UNIT 8 NEUTRALIZATION TITRATION-I

Transcription

1 UNIT 8 NEUTRALIZATION TITRATION-I Structure 8.1 Introduction Objectives 8. Basic Concepts of Titrimetry, Primary and Secondary Standards 8.3 Titration Curves Titration of A Strong Acid Versus Strong Base Titration of Weak Acid Versus Strong Base Titration of a Weak Base Versus Strong Acid Titration of Weak Acid Versus Weak Base Titration of Sodium Carbonate Versus Strong Acid Titration of Polyprotic Acid Versus Strong Base 8. Theory of Indicators Ostwald s Theory Modern Quinoid Theory 8.5 Colour Change Range of an Indicator 8.6 Selection of Indicator and Indicator Error 8.7 Summary 8.8 Terminal Questions 8.9 Answers 8.1 INTRODUCTION The term titrimetric analysis refers to quantitative chemical analysis carried out by determining the volume of a solution of accurately known concentration, which is required to react quantitatively with the solution of the substance to be determined. The solution of accurately known strength is called the standard solution. The weight of the substance to be determined is calculated from the volume of the standard solution used and the known laws of stoichiometry. The standard solution is usually added from a burette. The process of adding the standard solution until the reaction is just complete is known as titration, and the substance to be determined is titrated. The point at which this occurs is called the equivalence point or the theoretical (or stoichiometric) end-point. The end point is judged usually, by the addition of an auxiliary reagent, known as indicator. The neutralization titrations include the titration of free bases, or those formed from the salts of weak acids by hydrolysis with standard acids (acidimetry), and the titration of acids or those formed by the hydrolysis of salts of weak bases, with standard base (alkalimetry). These reactions involve the combination of hydrogen and hydroxide ions to form water. This chapter describes the various types of acid base neutralization titrations, including the titration of strong acids or bases and weak acids or bases. Through a description of the theory of indicators, the selection of a suitable indicator for detecting the completion of a particular titration has been discussed. Objectives After studying the unit, you should be able to: State and explain the concept of titrimetry explain the nature of neutralization titration curves understand primary and secondary standards describe the theory of indicators select the appropriate indicator for a particular titration calculate the ph of acid, base and buffer solution. 51

2 Estimations Based On Kinetic and Acid-Base Equilibria Studies 8. BASIC CONCEPT OF TITRIMETRY, PRIMARY AND SECONDARY STANDARDS The acid base titration involves a neutralization reaction in which an acid is reacted with an equivalent amount of base. The titrant is always a strong acid or a strong base. The object of neutralization, say, an alkaline solution with a standard solution of an acid in the determination of the amount of acid which is exactly equivalent chemically to the amount of base present. The point at which this is reached is equivalence point or theoretical end point; an aqueous solution of the corresponding salt results. If both the acid and base are strong electrolytes, the resultant solution will be neutral and have a ph of 7; but if either the acid or base is a weak electrolyte, the salt will be hydrolyzed to a certain degree, and the solution at the equivalence point will be either slightly alkaline or slightly acidic. The exact ph of the solution at the end point can readily be calculated from the ionization constant of the weak acid or the weak base and the concentration of the solution. For any actual titration the correct end point will be characterized by a definite value of the hydrogen ion concentration of the solution, the value depending upon the nature of the acid and the base and the concentration of the solution. Primary Standard A primary standard is a substance which satisfies the following requirements: 1. It must be easy to obtain, to purify, to dry (preferably at o C) and to preserve in pure state.. The substance should remain unaltered during weighing i.e., it should not be hygroscopic, or oxidized by the air, or affected by carbon dioxide. 3. The substance should be capable of being tested for impurities by qualitative and other tests of known sensitivity.. It should have a high equivalent so that the weighing errors may be negligible. 5. The substance should be readily soluble under the experimental conditions. 6. The reaction with the standard solution should be stoichiometric and practically instantaneous. In practice, it is difficult to obtain a primary standard, and a compromise between the above ideal requirements is necessary. The commonly employed primary standards include sodium carbonate, sodium tetraborate, potassium hydrogen phthalate, constant-boiling-point hydrochloric acid, potassium hydrogen iodate, and benzoic acid. Secondary Standard A substance, which fulfills the requirement that it can be weighed accurately to provide a known amount of reactant but which is not a pure substance, is called secondary standard. It may be used for standardizations, and whose content of the active substance has been found by the comparison against a primary standard. SAQ 1 a) Mention any two requirements of a primary standard. b) How can secondary standards be used for standardizations? 5

3 8.3 TITRATION CURVES A titration curve is constructed by plotting ph of the solution during titration as ordinates and the amount of acid or base added as abscissa. These curves are useful to indicate equivalence point graphically. The change in ph in the neighborhood of the equivalence is of greatest importance, as it enables us to select an indicator, which will give the smallest error. The nature of titration curve depends on the ionization constants of acid and base employed in titration i.e., their strength. The principles of acid base equilibria are important for the construction and interpretation of titration curves in neutralization titrations Titration of A Strong Acid Versus Strong Base In the case of a strong acid versus strong base, both the titrant and analyte are completely ionized. An example is the titration of hydrochloric acid with sodium hydroxide. H Cl Na OH H O Na Cl The H and OH combine to form H O, and the other ions (Na and Cl ) remain unchanged, so the net result of neutralization is conversion of the HCl to a neutral solution of NaCl. The calculation of the titration curves involves computation of the ph for the concentration of the particular species at the various stages of the titration. The ph during neutralization processes will be calculated as follows: 1. Up to the equivalence point the ph of the solution is determined by the amount of the strong acid remaining present.. At the equivalence point ph is After passing the equivalence point the ph value is defined by the excess of the base. Before any NaOH is added to 1.0 M HCl solution, its ph is zero. As the titration proceeds, part of H is removed from the solution as water. So the concentration of H gradually decreases but this decrease is not significant, probably, due to the reason that the strong acids are good buffers at low ph. The results of calculation of ph during the titration of 100 cm 3 of HCl with NaOH of equal concentration are presented in Table 8.1. Fig. 8.1 is the graphical representation of these data. Fig. 8.1: Neutralisation curves of 100 cm 3 of HCl with NaOH of same concentration (calculated) 53

4 Estimations Based On Kinetic and Acid-Base Equilibria Studies After 50 cm 3 addition of NaOH solution, 50 cm 3 of un-neutralized acid will be present in a total volume of 150 cm 3. [H ] = 50 x 1/150 M = 0.33 M, or ph = 0.8 After 75 cm 3 addition of NaOH, [H ] = 5 1/175 M = 0.13 M, or ph = 0.8 After 90 cm 3 addition of NaOH, [H ] = 10 1/190 M = M, or ph = 1.3 After 99 cm 3 addition of NaOH, [H ] = 1 1/199 M = M, or ph =.3 After 99.9 cm 3 addition of NaOH, [H ] = 0.1 1/199.9 M = M, or ph = 3.3 After 100 cm 3 addition of base, the ph will change sharply to 7, i.e., the theoretical end point provided carbon dioxide is absent; the resulting solution is simply one of sodium chloride. Where the ph for the over titrations of 0.10 and 1.0 cm 3 corresponding to ph 10.7 and 11.7, respectively can be calculated as given below: After cm 3 addition of NaOH, [OH ] = 0.1/00.1 M = M, poh = 3.3, and ph = 10.7 With 101 cm 3 of base, [OH ] = 1/01 = M, poh =.3, and ph = 11.7 These results show that as the titration proceeds, initially the ph rises slowly, but between the addition of 99.9 and cm 3 of alkali, the ph of the solution rises from 3.3 to 10.7,.3 to 9.7 and 5.3 to 8.7 in case of 1 M, 0.1 M and M solutions respectively. Further addition of base does not cause a significant change in ph. The results of titration are presented in Table 8.1 for 1M, 0.1 M and 0.01 M solutions of acid and base, respectively. Table 8.1: ph during Titration of 100 cm 3 of HCl with NaOH of Equal Concentration NaOH added cm M solution ph 0.1 M solution ph 0.01 M solution ph 5

5 In the quantitative analysis we are especially interested in the change of ph near the equivalence point. This part is accordingly shown on a large scale in Fig. 8., on which are also indicated the colour change intervals of some of the common indicators. The magnitude of the break will depend on both the concentration of the acid and the concentration of the base. The reverse titration will be the mirror image of these titrations. Fig. 8.: Neutralisation curves of 100 cm 3 of HCl with NaOH of same concentration in vicinity of equivalence point (calculated) 8.3. Titration of Weak Acid Versus Strong Base We illustrate this case by the titration curve of 0.1 M acetic acid with 0.1 M sodium hydroxide as shown in Fig The neutralization reaction is: CH 3 COOH Na OH H O Na CH 3 COO The acetic acid, which is only few percent ionized, depending on the concentration, is neutralized to water and an equivalent amount of the salt, sodium acetate. Before the titration is started, we have 0.1 M CH 3 COOH. As soon as the titration is started some of the CH 3 COOH is converted to CH 3 COONa, and a buffer system is set up. As the titration proceeds, the ph slowly increases as the ratio [CH 3 COO ] / [CH 3 COOH] changes. At the mid point of titration, [CH 3 COO ] = [CH 3 COOH], and the ph is equal to pka. At the equivalence point we have a solution of CH 3 COONa. Since this is Bronsted base (it hydrolyzes), the ph at the equivalence point will be alkaline. The ph will depend on the concentration of CH 3 COONa. The greater the concentration, the higher the ph. As excess of NaOH is added beyond the equivalence point, the ionization of base CH 3 COO is suppressed to a negligible amount, and the ph is determined only by the concentration of excess OH. Therefore, the titration curve beyond the equivalence point follows that for the titration of a strong acid. 55

6 Estimations Based On Kinetic and Acid-Base Equilibria Studies Fig. 8.3: The ph titration curve of weak acid (CH 3 COOH) and strong base (NaOH) The slowly rising region before the equivalence point is called the buffer region. It is flattest at the midpoint, and so the buffer capacity is greatest at a ph corresponding to pk a. The buffering capacity also depends on the concentrations of CH 3 COOH and CH 3 COO, and the total buffering capacity increases as the concentration increases. In other words, the distance of the flat portion on either side of pk a will increase as [CH 3 COOH] and [CH 3 COO ] increase. For plotting the titration curve the ph values can be calculated as: 1. ph of weak acid can be calculated from the following equation: ph = ½ pk a ½ log [acid] (8.1) Strictly, [H ] = 0.5 ( (K a K a [acid] 0.5 ) K a ). Up to equivalence point ph of the solution is determined by the dissociation exponent of the weak acid and by the ratio of the concentration of free acid (HA) and titrated acid (A = salt) (buffer solution): [ salt] [ acid] ph = pk a log (8.) 3. The ph at the equivalence point is greater than 7 due to the alkaline hydrolysis of the resulting salt: ph = ½ pk w ½ pk a ½ log [salt] or, ph = 7 ½ pk a ½ pc (8.3) Concentration of the salt is c mol dm 3 to be consistent we use here pc = -log c.. After passing the equivalence point the excess of the base determines the ph of the solution as if the hydrolyzing salts were not present at all. The initial ph of 0.1 M acetic acid solution is computed from Eq. (8.1); the 56

7 dissociation of acid is relatively so small that it may be neglected in expressing the concentration of acetic acid. ph = ½ (-log ) ½ log 0.1 or, ph =.87 when 50 cm 3 of 0.1 M alkali has been added, [salt] = /150 = and [acid] = /150 = ph can be computed from Eq. (8.) ph = log 1.8 x log or, ph =.7 The ph of solution at the equivalence point can be calculated using Eq. (8.3). ph = 7 1/ log / log 5 10 = / (1.3) = 8.7 The ph values at other points on the titration curve can be similarly calculated. After the equivalence point has been passed, the solution contains excess of OH ions which will repress the hydrolysis of the salt; the ph may be assumed, with sufficient accuracy for our purpose, to be that due to the excess of base present, so that in this region the titration curve will almost coincide with that for 0.1 M hydrochloric acid. The results of titration are presented in Table 8.. The results for the titration of 100 cm 3 of a weaker acid (K a = ) with 0.1 M sodium hydroxide at the laboratory temperature are also included. Table 8.: of 100 cm 3 of 0.1 M acetic acid (K a = ) and of 100 cm 3 of 0.1 M HA (K a = ) with 0.1 M sodium hydroxide. 0.1 M NaOH used cm M acetic acid ph 0.1M-HA(K a = ) ph

8 Estimations Based On Kinetic and Acid-Base Equilibria Studies For 0.1 M acetic acid and 0.1 M sodium hydroxide, it is evident from the titration curve that neither methyl orange nor methyl red can be used as indicators. The equivalence point is at ph 8.7, and it is necessary to use an indicator with a ph range of slightly alkaline side, such as phenolpthalein, thymolpthalein, or thymol blue (ph range, as base, ). For the acid with K a = the equivalence point is at ph = 10.0, but here the rate of change of ph in the neighborhood of the stoichiometric point is very less pronounced, owing to considerable hydrolysis. Phenolphthalein will commence to change colour after 9 cm 3 of alkali have been added, and this change will occur to the equivalence point; thus the end point will not be sharp and the titration error will be appreciable. With thymolphthalein, however, the colour change covers the ph range ; this indicator may be used, the end-point will be more sharp than for phenolphthalein, but nevertheless somewhat gradual, and the titration error will be about 0. per cent. Acids having K a < cannot be satisfactorily titrated in 0.1 M solution with simple indicator. In general it may be stated that weak acids (K a > ) should be titrated with phenolphthalein, thymolphthalein, or thymol blue as indicators. Fig. 8.: Titration of 50 cm 3 of 0.1 M-H 3 PO with 0.1 M-KOH Fig. 8. shows the titration curves for 50 cm 3 of 0.1 M solutions of weak acids of different K a values titrated with 0.1 M KOH. The sharpness of the end point decreases as K a decreases. As in Fig. 8. sharpness will also decrease as the concentration decreases. Generally for macro titrations (ca. 0.1 M), acids with K a values of 10 6 can be titrated accurately with a visual indicator; and with suitable colour comparisons, those with K a values approaching 10 8 can be titrated with reasonable accuracy. A ph meter can be used to obtain better precision for the very weak acids by plotting the titration curve. Weaker acids can be titrated in nonaqueous solvents that do not possess the acidity or basicity of water. 58

9 8.3.3 Titration of a Weak Base Versus Strong Acid The titration of a weak base with a strong acid is similar to the above case, but the titration curves are reverse of those for a weak acid versus strong base. This neutralization titration can be illustrated by the titration curve for 100 cm 3 of 0.1 M ammonia with 0.1 M hydrochloric acid as shown in Fig The neutralization reaction is, NH 3 H Cl NH Cl Fig. 8.5: The ph titration curve of 100 cm 3 0.1M ammonia (NH OH) with 0.1M hydrochloric acid (HCl) At the beginning of the titration, we have 0.1 M NH 3 and the ph is calculated for a weak base. As soon as some acid is added, some of the NH 3 is converted into NH and the buffer region is formed. End point is judged by inflection point of the titration curve but the jump in ph is smaller than in case of strong base - strong acid titration. The ph values obtained in the course of titration can be calculated as given below: 1. The ph of the solution of a weak base: ph = pk w ½ pk b 1/ log [base] (8.) Strictly, [H ] = K w / (0.5 ( ( K K b [base]) 0.5 K b ). The ph upto the equivalence point: b [ base] [ salt] ph = pk pk log (8.5) w b 3. The ph of the equivalence point is lower than 7, due to the hydrolysis of resulting salt: ph = 1/ pk w 1/ pk b 1/ log [salt] or, ph = 7 1/ pk b 1/ pc (8.6) C and pc have the same meaning as used in Eq. (8.3).. After the attainment of equivalence point, the solution contains the excess of H ions, hydrolysis of the salt will be repressed, and the subsequent ph change may be assumed with sufficient accuracy for our purpose, to those due to the excess of acid present. 59

10 Estimations Based On Kinetic and Acid-Base Equilibria Studies Example Titration of 100 cm 3 of 0.1 M aqueous ammonia (K b = ) with 0.1 M HCl at the ambient temperature. The ph of the solution at the equivalence point is given by the Eq. (8.6): ph = 7 1/ pkb 1/ pc = 7.37 ½ (1.3) = 5.8 For the titration of weak base and strong acid, those indicators are generally used which change their colour in the acidic ph range. It is clear from the Fig. 8. that neither thymolphthalein nor phenolpthalein can be employed in the titration of 0.1M aqueous ammonia. The equivalence point is at ph 5.3, and it is necessary to use an indicator with a ph range on the slightly acid side (3 6.5), such as methyl orange, methyl red, bromophenol blue, or bromocresol green. The last named indicators are applicable for the titration of all weak bases (K b > ) with strong acids For the weak base (K b = ), bromophenol blue or methyl orange may be used; no sharp colour change will be obtained with bromo-cresol green or with methyl red, and the titration error will be considerable Titration of Weak Acid Versus Weak Base This can be illustrated by titration of 100 cm 3 of 0.1 M acetic acid (K a = ) with 0.1 M aqueous ammonia (K b = ). The ph at the equivalence point is given by: ph = 1/ pk w ½ pk a ½ pk b (8.7) = = 7.0 The neutralization curve upto the equivalence point is almost identical with that using 0.1 M sodium hydroxide as the base; beyond this point the titration is virtually the addition of 0.1 M aqueous ammonia solution to 0.1 M ammonium acetate solution and Eq. (8.5) is applicable to the calculation of the ph. The titration curve for the neutralization of 100 cm M acetic acid with 0.1 M aqueous ammonia at the laboratory temperature is shown in Fig The main feature of the curve is that the change of ph near equivalence point and, indeed, during the whole of the neutralization curve is very gradual. Fig. 8.6: Titration curve of weak base and weak acid 60

11 In this sort of titration, no sudden change in ph, and hence no sharp end point can be found with any simple indicator. A mixed indicator which exhibits a sharp colour change over a very limited ph range, may sometimes be found which is suitable. Thus for acetic acid ammonia solution titrations, neutral red methylene blue indicator may be used, but on the whole, it is best to avoid the use of indicators in titrations involving both a weak acid and a weak base Titration of Sodium Carbonate Versus Strong Acid Sodium carbonate is a Bronsted base that is as a primary standard for the standardization of strong acids. It hydrolyses in two steps: CO 3 H O HCO 3 K OH w KH 1 = K = =.1x10.. (8.8) b1 K a HCO 3 H O CO H O OH Kw KH 8 = K = =.3x10... (8.9) b K a 1 where K a and K 1 a refer to the K a values of H CO 3. HCO 3 is the conjugate acid of CO 3 and H CO 3 is the conjugate acid of HCO 3 and K b values are calculated for salts of weak acids and bases (i.e., from K a K b = K w ). A titration curve for Na CO 3 with HCl is shown in Fig. 8.7 (solid line). Even though K is considerably larger than the 10 6 required for a sharp end point, the ph break b 1 is decreased by the formation of CO beyond the first equivalence point. The second end point is not very sharp either, because K is smaller than Fortunately, this end point can be sharpened because the CO produced from the neutralization of HCO 3 is volatile and can be boiled out of the solution. This is described below. b Fig. 8.7: Titration of 100 cm 3 of 0.1 M-Na CO 3 with 0.1 M-HCl 61

12 Estimations Based On Kinetic and Acid-Base Equilibria Studies At the start of the titration, the ph is determined by the hydrolysis of the Bronsted base CO. After the titration is begun, part of the CO is converted to HCO 3 ; and 3 - CO /HCO 3 3 buffer region is established. At the first equivalence point, there remains a solution of HCO 3, and [H ] K K a 1 a 3. Beyond the first equivalence point, the HCO 3 is partially converted to H CO 3 (CO ) and a partial buffer region is established, the ph being established by [HCO 3 ]/[CO ]. The ph at the second equivalence point is determined by the concentration of the weak acid CO. Phenolphthalein is used to detect the first equivalence point, and methyl orange is used to detect the second one. Neither the point, however, is very sharp. In actual practice, the phenolphthalein end point is used only to get an approximation of where the second end point will occur; phenolphthalein is colourless beyond the first end point and does not interfere. The second equivalence point, which is used for accurate titrations, is normally not very accurate with methyl orange indicator because the gradual changes in colour of the methyl orange. This is caused by the gradual decrease in ph due to the HCO 3 /CO buffer system beyond the first end point. If beyond the first end point the solution is boiled after each addition of HCl to remove the CO from the solution, the buffer system of HCO 3 /CO would be removed, leaving only HCO 3 in the solution. This is both a weak acid and a weak base whose ph ( 8.3) is independent of concentration ([H ] = K K or[oh ] = K K. a 1 a b 1 b Then the ph would remain essentially constant until the equivalence point when we are left with a neutral solution of water and NaCl (ph = 7.0) Titration of Polyprotic Acid Versus Strong Base Diprotic acids can be titrated stepwise, if mixture of two acids with constants acid (H SO 3 ), K a 1 = and Ka 1 and K a Ka 1 10 x Ka K a, the solution behaves like a respectively. Thus for sulphurous = , it is evident that there will be a sharp change of ph near the first equivalence point, but for the second stage the change will be less pronounced, yet first sufficient for the use of, say thymolphthalein as indicator. For carbonic acid (H CO 3 ), K = and K = , only the first stage will be just discernible in the neutralization curve, the second stage is far too weak to exhibit any point of inflexion and there is no suitable indicator for direct titration. Now the titration curve for diprotic acid (H A) versus Na CO 3 is to be discussed. During titration upto first equivalance point, a solution of HA / H A buffer region is established. At the first equivalance point, a solution of HA exists, and 1 1 [H ] K K or ph = pk pk. Beyond this, a A / HA buffer exists; a 1 a a 1 a1 a and finally the second equivalance point, the ph is determined from the hydrolysis of A. For triprotic acid (H 3 PO ), the ionization is given below: a H 3 PO H H PO K a 1 = = [H ][HPO [H PO ] 3 - ]... (8.10) H PO H HPO K a = = - [H ][HPO ] [H PO ] (8.11) 6

13 HPO H 3 PO K a 3 = = 3- [H ][PO ] [HPO ] (8.1) The overall ionization constant is the product of the individual ionization constants: H 3 PO 3H PO [H ] [PO ] K a = Ka Ka Ka =.3 x 10 = (8.13) 1 3 [H PO ] Orthophosphoric acid will behave like a mixture of three monoprotic acids. The ph of the first equivalance point for 0.1 M H 3 PO with 0.1 M NaOH is given approximately by: pk 1 a pk.6 1 a (8.1) 1 = and that of the second equivalance point by: pk 1 a pk 9.7 a (8.15) 3 1 = In the very weak third stage, the curve is flat and no indicator is available, the third equivalance point may be computed approximately from the following equation: ph = 1 pk 1 pk 1 pc (8.16) w a3 = ½ (1.6) = 1.35 (Here the terms have their usual meaning) The experimental neutralization curve of 50 cm 3 of 0.1 M H 3 PO with KOH and suitable indicators are shown in Fig. 8.. SAQ a) Calculate the ph at 0, 10, 90, 100 and 110 % of the titration for the titration of 50 cm 3 of 0.5 M HCl with 0.5 M NaOH. b) Calculate the ph at 0, 5 and 50 cm 3 titrant in the titration of 50 cm 3 of 1 M acetic acid with 1 M NaOH. c) Calculate the concentration of OH and the ph of a solution that is 0. M in aqueous NH 3 and 0.1 M in NH Cl. 3 63

14 Estimations Based On Kinetic and Acid-Base Equilibria Studies d) What is the buffer region? 8. THEORY OF INDICATORS Indicators are the organic substances, the presence of very small amount of which indicates the termination of a chemical reaction by a change of colour. Indicators are of various types, e.g., acid-base indicators, redox indicators, adsorption indicator, etc. Acid-base indicators are the organic substances, which have one colour in acid solution while different colour in alkaline solution. The following theories have been put forward to explain the colour change of the acid base indicator Ostwald s Theory According to the Ostwald theory indicators are such weak acids ( HIn) or bases (InOH) whose colours are different from that of the indicator-ion formed by their dissociation. The equilibria in the aqueous solution may be written as: HIn H In InOH OH In Unionized ionized colour colour If the indicator is a free amine or substituted amine the equilibrium is: In H O OH HIn Indicator-acids HIn dissociate in aqueous solution as follows: HIn H In Applying the law of mass action to this dissociation from which - [H ][In ] = K [HIn] a (8.17) [HIn] H ] = K (8.18) [In ] [ a The actual colour of the indicator, which depends upon the ratio of the concentration of the ionized and unionized forms, is thus directly related to the hydrogen ion concentration. Eq. (8.18) may be written as [In ] ph = log pk a [HIn]... (8.19) In this equation [HIn] represents the concentration of the undissociated indicatormolecule whose colour is called acid colour while [In ] denotes the concentration of the indicator-anions, the colour of which is called alkaline colour. K a is the dissociation constant of the indicator- acid. The indicator base may be characterized similarly to the indicator acid InOH In OH 6

15 [In ] [OH ] = K [InOH] taking the ionic product of water into consideration [In ] K w = K [H ] [InOH] K [H ] = K w b b b [ In ] [ InOH] (8.0) (8.1) (8.) [InOH ] ph = pk w pk b log (8.3) [In ] where K w represents the ionic product of water, K b denotes the dissociation constant of the indicator base, the colour of which is the alkaline colour ; the acid colour is due to the In ions. 8.. Modern Quinoid Theory According to the modern quinoid theory an acid-base indicator is a dynamic equilibrium mixture of two alternative tautomeric forms; ordinarily one form is benzenoid while the other is quinoid. Out of these one form exists in the acidic solution, while the other in alkaline solution. Change in ph causes the transition of benzenoid form to quinoid form and vise versa and consequently a change in colour. The colour changes in case of methyl orange and phenolphthalein are given below: Methyl Orange Na O 3 S N=N N(CH 3 ) Yellow Benzenoid form (in bases) H H Na O 3 S N N= H N (CH 3 ) Phenolphthalein OH Red Quinoid form (in acids) O OH H C O HO C H C=O COO Colourless, benzenoid form (in acid) Red, Quinoidform (in alkali) 65

16 Estimations Based On Kinetic and Acid-Base Equilibria Studies SAQ 3 Why the quinoid form of the indicator is coloured or darker than benzenoid form? 8.5 COLOUR CHANGE RANGE OF AN INDICATOR A large number of acid-base indicators are available, which possess different colour according to the hydrogen-ion concentration in the solution. The important characteristics of these indicators is that the change from a predominantly acid colour to predominantly alkaline colour is not sudden and abrupt, but takes place within a small interval of ph (generally about two ph units) termed the colour-change interval of the indicator. In this ph interval the indicator shows mixed colours of different shades of the acid and alkaline colours, i.e. the colour intensity of one colour indicators increases gradually. The acidic red colour of methyl orange and methyl red is more sensitively perceivable beside the alkaline yellow colour, since the colour intensity of red form is greater. The position of the colour-change interval in the ph scale varies widely with different indicators. For most acid-base titrations we can therefore select an indicator, which exhibits a distinct colour change at a ph close to that obtained at the equivalence point. Table 8.3 summarizes a selected list of indicators suitable for neutralization titration. Table 8.3: The ph Transition Ranges and Colours of Some Indicators Indicator ph range Colour in Colour in acid solution alkaline solution pkʹin Brilliant cresylblue (acid) Red-orange Blue - Cresol red (acid) Red Yellow - Thymol blue (acid) Red Yellow 1.7 m-cresol purple Red Yellow - Bromo-phenol blue Yellow Blue.1 Methyl Yellow.9-.0 Red Yellow 3.3 Ethyl Orange Red Orange - Methyl orange Red Orange 3.7 Congo red Blue Red - Bromo-cresol green Yellow Blue.7 Methyl red.-6.3 Red Yellow 5.0 Ethyl red Red Orange - Propyl red Red Yellow - Chlorophenol red.8-6. Yellow Red 6.1 -Nitrophenol Colourless Yellow 7.1 Bromo-cresol purple Yellow Purple 6.1 Bromo-phenol red Yellow Red - Bromo-thymol blue Yellow Blue 7.1 Neutral red Red Orange - Phenol red Yellow Red 7.8 Cresol-red (base) Yellow Red 8. m-cresol purple Yellow Purple - Thymol blue (base) Yellow Blue 8.9 o-cresol-phthalein Colourless Red - Phenol-phthalein Colourless Red 9.6 Thymolphthalein Colourless Blue 9.3 Alizarin yellow R Yellow Orange red - Brilliant crystal blue (base) Blue Yellow - Tropaeolin O Yellow Orange - 66

17 SAQ What is the usual ph range for the colour change of an indicator at the end point? 8.6 SELECTION OF INDICATOR AND INDICATOR ERROR The primary consideration in choosing an indicator for a given titration is that the indicator should change, that is the end point should occur, within the required increment, v of the equivalence point. As a general rule it may be stated that for a titration to be feasible there should be change of approximately two units of ph at or near the stochiometric point produced by the addition of small volume of the reagent. The ph range at either side of the equivalence point (0.1-1 cm 3 ) may be calculated, and the difference will indicate whether the change is large enough to permit a sharp end point to be observed. Alternatively, the ph change on the both sides of the equivalence point is noted from the neutralization curve determined by potentiometric titration. If the ph change is satisfactory an indicator should be selected that changes its colour at or near the equivalence point. Next, the molar absorptivity of the indicator in its two forms should be known in order that the amount of indicator required to give a readily observable change can be determined, and this amount should be small to avoid consumption of titrant. Finally the personal preference to the operator can be consulted as a selection of colours, etc., most studied to his vision. Yet, again it should be added that the indicator reaction should be fast. Indicator Error The indicator error follows from the fact, that the indicator itself will consume a certain amount of the standard solution. The amount of this consumption of standard solution depends first of all upon the nature of indicator and its concentration; whether it is alkaline for instance, etc. If c t. v t >> c ind. v ind where c t and c ind are the concentrations of titrant and the indicator solution; v t and v ind are the volumes used of titrant and indicator, respectively, then the indicator error will be negligible. SAQ 5 What is the indicator error? 8.7 SUMMARY In this unit we have described the basic concept of titrimetry. The neutralization curves for strong acid with strong base, weak acid with strong base, weak base with strong acid, weak acid with weak base, sodium carbonate with strong acid, and polyprotic acid with strong base have been illustrated. The nature of titration curves 67

18 Estimations Based On Kinetic and Acid-Base Equilibria Studies has been explained. The calculations of ph of the solution before titration, at the equivalence point, and after the equivalence point have been illustrated. The theory of indicators and the selection of the appropriate indicator for a particular titration have been discussed. 8.8 TERMINAL QUESTIONS 1. Only strong acid or base is used as the titrant, why?. Calculate the [OH ], poh, ph and percent ionization for 0. M aqueous NH Calculate the concentration of the species in a 0.1 M H SO. K = When the mixtures of acids (or bases) can be titrated stepwise? 5. Calculate the ph of a solution prepared by adding 5 cm 3 of 0.10 M NaOH to 30 cm 3 of 0.0 M acetic acid. 6. Write the criteria for choosing an indicator. 8.9 ANSWERS Self Assessment Questions 1. a) A primary standard is a substance which satisfies the following requirements: It must be easy to obtain, to purify, to dry ( preferably at o C) and to preserve in pure state. The substance should remain unaltered during weighing i.e., it should not be hygroscopic, or oxidized by the air, or affected by carbon dioxide. The substance should be capable of being tested for impurities by qualitative and other tests of known sensitivity. b) Secondary standards may be used for standardizations by finding the content of the active substance comparing against a primary standard.. a) 0.30, 0.39, 1.58, 7.00, and 1.8 b).37,.7, and 9.07 c) [OH ] = 3.6 x 10 5 M and poh = 9.56 d) The slowly rising region before the equivalence point is called the buffer region. It is flattest at the midpoint, and so the buffer capacity is greatest at a ph corresponding to pk a. 3. The increased per cent conjugation in the quinoid form of the indicator results the shifting of the λ max from a shorter to a greater wave length (i.e., from the ultra violet region to the visible region).. The point at which the colour change for the indicator occurs in a titration is called end point. Typically, colour changes occur over a range of 1.5 to Section 8.6, nd Para. 68

19 Terminal Questions 1. The magnitude of the break at the equivalance point is significant and the end point can be obtained with greater accuracy.. [OH ] = 1.9 x 10 3 M, poh =.7, ph = 11.8 and % ionization = 0.95% ionized. 3. In first step, complete ionization of H SO is complete. H SO H O H O HSO 0.10 M 0.10 M 0.10 M 3 In second step, ionization is not complete. HSO HO [H3O ] [SO ] H 3O SO and K = =1. 10 [HSO ] Let x = HSO ] that ionizes. Therefore at equlibrium [ HSO HO H 3O SO (0.1-x) M (0.10 x )M xm I step II step [ H3O ] [SO ] (0.10 x) ( x) = = 1. [HSO ] (0.10 x) K = 10 x cannot be ignored because K is too large [H 3 O ] = [ SO ] = 0.01M. The concentrations of species in 0.1 M H SO are: [H SO ] 0.0 M ; HSO = (0.10-x) M = 0.09 M; [ SO [H 3 O ] = (0.10 x) M =0.11M [OH ] = K w / [H 3 O ] = /0.11 = ] = 0.01 M. There should be an appreciable difference in their strength and one acid or base should be at least 10 times weaker than the other to titrate separately. 5. CH 3 COOH NaOH CH 3 COONa H O m mol of CH 3 COOH formed = m mol of NaOH added =.5 m mol Unneutralized CH 3 COOH = = 3.5 m mol ph=.76 log.5/3.5 = Choose an indicator with a pk a near the equivalence. 69