Inorganic Pharmaceutical Analysis. Pharmacochemistry Research Group School of Pharmacy

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1 ACID BASE TITRATION An application method of Inorganic Pharmaceutical Analysis Lecturer : Dr. Tutus Gusdinar Pharmacochemistry Research Group School of Pharmacy INSTITUT TEKNOLOGI BANDUNG

2 The application of neutralization reaction Neutralization reaction could be used in determination of either natural or tried acid/basic analytes. Water is commonly used as a good solvent because of its cheap cost, easy to be obtained and prepared, non toxic, and low coeficient of expansion (Aquous Titration). Severalanalytescannotbetitrated analytes titrated inwater water, caused of its low solubility or too weak acid/base properties, these should be titrated in non aquaous solvent (Non Aquaous Titration).

3 Neutralization titration reagents Standard d acid solution o should oudbe standardized with a Primary Standard Bases suchas Na carbonate, TRIS atau THAM (tris hydroximethyl aminomethane), Natt tetraborate, t Mercuricoxide. id Standard basic solution (attention CO 2 effect to distilled water) should be standardized with a Primary Standard Acid such as KH Phtalate, Benzoic acid, Sulphamic acid, KH iodate, Sulphosalicilic li ili acid.

4 Sulfuric acid molecular model

5 ACID BASE INDICATORS

6 What is an acid base indicator? An acid base indicator is a weak acid or a weak base. The undissociated form of the indicator is a different color than the iogenic form of the indicator. An Indicator does not change color from pure acid to pure alkaline at specific hydrogen ion concentration, but rather, color change occurs over a range of hydrogen ion concentrations. This range is termed the color change interval. It is expressed as a ph range.

7 How is an indicator used? Weak acids are titrated in the presence of indicators which change under slightly alkaline conditions. Weak kbases should ldbe titrated t in the presence of indicators which change under slightly acidic conditions.

8 What are some common acid base indicators? Several acid base indicators are listed below, some more than once if they can be used over multiple ph ranges. Quantity of indicator in aqueous (aq.) or alcohol (alc.) solution is specified. Ti Tried and true d indicators include: thymol blue, tropeolin OO, methyl yellow, methyl orange, bromphenol blue, bromcresol green, methyl red, bromthymol blue, phenol red, neutral red, phenolphthalein, thymolphthalein, alizarin yellow, tropeolin O, nitramine, and trinitroben zoic acid. Data in this table are for sodium salts of thymol blue, bromphenol blue, tetrabromphenol blue, bromcresol green, methyl red, bromthymol blue, phenol red, and cresol red.

9 Acidic : HIn + H 2 O H 3 O + + In - Basic : In - + H 2 O HIn + OH - [H 3 O + ][In - ] [HIn] Ka = ph = pka - log [HIn] [In - ] End point colour depends to the most dominant concentration of such indicator form. Ex : if HIn is red and In - is yellow, then : At low ph in which [HIn] is dominant, the ratio of 10/1 (red). At high ph in which [In - ] is dominant, the ratio of 1/10 (yellow). At medium ph in which [HIn] = [In - ], the ratio of 1 (orange).

10 Indicator ph range Yll Yellow coloured solution : ph yellow = pka + log 10/1 = 5 +1 = 6 Red coloured solution : ph red = pka + log 1/10 = 5 11 = 4 ΔpH = ph yellow ph red = 2 ph range = 4 6

11 PHENOLPHTALEIN H 2 In HIn - HIn In -2

12 The first useful theory of indicator action was suggested by W. 0stwald based upon the concept that indicators in general use are very weak organic acids or bases. The simple Ostwald theory of the colour change of indicators has been revised, and the colour changes are believed to be due to structural t changes, including the production of quinonoid and resonance forms; these may be illustrated by reference to phenolphthalein, the changes of which are characteristic of all phthalein indicators. In the presence of dilute alkali the lactone ring in (I) opens to yield (II), and the triphenylcarbinol structure (II) undergoes loss of water to produce the resonating ion (III) which h is red. If phenolphthalein hth l is treated t with excess of concentrated alcoholic alkali the red colour first produced disappears owing to the formation of (IV).

13 The chemical structure change of phenolphtalein indicator

14 PHENOL RED H 2 In + (RED) HIn (YELLOW) In - (RED)

15 METHYL ORANGE ( HELIANTHINE ) HIn + (RED) In (YELLOW)

16 Acid Base indicators list NAME ph range pka (μ = 0,1 M) Thymol Blue COLOUR CHANGE Red Yellow Yellow Blue TYPE Acid Methyl Yellow Red Orange Basic Mehtyl Orange * Red Orange Basic Bromcresol Green Yellow Blue Acid Methyl Red * Red Yellow Basic Bromcresol Violet Yellow Violet Acid Bromthymol Blue Yellow Blue Acid Phenol Red Yellow Red Acid Cresol Violet Yellow Violet Acid Phenolphtalein Colourless Red Acid Thymolphtalein Colourless Blue Acid Alizarin Yellow Yellow Violet Basic

17 1 M 0.1 M 0.01 M

18 ALCALIMETRY

19 ACIDIMETRY

20 Mixed Indicator When sharp indicator colour change could not be obtained at the end point, a mixture of two indicators or an indicator mixture could be used. INDICATOR COLOUR CHANGE Methyl orange 1 gram violet grey green Indigo carmine 2.5 gram (acid) ph=4 (basic) Dissolved in 1 liter of water Bromcresol Green 01% part red green Methyl Red 0.1 % 2 part (acid) ph=5.1 (basic) Phenolphtalein 01% part green pale blue violet Methylene Green 0.1 % 2 part (acid) ph=8.8 (basic, ph>9) Cresol Red 0.1 % 1 part yellow pink violet Thymol Blue 0.1 % 3 part (acid) ph=8.2 (basic, ph>8.4) Cresol Red 0.1 % 16 ml MthlRd Methyl Red 01% ml Methylene Blue 0.2 % 4 ml green greenish grey greenish violet violet ph=8.85 ph=8.35 ph=8.6 ph=8.8

21 Carbonate Titration CO H + 3 O HCO 3 + H 2 O pka 2 = 6.34 HCO 3 + H 3 O + H 2 CO 3 + H 2 O pka 1 = By ΔpKa = 4.02 unit the equivalent point might be sharp, but Ka 1 is too low and a sharp equivalent point could not be obtained. Carbonate titration should be titrated with a strong acid using phenolphtalein indicator acid (ph = ) : ph NaHCO3 = ½ [pka 1 + pka 2 ] = 8.35 Methyl orange (ph = ) at the second equivalent point TE 2, saturated solution of CO 2 has ph = 3.9. The CO 2 gaz could be removed by these two techniques : Neutralization of sample using methyl orange indicator, or Removing CO 2 by boiling process of distilled water.

22 Titration ti curve of carbonate Phenolphtalein ph Methyl Orange ml HCl

23 Titration of Carbonate Bicarbonate Mixture First end point (phenolphtalein) : completed neutralization of NaOH, Na 2 CO 3 is half neutralized, HCO 3 has not reacted yet. Second end point t( (methyl orange) : all of HCO 3 was neutralized, small drops of titrant (HCl) can change the ph from 8 to 4 (could be corrected withindicatorblanco). indicator The mixture of NaOH + NaHCO 3 solution could not be prepared because of : HCO 3 + OH CO H 2 O The reaction results should be a mixture of HCO 3 + CO 2 3 or the only CO 2 3, depends on the relative amount of the compound content in a sample.

24 Titration curve of a mixture of carbonate + Bicarbonate OH - + H 3 O + H 2 O CO H 3 O + HCO H 2 O ph V 1 HCO H 3 O + H 2 CO 3 + H 2 O V ml HCl

25 Titration of mixed two acids Like as diprotic acid : [HX] initial = [HY] initial If HX (with its Ka 1 ) is strong acid and HY (with its Ka 2 ) is weak acid, then a feasible titration i can occured at pka 1 pka 2 > 4 unit For unequal initial concentration, first equivalent point could be calculated by such following steps : 1) Charge balance [Na + ] + [H 3 O + ] = [OH ] + [X ] + [Y ] 2) [Na + ] = acid formal concentration = [HX] + [X ] 3) 1)+2) : [H 3 O + ] = [OH ] + [Y ] [HX] 4) Substitute [OH ], [Y ] and [HX] from Kw, Ka 1, Ka 2 : [H 3 O + ] = Kw/ [H 3 O + ] + Ka 2 [HY]/[H 3 O + ] [H 3 O + ][X ]/Ka 1 5) [H 3 O + ] = {Ka 1 Kw + Ka 1 Ka 2 [HY]}/Ka 1 +[X ] 6) If Ka 2 [HY] >>>Kw and [X ] >>Ka 1 then [H 3 O + ] = Ka 1 Ka 2 [HY]/[X ] 7) ph = ½ (pka 1 + pka 2 ) ½ log [HY]/[X ]

26 Titration of mixed HCl + HAc HCl titrated first, ph is not influenced by H 3 O + from HAc [Le Chatelier principle p : excess proton will pressure weak acid dissociation]. This assumption is less valid at near the equivalent point, caused of excess proton concentration increases. At first equivalent point tthe HCl has all titrated t and theph depends d only to HAc dissociation. After first equivalent point HAc titrationwillstart start. Inthe next curve this first equivalent point is not clear because of lower (not enough) ΔpH/ΔV. While the ph of HAc M is about 3. Because of ph< 4 this titration i is not feasible (by using indicator). In the next step the weak acid titrated by strong base, the titration will be feasible. Example: Titration of 50 ml mixture of HCl 0.10 M and HAc 0.10 M with a solution of NaOH 0.10M.

27 Titration curve of acids mixture 50 ml of HCl 0,10 M and HAc 0.10 M (Ka= ) titrated with NaOH 0.20 M ph 50 ml of HCl 0,10 M and HX 0.10 M (Ka= ) titrated with NaOH 0.20 M ml NaOH

28 TITRATION ERROR A feasible titration can be performed in a complete reaction at the end point, resulting a sharp ph depletion (vertical curve). A complete reaction should be obtained in a high value of K (equilibrium constant), great change of ph at near to end point, to obtain easily a high precision of the end point.

29 Titration of strong acid with strong base gives a very high K value : H O + OH H 2 O K = 1/Kw = 10 A high ΔpH occurs at equivalent point, i.e unit ph for a Δ V = 0.10 ml. At this high ΔpH any indicator could be used for obtaining a high precision of titration (ppm). This is a feasible titration. But it is difficult to calculate exactly the K value for a feasible titration, because of the influence of analyte and titrant concentrations to the ΔpH. In certain condition, titration could be perform without high precision.

30 It is predicted about 99.9% and 99.99% of analyte changes to its reaction product at the equivalent point. At this condition the K value could be predicted. There is a limited ability of human eyes to observe indicator colour change at the equivalent point, a few drops of titrant could change ph of 1 2 unit. Example : Titration of 50 ml HA 0.10 M with a strong base of 0.1 M. Calculate K minimum at an addition of ml titran (complete reaction), when addition of 2 drops (0,10 ml) of titrant after equivalent point changes ph at units. What is K minimum at ph change of 1 unit?

31 Influence of analyte and titrant concentration : ΔpH H decreases if the concentration ti of analyte lt and titrant decrease. Example : Weak acid titration : low of Ka, high of ph equivalent point and low of ΔpH. Increasing [HA] will decrease ΔpH. Increasing titrant volume should increase titration error. (resulting a smaller value than true end point). If [HA] is titrated at a smaller initial volume, then ΔpH increases, caused by little excessive titrant volume. If [titrant] increases then ΔpH increases, this diminues titrant Volume and increases titration error (higher than true end point).

32 In general: Precision (at ppm) could be produced by titration of weak acid/base solution of 0.05 M (dissociation constant of 1x10 6 M) with a titrant solution of M (K = 1x10 8). Salts of weak acid (Bronsted base) could be feasibly titrated by strong acid when its conjugated acid is too weak.

33 Example : An acid HA (Ka = 1x ) is too weak for a titration with a base which has dissociation constant of its conjugate base A = 1x10 5, because of Ka x Kb = 1x A could be titrated by strong acid. Same case for a weak base and its salt.

34 Titration error is the difference of reagent amount used between end point and equivalent point, in % oro/ooof o/oo of reacted compound equivalent quantity. Titration of strong acid of 0.01 N with a strong base gives an error of % at ph 5 or 9; but this error will be % at ph 6 or 8. Lower titration error should be obtained in a titration of carboxilic acid (Ka > 10 5 ) with a strong base. If Ca = acid analytical concentration and Cb = standad base analytical concentration, then at the equivalent point Ca = Cb, at the other points : Ca Cb = + or as titration error. During titration : Ca = [HA] + [A ] and Cb = [A ] + [OH ] [H + ] If weak base solution (near equivalent point) and [H + ] are neglected, hence Ca = Cb = [A ] + [HA] = [A ] + [OH ] [HA] = [OH ]

35 At equivalent point a pure cmpound NaA dissociates A + H 2 O HA + OH If acetic acid titration (Ka = 1.8 x ) without a significant volume Change, error at end point + 1 from equivelent point, At the equivalent point [HA] = [OH ] = 7.5 x 10 6 M ph = 8.8 If equivalent point occured at ph = 9.88 (or ph = 9.9) hence [HA] = 7 x 10 7 M or [OH ] = 76 x 10 5 Titration error = Cb Ca = {[OH ] [HA]}/Ca x 100% = 0.07%. If colour change at the equivalent point is clear and sharp, and the indicator is very good, titration error could be neglected.

36 Titration of Heavy Metal Cations An aquaous solution of heavy metal is a base (Bronsted), titrated witha strong base form an insoluble saltbase. Al OH Al(OH) 3 white precipitate i Cu OH Cu(OH) 2 blue precipitate The end point could be obtained before equivalent point caused of the early precipitated saltbase base. Itis beter to performe this withan indirect titration technique.

37 Titration of Borate (Borax) A borate (borax) dissolved in water could form a half neutralized boric acid Na 2 B 4 O H 2 O 2 H 3 BO H 2 BO Na + 2 H 2 BO H + 2 H 3 BO 3 A sharp end point should be obtained, this is an appropriate primary standard base for a standardization of HCl.

38 Titration of phosphoric acid Titration of phosphoric acid with NaOH is a mono/di protic (not triprotic) Eq point 1 : ph = ½ (pka 1 + pka 2 ) = 4.66 Indicator bromcresol green or methyl yellow. End point in detected using pure NaH 2 PO 4 as control compound. Eq point 2 : ph = ½ (pka 2 + pka 3 ) = 9.7 Indicator phenolphtalein orthymolblue blue, colour change occures in basic solution before equivalent point. Indicator timolphtalein would be better because of ph = 9.6 Third ionization product of phosphoric acid has Ka = 5 x Solution ofna 3 PO 4 is strong basic andeq. Point 3 is never obtained, except trivalent phosphate ion is removed such by CaCl 2 precipitation after Eq.Point 2 2 Na 2 HPO CaCl 2 Ca 3 (PO 4 ) NaCl + 2 HCl

39 Titration of carbonic acid (hydrated CO 2 ) As a diprotic acid : ph = ½ (pka1 + pka2) = 8.40 At Eq Point 1 the CO 2 could be titrated as monoprotic acid with a solution of NaOH using phenolphtalein or thymol blue (or its mixture), colour change is not sharp, needs control solution (pure NaHCO 3 + indicator at same quantity as for sample). Second ionization product, as a weak diprotic acid is too week for a direct titration. Carbonate ion has to be removed (precipitation) by addition of excess Ba(OH) 2 H 2 CO 3 + Ba(OH) 2 NaHCO 3 + Ba(OH) 2 BaCO 3 + 2H 2 O BaCO 3 + NaOH + H 2 O

40 Back titration with an acid standard solution could use phenolphtalein or thymol blue as the indicator without any filtering process. BaCO 3. NaHCO 3 Na 2 CO 3 H 2 CO 3 + NaHCO 3 ph NaHCO 3 + Na 2 CO ml NaOH 0.1 N Titration curve of 100 ml of carbonic acid 0.05 M with NaOH 0.1N

41 Summary Choice of indicator

42 Strong acid and strong base. For 0.1 M or more concentrated solutions, any indicator may be used which has a range between the limits ph 4.5 and ph 9.5. With 0.01 M solutions, the ph range is somewhat smaller ( ). If carbon dioxide is present, either the solution should be boiled while still acid and the solution titrated when cold, or an indicator with a range below ph 5 should be employed.

43 Weak acid and a strong base. The ph at the equivalence point is calculated from the equation: ph = ½ pkw + ½ pka ½ pc The ph range for acids with Ka > 10-5 is ; for weaker acids (Ka >10-6 ) the range is reduced (8-10). The ph range will cover most of the examples likely to be encountered; this permits the use of thymol blue, thymolphthalein, or phenolphthalein. hth l

44 Weak base and strong acid. The ph at the equivalence point is computed from the equation: ph = ½ pkw ½ pkb + ½ pc The ph range for bases with Kb > 10-5 is 3-7 7, and for weaker bases (Kb > 10-6 ) 3-5. Suitable indicators will be methyl red, methyl orange, methyl yellow, bromocresol green, and bromophenol blue.

45 Weak acid and weak base. There is no sharp rise in the neutralisation curve and, generally, no simple indicator can be used. The titration should therefore be avoided, if possible. The approximate ph at the equivalence point can be computed from the equation : ph = ½ pkw + ½ pka ½ pkb It is sometimes possible to employ a mixed indicator which exhibits a colour change over a very limited it ph range, for example, neutral red-methylene blue for dilute ammonia solution and acetic (ethanoic) acid.

46 Polyprotic acids (or mixtures of acids, with dissociation constants K 1, K 2, and K 3 ) and strong bases. The first stoichiometric end point is given approximately by ph = ½ (pk 1 + pk 2 ) The second stoichiometric end point is given approximately by ph = ½ (pk 2 + pk 3 )

47 Anion of a weak acid titrated with a strong acid. The ph at the equivalence point is given by ph = ½ pkw ½ pka ½ pc Cation of a weak base titrated with a strong base. The ph at the stoichiometric end point is given by ph = ½ pkw ½ pkb ½ pc

48 As a general rule, wherever an indicator does not give a sharp end point, it is advisable to prepare an equal volume of a comparison solution containing the same quantity of indicator and of the final products and other components of the titration i as in the solution under test, and to titrate to the colour shade thus obtained. In cases where it proves impossible to find a suitable indicator (and this will occur when dealing with strongly coloured solutions) then titration may be possible by an electrometric method such as conductimetric, potentiometric or amperometric titration. In some instances, spectrophotometric titration may be feasible. It should also be noted that if it is possible to work in a non-aqueous solution rather than in water, then acidic and basic properties may be altered according to the solvent chosen, and titrations which are difficult in aqueous solution may then become easy to perform. This procedure is widely used for the analysis of organic materials but is of very limited application with inorganic substances.

49 End

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