific reactions and also the various solvent systems. Have you heard about the term “dielectric constant”? Well in this unit you will also learn about its importance. After discussing Hammett’s acidity functions, titrants, end point detection, I will also bring to your notice certain applications for such titrations. This unit will help you to learn about the different aspects of nonaqueous neutralisation titrations which are utilised in analytical chemistry. Objectives After studying this unit you would be able to: • understand nonaqueous titrations • understand the difficulties encountered in the titration of a dilute solution of a weak acid with alkali solution • discuss the role of solvents in acid-base reactions with the help of acid-base concept of Bronsted and Lowry • discuss the conjugate acid-base pairs in a given Bronsted acid-base reaction • understand the three groups of nonaqueous solvents, i.e. amphiprotic, aprotic and basic • discuss the importance of dielectric constant • understand Hammett’s acidity function • discuss the different titrants and end-point detection • discuss some applications of nonaqueous titrants 70 Neutralization 9.2 NONAQUEOUS TITRATION Titrations-II There are limitations of using water as a solvent for acid-base titrations and often the usage of non-aqueous solvents is advantageous. Nonaqueous titration is the titration of substances dissolved in nonaqueous solvents. Why do you need to know about this sort of titrimetric procedure? Because it is suitable for the titration of very weak acids and very weak bases and it provides a solvent in which organic compounds are soluble. Titration of weak acid is very difficult and this difficulty may be overcome by using a basic solvent. In many cases, like many organic acids will dissolve in methanol. There are many problems in the alkalimetric determination of weak acid in aqueous medium. Acids and bases with ionization constants less than about 10 − 7 to 10 − 8 are too weak to be titrated accurately in aqueous solutions by conventional methods as discussed in the previous unit. If you choose a solvent less basic than water, it is possible to titrate much weaker bases. Same principle can be applied for titrating weak acids. So, if you choose a solvent less acidic than water, you can titrate a very weak acid. You can see by doing experiments too that, whenever a strong acid (like 0.1N HCl) is titrated with a strong base (0.1N NaOH), then the inflection or the jump in pH at the equivalence point is by almost 5.4 units of pH. So it is much easy to calculate the equivalence point in this case. But the problem arises when it is a weak acid, like CH 3COOH and the hump in pH is only by 2.3 pH units. So if you refer to Fig. 9.1, you can easily follow that the pH jump decreases as the strength of the acid (which is directly measured by dissociation constant) decreases. In other words you can say that as the dissociation constant decreases (for weak acids), the pH jump at the equivalence point also decreases. You have to carry out non aqueous titration because weak acid & weak bases are not completely ionised when dissolved in water at reasonable concentrations (~ 0.1 M), but when we use non aqueous solvent it is strongly acidic in nature or basic in nature. Due to this nature, it ionises the given organic or inorganic substance into it. So if you have to have complete ionisation of weak acid and bases, then you must use nonaqueous solvent. The problem occurs when the pK A or pK B of the material concerned is ≥ 7. The end point break of the titration curve of a weak acid or a weak base (Fig. 9.1) is not sharp enough when pK A is about 7. Can you find an answer as to why there are such limitations in these type of titrations? Now, consider the reaction of the salt formed at the end point with water: − A− + H O HA + OH 2 … (9.1) − When the pK A of HA is nearly 7, the basicity of A is considerably high so that it can readily accept a proton and form HA. So the neutralization is not 100%. In fact it is even less than 99% in these cases. But the limiting value for quantitative neutralization is 99.9%. This situation is further worsened with the increase in the dilution of the acid/alkali solutions when the inflection becomes still smaller. It is very difficult to make out the end point of titration when the inflection is very small, that is even less than 2 pH units. So there will be enormous error in such cases. 71 Estimations Based On Kinetic and Acid-Base Equilibria Studies Fig. 9.1: Titration curve of a weak acid or a weak base By this time you know that phenolphthalein is a suitable indicator in the titration of a weak acid with strong base. When a small amount of a weak acid is given we have to use a dilute solution of NaOH for the titration. It has been observed that if 0 ·01N or more dilute solution of NaOH is used, the pink colour of the indicator fades away rapidly at the end point. This causes difficulty in the recognition of the end point. (A faint pink colour first appears. Now, it is necessary to shake the solution so that there is a thorough mixing of the added titrant. When this is done the pink colour practically disappears. This creates a doubt that the end point has not been reached. In order to confirm the end point, a drop or two of the titrant are then added when solution again appears to be faint pink but on shaking it fades away rapidly. This is known as fleeting end point. Due to this problem, more titrant has to be added to locate the end point than required stoichiometrically). Another problem in the titration of a dilute solution of a weak acid is the interference due to atmospheric CO 2. Hence the titration is done at a higher temperature to drive out the dissolved CO 2. But if the acid under question is volatile, a part of the acid will be lost. The difficulties encountered in the titration of a dilute solution of a weak acid with alkali solution can be summarised as: i) the inflection on the pH-neutralisation curve is small, ii) the phenolphthalein colour fades away at the end point, and iii) interference due to atmospheric CO 2. The above difficulties can be overcome if by some means the dissociation constant of the weak acid can be increased so that it behaves like a strong acid. This can be done by using a suitable non-aqueous solvent in place of water. In order to understand the increase in the dissociation of a weak acid in a suitable solvent we must study the concept of acids and bases suggested by Bronsted. SAQ 1 a) Why the end-point of the titration curve of a weak acid or a weak base is not sharp enough, when pK A is about 7 ? …………………………………………………………………………………………... …………………………………………………………………………………………... 72 b) How can the difficulties encountered in the titration of a dilute solution of a Neutralization weak acid with alkali solution be overcomed? Titrations-II …………………………………………………………………………………………... …………………………………………………………………………………………... …………………………………………………………………………………………... …………………………………………………………………………………………... …………………………………………………………………………………………... 9.3 ROLE OF SOLVENTS IN ACID-BASE REACTIONS Acid-Base Concept of Bronsted and Lowry The Arrhenius theory successfully explains acid-base reactions in aqueous medium. But it cannot be extended to acid-base neutralization reactions in non-aqueous medium. Bronsted and Lowry in 1923 came up with a new concept defining acids and bases in terms of proton transfer; an acid being a substance which donates proton whereas a base acts as an acceptor of proton. This concept can be extended to many substances which could not be included as acids and bases in the original Arrhenius theory. For example, NH + + H O → NH + H O+ 4 2 3 3 … (9.2) In the above reaction, a proton is being donated from ammonium to a water molecule, + hence according to Bronsted-Lowry theory, NH 4 is an acid and H 2O is a base. Here the concept of base is also different. In the Bronsted model, any substance in any medium is a base if it can accept a proton. Thus in the following reactions, + → + + − NH 3 H 2O NH 4 OH … (9.3) NH + HCl → NH + Cl − 3 4 … (9.4) ammonia will be a base because it can accept proton although it does not contain hydroxyl group.The Bronsted theory can be applied to nonaqueous solvents.

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