Acid-Base Titrations - Pharmaceutical Analytical Chemistry I PDF

Summary

These notes cover the fundamental concepts of acid-base titrations within the context of pharmaceutical analytical chemistry. Topics include titration methods, standard solutions, and calculations. The document also touches on various pharmaceutical applications of analytical chemistry.

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Pharmaceutical Analytical Chemistry I (PC 101 & CPC 101) Level I Students Assoc. Prof. Mohamed El-Awady Acid-Base Titrations General Objectives of the Course After this course, you should know: Introduction - Cl...

Pharmaceutical Analytical Chemistry I (PC 101 & CPC 101) Level I Students Assoc. Prof. Mohamed El-Awady Acid-Base Titrations General Objectives of the Course After this course, you should know: Introduction - Classification of Quantitative Analysis. Volumetric Analysis (Titrimetric Analysis) (Titration) - Requirements of reactions in titrations. Standard Solutions [Primary & Secondary]. Methods of expressing concentration of standard solution - Important mathematical equations. Equivalence point and end point - Detection of the end point. Steps for the determination of a sample using titration. Types of titration methods: Direct / Back (Residual) / Displacement (Substitution). Acid-base equilibrium - pH of acids, bases and salts. Buffer solutions: (Definition, composition, action, pH, capacity). Acid-base indicators: (Definition, types). Titration curves & choice of acid-base indicators. Applications of acid-base titrations including determination of : [Acids (weak & strong), bases (weak & strong), salts, mixture of acids, mixture of bases, Kjeldahl method for determination of proteins and nitrogen compounds, and biphasic titration]. 2 ‫‪Acid-Base Titrations‬‬ ‫ هل دراستك للكيمياء التحليلية لها فائدة؟‬ ‫ ما هي التطبيقات الصيدلية للكيمياء التحليلية؟‬ ‫ ما هي عﻼقة الكيمياء التحليلية بعلوم الصيدلة اﻷخرى؟‬ ‫ هل لقسم الكيمياء التحليلية الصيدلية دور في خدمة المجتمع؟‬ ‫)نظري و عملي( التي يتم تدريسها‬ ‫ هل المناهج‬ ‫متواكبة مع متطلبات سوق العمل؟‬ ‫ هل في مصر أماكن شـغل للصيدلي غير اﻷماكن التقليدية‬ ‫)الصيدليـات و الدعايـا( و مرتبطة بدراستك للكيمياء التحليلية؟‬ ‫ كيف تؤهل نفسك للعمل في هذه اﻷماكن الغير تقليدية؟‬ ‫‪3‬‬ Acid-Base Titrations Pharmaceutical Applications of Analytical Chemistry 1) Quality Control in Pharmaceutical Companies: Raw Production Finished Material Patient ex : Aspirin Area Product Quality Quality Quality Quality Control Control Control Control 2) Water Analysis. 3) Food Analysis. 4) Forensic Analysis. 5) Cosmetics Industry. 6) Detergents & Chemicals Industry. 4 Acid-Base Titrations Career Opportunities Related to Analytical Chemistry 1) Quality control dept. in pharmaceutical co. ‫ات اﻷدو ة‬ ‫رقا ة الجودة‬ 2) Academy of scientific research & technology. ‫و التكنولوج ا‬ ‫أ اد م ة ال حث العل‬ 3) National research center. ‫المركز القو لل حوث‬ 4) National organization for drug control & research. ‫الهيئة القوم ة للرقا ة و ال حوث الدوائ ة‬ 5) Forensic medicine agency. ‫مصلحة الطب ال‬ 6) Quality control dept. in cosmetics industry. ‫ات التجم ل‬ ‫ات مستح‬ ‫رقا ة الجودة‬ 7) Quality control dept. in detergents & chemicals industry. ‫ات المنظفات و ال ماو ات‬ ‫رقا ة الجودة‬ 8) Quality control dept. in food industry. ‫رقا ة الجودة مصانع اﻷغذ ة‬ 9) Water purification & desalination stations. ‫محطات تنق ة و تحل ة الم اه‬ 10) Community pharmacist. ‫ص د المجتمع‬ Acid-Base Titrations Analysis of Anions & Cations Qualitative Analysis (= What?) Instrumental Analysis Pharm. Analytical depends on Volumetric Analysis Chemistry measuring (Titration) volume Acid-base titration (aq. & non-aq.) Precipitation titration Complex-formation titration Quantitative Analysis Redox titration or Determination or Estimation or Assay (=How much?) depends on Gravimetric Analysis measuring weight depends on Instrumental Analysis instruments 6 Acid-Base Titrations Introduction Pharmaceutical Analytical Chemistry: It is the branch of analytical chemistry that uses the general theories of analytical chemistry in the analysis of pharmaceutical compounds. Qualitative Analysis: It establishes the chemical identity of the species present in the sample. (=What?) Quantitative Analysis: It establishes the amount (quantity) of the species present in the sample. (=How much?) The following terms are commonly used to express the meaning of quantitative analysis: Quantitative Analysis = Determination = Assay = Estimation 7 Acid-Base Titrations Volumetric Analysis (= Titrimetric Analysis = Titration) It is a method of quantitative analysis that depends on measuring the volume of a solution of known concentration (standard solution) needed to react completely with a sample of unknown concentration. It has different types based on the type of reaction involved in the titration including: acid-base, precipitation, complex-formation & redox titrations. In its simplest form, the titration is done by placing the standard solution (Titrant) in the burette and the sample of unknown concentration is placed in a conical flask. (N.B. Analyte is the component of the sample that is determined, and it is sometimes called Titrand) 8 Acid-Base Titrations Titration 9 Acid-Base Titrations Titration Video: Running a titration analysis By Royal Society of Chemistry https://www.youtube.com/watch?v=RI14t0R1wMY 10 Acid-Base Titrations Requirements of reactions in titrations 1. The reaction must be rapid. NaOH + HCl NaCl + H2O 2. The reaction must proceed toward completion. 3. The reaction must be represented in the form of a balanced chemical equation. 4. The reaction must have an available accurate method for detecting the end point. 11 Acid-Base Titrations Standard Solutions Standard solutions are solutions of known concentration and composition. Standard solutions are classified into 2 types: (1) Primary Standard It is a solution of exactly known concentration and composition and its concentration remains constant for a long period. It can be directly prepared by weighing an exact weight of the primary standard substance in a measuring (volumetric) flask and dissolving it in a proper solvent. 12 Acid-Base Titrations Requirements of a primary standard substance: 1- It must be easily obtained in a very high grade of purity and of known composition. 2- Very stable. 3- Non-hygroscopic and non-volatile. 4- It can be dried at 105-110°C without decomposition. 5- It should have a relatively high molecular weight to minimize weighing error. 6- It must react with other substances in a quantitative way according to balanced chemical equations. Examples of primary standard substances: Acids: Potassium acid phthalate - Oxalic acid - Benzoic acid. Bases: Sodium carbonate - Borax (sodium borate). 13 Acid-Base Titrations (2) Secondary Standard It is a solution of an approximately known concentration. Its concentration can not be directly calculated from the weight of solute and volume of solution. [Remember: Solute + Solvent = Solution] Its exact concentration is determined by a process called “Standardization”. Standardization is a process used to determine the exact concentration of secondary standard solutions and correct any error if present. That is done by: Titrating the 2ry standard solution against a 1ry standard solution OR against a previously standardized 2ry standard. Examples: HCl, H2SO4, NaOH, KOH ………… and others. 14 Acid-Base Titrations Methods of expressing concentration of standard solutions The most common methods are : Molarity (M) & Normality (N). (1) Molarity or Molar concentration (M): It is the number of moles of the solute per 1 liter of the solution. Example: 0.5 M NaOH means 0.5 mole of NaOH / 1 L soln. (2) Normality or Normal concentration (N): It is the number of equivalent weights of the solute per 1 liter of the solution. Example: 2 N HCl means 2 equivalent weights of HCl / 1 L soln. 15 Acid-Base Titrations Remember: Definition of Mole (mol.) For molecules: The mole is gram-molecular weight {i.e. molecular wt. expressed in grams}. For atoms: The mole is gram-atomic weight {i.e. atomic wt. expressed in grams}. Examples: If the mol. wt. of NaOH = 40 then 1 mole of NaOH means 40 g NaOH. (i.e. The weight of 1 mole of NaOH is 40 g) If the atomic weight of Na = 23 then1 mole of Na means 23 g Na. (i.e. The weight of 1 mole of Na is 23 g) 16 Acid-Base Titrations Remember: Definition of Equivalent Weight General definition of the equivalent weight of a substance: It is the weight of the substance that is equivalent in its reactive power to 1 mole of hydrogen. The exact definition and calculation of equivalent weights differ according to the type of the reaction. We will focus in this lecture on acid-base reactions. Definition of the equivalent weight in acid-base reactions: It is the weight of the substance that will release, react with or be chemically equivalent to 1 mole of hydrogen ions (H+) in that reaction. 17 Acid-Base Titrations Remember: Calculation of the equivalent weight in acid-base reactions: 18 Acid-Base Titrations Remember: Advantage of using normal concentrations: The main advantage of using normal concentration over molar ones is the simple 1:1 ratio involved in all reactions of the same type and that facilitates our calculations. Example: 1 equiv. wt. of NaOH ≡ 1 equiv. wt. of HCl They are all equivalent in their reactive powers because they ≡ 1 equiv. wt. of H2SO4 are all referred to a common ≡ 1 equiv. wt. of H3PO4 standard [Hydrogen]. ≡ 1 mole of H+ In other words, solutions of the same normality react in 1:1 ratio. Example: 1 L of 1 N NaOH ≡ 1 L of 1 N HCl ≡ 1 L of 1 N H2SO4 ≡ 1 L of 1 N H3PO4 5 ml of 0.2 N NaOH ≡ 5 ml of 0.2 N HCl ≡ 5 ml of 0.2 N H2SO4 ≡ 5 ml of 0.2 N H3PO4 Note: “ = ” means equal (‫ ) ساوي‬while “ ≡ ” means equivalent (‫) ا أو يتفاعل مع‬. 19 Acid-Base Titrations REMEMBER Molar Concentration Normal Concentration (Molarity) (M) (Normality) (N) It is the number of moles of solute It is the number of equivalent weights of per 1 liter of the solution solute per 1 liter of the solution 𝑵𝒖𝒎𝒃𝒆𝒓 𝒐𝒇 𝒎𝒐𝒍𝒆𝒔 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆 𝑵𝒖𝒎𝒃𝒆𝒓 𝒐𝒇 𝒆𝒒𝒖𝒊𝒗. 𝒘𝒆𝒊𝒈𝒉𝒕𝒔 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆 𝑽𝒐𝒍𝒖𝒎𝒆 𝒊𝒏 𝑳 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 𝑽𝒐𝒍𝒖𝒎𝒆 𝒊𝒏 𝑳 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 How to calculate “Number of Moles”? How to calculate “Number of Equiv. Wts”? 20 Acid-Base Titrations REMEMBER Relation between Mole & Equivalent Weight and between Molarity & Normality of the same solution Value of 1 Piaster = Value of 1 Pound / 100 The number differs according to the reaction type. No. of Piasters = No. of Pounds x 100 The number differs according to the reaction type. Example: For H2SO4 in acid-base reactions 1 Mole H2SO4 Equiv. Weight of H2SO4 = Mol. Weight of H2SO4 / 2  While Normality of H2SO4 solution = Molarity x 2 2 Equiv. (Normality of 1 M H2SO4 = Molarity x 2 = 1 x 2 = 2 N) weights 21 Acid-Base Titrations Relation between concentrations of two reacting solutions OR How we calculate the concentration of the sample from the concentration of the standard? First Method (the method used in pharm. quality control) 1- Calculation of Equivalence Factor (F) {1 ml of standard  ? g of sample} 2- Calculation of Concentration in w/v {g/L, mg/L, g%, mg%} {for more details refer to the practical course} 22 Acid-Base Titrations Relation between concentrations of two reacting solutions OR How we calculate the concentration of the sample from the concentration of the standard? Second Method e.g. If we use normal concentrations: H2SO4 + 2 NaOH Na2SO4 + 2 H2O N1 V1 = N2 V2 N1 V1 = N2 V2 If we use molar concentrations: H2SO4 NaOH M1 V1 M2 V2 M1 V1 M2 V2 𝟏 𝟐 = 𝟏 𝟐 n1 and n2 are the number of moles in a balanced chemical equations. 23 Acid-Base Titrations An important equation used in case of “Dilution of One Solution” C V(before dilution) = C’ V’ (after dilution) C and C’ can be of any concentration unit (M, N, g%, mg% …… etc) but both of them should be of the same unit. V and V’ can be of any volume unit (mL, dL, L …… etc) but both of them should be of the same unit. e.g. What will be the concentration of 96 g% alcohol after diluting 73 mL of it with water to a final volume 100 mL. Answer: 96 x 73 = ? x 100 so the concentration after dilution = 70.08 g% 24 Acid-Base Titrations Important Terms in Titration The Equivalence Point: [ Theoretical ] - It is the point at which the added amount of the standard solution is chemically equivalent to the amount of the substance being determined. - Its real position can only be theoretically calculated. - It is expressed by the volume of the titrant added. The End Point: [ Experimental ] - It is the point at which the visual change of the indicator takes place in the titration. - It can be determined experimentally. - It is expressed by the volume of the titrant added. The Titration Error: - It is the difference between the end point & the equivalence point. - The more accurate the titration, the smaller the titration error. 25 Acid-Base Titrations Detection of the End Point The end point of the titration can be detected by one of the following methods: (1) Color change of an indicator [a reagent that changes its color at the end point]. (2) Some titrants are colored and can be used as self-indicators (ex: KMnO4 in redox titrations). (3) Formation or disappearance of turbidity. (4) Change in the refractive index of the solution. (5) Change in electric potential or conductivity of the solution. 26 Acid-Base Titrations Steps for the determination of a sample using titration 1st Step: Preparation of the standard solution (usually 2ry standards like HCl, H2SO4, NaOH, KOH) and standardization of it by titration against a 1ry standard or a previously standardized 2ry standard & then calculation of the correction factor (f) as follows: Correction factor (f) f 1.05] [ ] [0.95  27 Acid-Base Titrations If f = (0.95-1.05):  No modification in the preparation (i.e. the prepared standard solution 0.1 N NaOH is ready for use), but the concentration of the (f = 0.985) prepared solution should be multiplied by “ f " before using it in any calculations. For example, we use (0.1 x f) N instead of 0.1 N. In calculations, If f (< 0.95) or (> 1.05):  certain modification we use (0.1 x 0.985) should be done in the preparation by diluting or instead of (0.1) concentrating the solution & then (f) is recalculated to ensure that it becomes 0.95-1.05. 2nd Step: Titration of the sample with the prepared standard solution and observation of the E.P. 3rd Step: Calculation of the sample concentration. (refer to the practical course) 28 Acid-Base Titrations Types of Titration Methods 1. Direct titration. 2. Back (Residual) titration. 3. Displacement (Substitution) titration. 1. Direct titration It means stepwise addition of the standard solution (titrant) from the burette into the solution being analyzed until the end point is reached. i.e. Titration of sample ≠ standard (titrant)  Vol. taken from the sample [e.g. 10 ml] ≡ (E.P.) ml of the standard 29 Acid-Base Titrations 2. Back (Residual) titration In this type, a known excess of a standard (more than sufficient to react with the sample) is added to the sample. Then after the reaction is complete, the remaining unreacted excess (residual) of the standard is back titrated against another suitable standard. The amount of the first standard that reacted with the sample can be calculated by subtracting the E.P. which is equivalent to the unreacted amount of the first standard from the initially added volume (the known added excess).  Sample ≡ reacted part ≡ (known excess added of 1st stand. – E.P.) 30 Acid-Base Titrations Diagrammatic Representation of Back (Residual) Titration Sample + 25 mL Standard I  then titration of the remaining Known excess unreacted part ≠ Standard II (e.g. 25, 50, 100 mL)  E.P. 10 mL Sample 25 mL Standard I reacted remaining part unreacted part Titration ≠ Standard II   E.P. Since the remaining unreacted part of Standard I ≡ E.P. only  Sample ≡ reacted part of Standard I ≡ (25 – E.P.) 31 Acid-Base Titrations Back (residual) titration is used in the following cases: (1) If the sample is volatile. [e.g. detn. of ammonia (NH3), formic acid (HCOOH)] (2) If the sample is insoluble in water. [e.g. detn. of zinc oxide (ZnO), calcium carbonate (CaCO3)] (3) If the reaction between the sample & the standard needs heating. [e.g. detn. of ammonium chloride (NH4Cl)] (4) If the reaction between the sample & the standard is slow. (5) If the reaction between the sample & the standard needs the presence of an excess of the standard to achieve stoichiometry. [e.g. detn. of lactic acid] (6) If no suitable indicator is available for direct titration. 32 Acid-Base Titrations 3. Displacement (Substitution) titration In this type, a very weak acid or base is displaced from its salt by a strong acid or base. Typical examples include the titration of sodium carbonate or sodium borate (borax) against standard HCl. Na2CO3 + 2 HCl 2 NaCl + H2O + CO2 HCl (strong acid) displaces carbonic acid (very weak acid) in its salt Note: 33 Acid-Base Titrations Examples of acids and bases in our life 34 Acid-Base Titrations Acid-base Equilibrium How to apply Law of mass action (law of chemical equilibrium) in acid-base reactions? Law of mass action: The rate of a chemical reaction is proportional to the product of multiplication of the active masses (molar concentrations) of the reacting substances. Forward Direction A+B C+D According to the law of mass action: Backward Direction Rate of forward reaction (Rf) = Kf [A][B] Rate of backward reaction (Rb) = Kb [C][D] Kf [C][D] At equilibrium: Rf = Rb so Kf [A][B] = Kb [C][D] Keq = = Kb [A][B] [C]c[D]d In the more general case: aA + bB cC + dD Keq = [A]a[B]b Keq is a constant that is affected only by temperature & pressure. 35 Acid-Base Titrations In acid-base reactions: In case of weak acids: (e.g. CH3COOH) CH3COOH CH3COO- + H+ Keq = Ka = Ionization constant of the acid or Acid dissociation constant [CH3COO−][H+] where (for acetic acid): Ka = & pKa = - logKa [CH3COOH] In case of weak bases: (e.g. NH4OH) NH4OH NH4+ + OH- Keq = Kb = Ionization constant of the base or base dissociation constant [NH4+][OH −] where (for amm. hydroxide): Kb = & pKb = - logKb [NH4OH] N.B. Ka & Kb are calculated only for weak acids and weak bases. 36 Acid-Base Titrations In case of water (H2O): Pure water is a very weak electrolyte (very limited ionization). H2O H+ + OH- [H+][OH−] Dissociation constant of water (Kw) = = [H+][OH-] [H2O] For pure water at room temperature (25°C): [H+][OH-] was experimentally found to be 1 x 10-14 M  Kw = [H+][OH-] = 1x10-14 M -logKw = (-log[H+]) + (-log[OH-]) = - log(1x10-14) pKw = pH + pOH = 14 At neutral medium, [H+] = [OH-] = 1x10-7 M  -log  pH = pOH = 7 At acidic medium, [H+] > 1x10-7 M  -log  pH < 7 At basic medium, [H+] < 1x10-7 M  -log  pH > 7 37 Acid-Base Titrations Relation between: pKa of an acid and pKb of its conjugate base pKa + pKb = pKw = 14 e.g. CH3COOH CH3COO- + H+ pKa (acetic acid)+ pKb (acetate) = 14 pKb of a base and pKa of its conjugate acid pKb + pKa = pKw = 14 e.g. NH4OH NH4+ + OH- pKb (amm hydroxide)+ pKa (ammonium) = 14 -------------------------------------------------------------------------------------------- Note that small “p” means “-log” “[ ]” means “molar concentration” 38 Acid-Base Titrations pH of acids and bases pH of a Strong Acid pH = - log[H+] pH of a Strong Base pOH = - log[OH-] and pH = 14 - pOH pH of a Weak Acid pH = ½ (pKa + pCa) {Ca is the molar concentration of the acid} pH of a Weak Base pOH = ½ (pKb + pCb) and pH = 14 - pOH {Cb is the molar concentration of the base} 39 Acid-Base Titrations pH of Salts Salt of a Strong Acid and a Strong Base pH = 7 e.g. NaCl, K2SO4 Salt of a Strong Acid and a Weak Base pH < 7 e.g. NH4Cl Salt of a Weak Acid and a Strong Base pH > 7 e.g. CH3COONa Salt of a Weak Acid and a Weak Base It depends on the pKa of the acid and pKb of the base 40 Acid-Base Titrations Buffer Solutions Definition: The buffer is a solution which resists changes in pH upon addition of a small amount of a strong acid or a strong base. Buffers are very important to chemical and biological systems. The pH in the human body varies from one fluid to another, for example: the pH of blood is about 7.4, whereas the gastric juice has a pH about 1.5. These pH values, which are critical for enzyme function, are maintained by biological buffers. Buffer consists of either: A weak acid and its salt with a strong base. {Acidic Buffers} [e.g. acetic acid and sodium acetate (acetate buffer)] or A weak base and its salt with a strong acid. {Basic Buffers} [e.g. amm. hydroxide and amm. chloride (ammonia buffer)] 41 Acid-Base Titrations 42 Acid-Base Titrations Buffer action: (How does the buffer resist change in pH?) For acidic buffers [e.g. acetate buffer (CH3COOH & CH3COONa)]  At addition of a strong acid: (e.g. HCl) CH3COO- + H+ CH3COOH (weak acid)  At addition of a strong base: (e.g. NaOH) CH3COOH + OH- CH3COO- + H2O For basic buffers [e.g. ammonia buffer (NH4OH & NH4Cl]  At addition of a strong acid: (e.g. HCl) NH4OH + H+ NH4+ + H2O  At addition of a strong base: (e.g. NaOH) NH4+ + OH- NH4OH (weak base) 43 Acid-Base Titrations pH of Buffers: (Henderson Equation) For acidic buffers [e.g. acetate buffer (CH3COOH & CH3COONa)] [salt] pH = pKa + log [acid] For basic buffers [e.g. ammonia buffer (NH4OH & NH4Cl] [salt] pOH = pKb + log and pH = 14 – pOH [base] 44 Acid-Base Titrations pH range of Buffers: It is the optimum pH range within which the buffer can be effectively used to neutralize added acids or bases, while maintaining a relatively constant pH. at pH = pKa pH range of buffer = pKa ± 1 the buffer has maximum capacity Examples: pKa of acetic acid = 4.74 so acetate buffer is effective in the pH range from 3.74 to 5.74. pKa of ammonium hydroxide = 9.24 (= 14 - pKb) so ammonia buffer is effective in the pH range from 8.24 to 10.24. 45 Acid-Base Titrations Buffer Capacity: It is a measure of the ability of the buffer to resists changes in pH upon addition of a small amount of a strong acid or a strong base. [salt] [salt] The buffer has maximum capacity if the ratio or = 1. [acid] [base] Another quantitative definition of Buffer Capacity is as follows: It is the number of moles of strong acid or strong base required to change pH of 1 liter of the buffer solution by 1 pH unit. 46 Acid-Base Titrations Acid-Base Indicators Types of acid-base indicators: (A) Color Indicators: The end point is detected by change in the indicator color. (B) Turbidity Indicators: The end point is detected by appearance of turbidity. (C) Fluorescence Indicators: The end point is detected by emission of light. (A) Color Indicators They are highly colored organic compounds with weakly acidic or basic characters that undergoes color change at a certain interval of pH called “pH range of the indicator”. They are classified into: Simple color indicators – Screened indicators – Mixed indicators – Universal indicators 47 Acid-Base Titrations (A) Color Indicators 1- Simple color indicators: They are organic compounds that change their colors at the end point. They may be one-colored, such as phenolphthalein (ph.ph) or two-colored, such as methyl orange (M.O) and methyl red (M.R). Phenolphthalein (ph.ph): pH range ≈ (8.5 – 10) approximately Methyl Orange (M.O): pH range ≈ (3 – 4.5) approximately 48 Acid-Base Titrations 2- Screened indicators: = (mixture of acid-base indicator + an inert dye). 3- Mixed indicators: = (mixture of two acid-base indicators) having the same pH range but contrast colors. 4- Universal indicators: = (mixture of multiple acid-base indicators) having different pH ranges. They give different colors over wide pH range, so they can be used for approximate determination of pH in routine work. The main advantage of using screened, mixed or universal indicators is to make the color change sharper than simple indicators. 49 Acid-Base Titrations Titration Curves Titration curve is the curve obtained by plotting pH of the titrated solution (y-axis) against the volume of the titrant added (x-axis). The titration curve is characterized by its sigmoid shape (S-shaped). The midpoint of the vertical part of the curve corresponds to the equivalence point. The shape of the titration curve depends on: (1) Nature of the titrated solution & the titrant. {Acid/Base - Strong/Weak} (2) Concentration of the titrated solution & the titrant. { concentration  height of the vertical part } The titration curve helps in studying the titration reaction and selecting a suitable indicator. 50 Acid-Base Titrations How to predict the shape of the titration curve? 51 Acid-Base Titrations Choice of Acid-Base Indicators  For accurate titrimetric method, the indicator should be selected so as to make the titration error small as possible, which means that the indicator should be selected so that it changes its color as close as possible to the equivalence point.  To achieve that the pH-range of the indicator (i.e. pH-interval within which the indicator changes its color) should fall in the vertical part of the titration curve which is the part marked by sharp change of pH. 52 Acid-Base Titrations Example: Titration of NaOH with standard HCl using M.O. or ph.ph. indicator: NaOH + HCl → NaCl + H2O (pH = 7) The equivalence point (theoretical) corresponds to pH=7, but since the vertical part of this titration curve ranges from about 2 to 12 so M.O. (pH range ≈ 3.0 - 4.5), ph.ph. (pH range ≈ 8.5 - 10) or any acid-base indicator can be used as a suitable indicator for this titration achieving very small titration error. 53 Acid-Base Titrations Titration curves of strong or weak acids & strong or weak bases and suitable acid-base indicators The following slides illustrate the titration curves of different titration cases including: Titration of a strong acid ≠ a strong base OR a strong base ≠ a strong acid. Titration of a weak acid ≠ a strong base. Titration of a weak base ≠ a strong acid. Titration of a weak acid ≠ a weak base OR a weak base ≠ a weak acid. The following three indicators are used as examples of different pH ranges: {1} M.O. indicator [pH range in the acidic region (≈ 3.0 - 4.5)]. {2} Bromothymol blue indicator [pH range in the neutral region (≈ 6.0 – 7.6)]. {3} ph.ph. indicator [pH range in the basic range (≈ 8.5 - 10)]. 54 Acid-Base Titrations Titration of strong acid and strong base HCl + NaOH → NaCl + H2O e.g. At the equivalence point, pH = 7 (Neutral) The pH changes very sharply between about 2 and 12 Titration curve  Almost any indicator can be used like M.O., Bromothymol blue, ph.ph., etc. Indicator  With very dilute solutions, the vertical part decreases & so the choice of the indicator is more limited. Only Bromothymol blue can be employed. 55 Acid-Base Titrations Titration of weak acid against strong base CH3COOH + NaOH CH3COONa + H2O At the equivalence point, e.g. sample titrant (Basic) pH > 7 The pH increases Titration curve rapidly in the alkaline side from about 7.5 to 10.5 Indicator The suitable indicator should have a pH-range in the alkaline side ph.ph. is suitable, Bromothymol blue is unsatisfactory, M.O. is totally unsuitable. 56 Acid-Base Titrations Titration of weak base against strong acid NH4OH + HCl NH4Cl + H2O At the equivalence point, e.g. sample titrant (Acidic) pH < 7 The pH decreases Titration curve rapidly in the acidic side from about 6.5 to 3.5 Indicator The suitable indicator should have a pH-range in the acidic side M.O. is suitable, Bromothymol blue is unsatisfactory, ph.ph. is totally unsuitable. 57 Acid-Base Titrations Titration of weak acid and weak base The pH at equivalence point depends e.g. CH3COOH + NH4OH CH3COONH4 + H2O on pKa of the acid & pKb of the base The titration curve Titration curve does not show any sharp change of pH. Indicator No suitable indicator and so this titration should be avoided 58 Acid-Base Titrations Summary 59 Acid-Base Titrations Titration curves for polyprotic (polybasic) acids A polyprotic acid can be titrated stepwise with good end-point breaks (inflections) for each proton if: the difference in pKa values is at least 4 [i.e. (pKa2 ̶ pKa1) >= 4 ] 4 Ka1 or the ratio of Ka values is at least 10 [i.e. >= 104 ] Ka2 of each two successive protons. H3A H+ + H2A- ………………….. Ka1 , pKa1 more than less than - 2- H2A H+ + HA ………………….. Ka2 , pKa2 more than less than 2- 3- HA H+ + A ……………………. Ka3 , pKa3 60 Acid-Base Titrations Example: Phosphoric acid (H3PO4): H3PO4 is a weak triprotic acid. Its ionization can be expressed as follows: - H3PO4 H+ + H2PO4 …………………..…. pKa1 = 1.96 (≈ 2) H2PO4- H+ + HPO42- ……………………. pKa2 = 7.12 (≈ 7) HPO42- H+ + PO43- …………………..….. pKa3 = 12.32 (≈ 12)  The difference between each 2 successive pKa values is more than 4 so stepwise titration is possible, but we can titrate only the first two protons of H3PO4 separately. The 1st H+ is titrated using M.O. & the 2nd H+ using ph.ph.  The third proton is too weak [pKa3 is very high (>9) or Ka3 is very small (9) or Ka3 is very small (10) Titrant ( pH ≈ 8.3 ) Equations 2nd step: NaHCO3 + HCl → NaCl + CO2 + H2O.....Complete neutralization Titrant ( pH ≈ 3.8 ). Na2CO3 + 2HCl → 2NaCl + CO2 + H2O Volume Volume consumed consumed Titration curve of HCl of HCl (E.Pph.ph.) (E.PM.O.) makes makes half complete neutraliza- neutraliza- tion tion Conclusion E.Pph.ph. ≡ ½ CO32- & E.PM.O. ≡ all CO32- i.e. E.PM.O. = 2 E.Pph.ph. 69 Acid-Base Titrations 2 Titration of Na2CO3 against standard HCl ph.ph. M.O. Representative Diagram 70 Acid-Base Titrations 3 Equations Titration of NaHCO3 against standard HCl NaHCO3 + HCl → NaCl + CO2 + H2O At the equivalence point, Sample (pH ≈ 8.3) Titrant (pH ≈ 3.8 ) pH ≈ 3.8 Titration curve Conclusion E.PM.O. ≡ HCO3- ( ph.ph. can not be used for determination of HCO3- sample ) 71 Acid-Base Titrations 3 Titration of NaHCO3 against standard HCl ph.ph. M.O. Representative Diagram 72 Acid-Base Titrations Summary OH- CO32- HCO3- E.Pph.ph. (1st flask): all OH- ½ CO32- No E.P. E.PM.O. (2nd flask): all OH- all CO32- all HCO3-  The determination of their mixtures depends on using “Double-indicators titration or Two-indicators titration” in which a mixture of two acids or two bases is determined by titration using 2 indicators. i.e. We make 2 titrations in 2 different flasks, one using ph.ph. & the other using M.O., and from the observed end points the concentrations of the mixture components can be calculated.  The reaction between HCl & Na2CO3 or between HCl & NaHCO3 is an example of displacement (substitution) titration where HCl [strong acid] displaces carbonic acid [very weak acid] in its salts. 73 Acid-Base Titrations Determination of NaOH & Na2CO3 mixture The principle depends on double-indicators (two-indicators) method. 1st Flask: Direct titration of the mixture ≠ stand. HCl using ph.ph. ind. E.Pph.ph. ≡ OH- + ½ CO32- 2nd Flask: Direct titration of the mixture ≠ stand. HCl using M.O. ind. E.PM.O. ≡ OH- + CO32- 2(E.PM.O. - E.Pph.ph.) ≡ CO32- E.PM.O. - 2(E.PM.O. - E.Pph.ph.) ≡ OH- Equations: For NaOH: NaOH + HCl → NaCl + H2O For Na2CO3: Na2CO3 + HCl → NaCl + NaHCO3 …... Half neutralization NaHCO3 + HCl → NaCl + CO2 + H2O..Complete neutralization Na CO + 2HCl → 2NaCl + CO + H O 74 Acid-Base Titrations Determination of NaOH & Na2CO3 mixture (Continued) Titration Curves: For NaOH: For Na2CO3: 75 Acid-Base Titrations Determination of Na2CO3 & NaHCO3 mixture The principle depends on double-indicators (two-indicators) method. 1st Flask: Direct titration of the mixture ≠ stand. HCl using ph.ph. ind. E.Pph.ph. ≡ ½ CO32- 2nd Flask: Direct titration of the mixture ≠ stand. HCl using M.O. ind. E.PM.O. ≡ CO32- + HCO3- 2E.Pph.ph ≡ CO32- E.PM.O. - 2E.Pph.ph. ≡ HCO3- Equations: For Na2CO3: Na2CO3 + HCl → NaCl + NaHCO3 …... Half neutralization NaHCO3 + HCl → NaCl + CO2 + H2O..Complete neutralization Na2CO3 + 2HCl → 2NaCl + CO2 + H2O For NaHCO3: NaHCO3 + HCl → NaCl + CO2 + H2O 76 Acid-Base Titrations Determination of Na2CO3 & NaHCO3 mixture (Continued) Titration Curves: For Na2CO3: For NaHCO3: 77 Acid-Base Titrations Determination of HCl & CH3COOH mixture The principle depends on double-indicators (two-indicators) method. 1st Flask: (E.PM.O. for HCl only) Direct titration of the mixture ≠ stand. NaOH using M.O. ind. E.PM.O. ≡ HCl only 2nd Flask: (E.Pph.ph. for total HCl & CH3COOH) Direct titration of the mixture ≠ stand. NaOH using ph.ph. ind. E.Pph.ph. ≡ HCl + CH3COOH (E.Pph.ph. - E.PM.O.) ≡ CH3COOH Equations: For HCl: HCl + NaOH → NaCl + H2O For CH3COOH: CH3COOH + NaOH → CH3COONa + H2O 78 Acid-Base Titrations Determination of HCl & CH3COOH mixture (Continued) Titration Curve: 79 Acid-Base Titrations Determination of HCl & CH3COOH mixture (Continued) Notes: During the determination of HCl only (E.PM.O.), NaOH neutralizes HCl only without interference of acetic acid because the proton (H+) of HCl suppresses the ionization of acetic acid by common-ion effect as shown in the following equations: HCl H+ + Cl- CH3COOH H+ + CH3COO- and when the neutralization of HCl is completed, the pH of the solution reaches the pH-range of M.O. and so M.O. changes its color. While during the determination of total HCl & CH3COOH (E.Pph.ph.), NaOH neutralizes HCl first without any change in the color of ph.ph. and then the titration is completed to make neutralization of acetic acid also. At this point, the pH reaches the pH-range of ph.ph. and so ph.ph. changes its color. 80 Acid-Base Titrations Determination of HCl & CH3COOH mixture (Continued) Notes: The color change of M.O. at the first E.P. is not very sharp as in case of titrating HCl alone against NaOH. This is due to the formation of acetate buffer (CH3COOH & CH3COONa) after finishing the reaction of the titrant (NaOH) with HCl and starting the reaction with CH3COOH. The formed buffer resists the change in pH and so resists color change of M.O. The principle of this determination can be applied for the determination of many (Strong acid/Weak acid) mixtures where: E.PM.O. ≡ Strong acid E.Pph.ph. ≡ Total (Strong acid + Weak acid) Examples of strong acids: HCl, H2SO4 Examples of weak acids: several organic acids like acetic acid, butyric acid, phthalic acid, ….. etc. 81 Acid-Base Titrations Determination of Ammonium Chloride (NH4Cl) NH4Cl is acidic because it is a salt of strong acid (HCl) & weak base (NH4OH). NH4Cl can be determined by 2 methods: [A] Formol method [specific method for ammonium salts]. [B] Back (residual) titration method. (A) Formol Method: 4 NH4Cl + 6 HCHO (CH2)6N4 + 4 HCl + 6 H2O Sample Neutral Hexamine or Equivalent to Formalin Hexamethylene tetramine NH4Cl sample It depends upon addition of HCHO (neutral formalin)  Titr. ≠ stand. NaOH to NH4Cl sample resulting in the formation of using ph.ph. ind. an equivalent amount of HCl that can be titrated against standard NaOH using ph.ph. indicator. 82 Acid-Base Titrations Determination of Ammonium Chloride (NH4Cl) (Continued) (B) Back (residual) titration method: NH4Cl + NaOH NaCl + NH3 + H2O Sample known excess of stand. NaOH NaOH + HCl NaCl + H2O remaining titrant unreacted part It depends upon boiling NH4Cl with a known excess of standard NaOH producing ammonia which is removed by boiling. Then the remaining unreacted part of NaOH is back titrated against standard HCl using M.O. or ph.ph. indicator. 83 Acid-Base Titrations Determination of Ammonium Chloride (NH4Cl) (Continued) (B) Back (residual) titration method: Notes: Boiling is necessary for the reaction between NH4Cl and NaOH to remove the produced ammonia. This method requires Blank Determination, which is a separate determination in which we repeat the same steps of the experiment but without the sample and the end point in this case is called (Blank Reading). Why blank determination is required in this experiment? Because when NaOH is heated and then cooled, its strength is decreased due to: 1- Interaction of NaOH with the glass of the flask. 2- Absorption of atm. CO2. 84 Acid-Base Titrations Diagrammatic Representation of the Blank Determination 85 Acid-Base Titrations Determination of Aspirin (Acetylsalicylic acid) Aspirin is a well-known non-steroidal anti-inflammatory drug (NSAID) used as analgesic, antipyretic and anti-inflammatory. It is a weakly acidic compound. It can be determined by back (residual) titration method. NaOH + HCl NaCl + H2O remaining titrant unreacted part The method depends on boiling aspirin with a known excess of standard NaOH where aspirin is readily hydrolysed into acetic acid and salicylic acid which react with NaOH giving sod. acetate and sod. salicylate. Then the remaining unreacted part of NaOH is back titrated against standard HCl using ph.ph. indicator. 86 Acid-Base Titrations Determination of Aspirin (Acetylsalicylic acid) (Continued) Notes: Boiling is necessary for the reaction between aspirin and NaOH to increase the rate of the hydrolysis reaction. This method requires blank determination because when NaOH is heated and then cooled, its strength is decreased due to: 1- Interaction of NaOH with the glass of the flask. 2- Absorption of atm. CO2. 87 Acid-Base Titrations Determination of Boric Acid (H3BO3) & Determination of Borax (Na2B4O7) Both boric acid & borax have a great pharmaceutical importance. Boric acid is a widely-used antiseptic used to treat different bacterial and fungal infections. Borax is used similarly to boric acid and has also been used externally as a mild astringent. 88 Acid-Base Titrations Determination of Boric Acid (H3BO3) Boric acid is a very weak monoprotic acid (pKa = 9.2) so it can not be directly titrated against a standard alkali. H3BO3 BO2- + H+ + H2O metaborate ion But if a sufficient amount of glycerol, mannitol or similar organic polyhydroxy compound is added to boric acid, it will be converted into a relatively stronger monoprotic acid that can be directly titrated against standard NaOH using ph.ph. indicator. (very weak acid) (relatively stronger acid) Glyceroboric acid + NaOH Sodium glyceroborate + H2O 89 Acid-Base Titrations Determination of Borax (Na2B4O7) (Borax = Sodium Borate = Sodium Tetraborate = Disodium Tetraborate) Borax is hydrolyzed in water as follows:. 90 Acid-Base Titrations Determination of Borax (Na2B4O7) (Continued) The liberated NaOH can be titrated against 0.1 M HCl using M.O. indicator without interference of boric acid as it is a very weak acid. E.P1 ≡ 2 NaOH (from borax) Then borax can be again determined through the liberated boric acid by adding glycerol to the above titrated solution (i.e. the solution obtained after E.P1) and then titration against 0.1 M NaOH using ph.ph. indicator. E.P2 ≡ 4 H3BO3 (from borax) Borax ≡ 2 NaOH ≡ 2 HCl (1st titrant) E.P1 Borax ≡ 4 H3BO3 ≡ 4 Glyceroboric acid ≡ 4 NaOH (2nd titrant) E.P2 If the 1st titrant (HCl) & the 2nd titrant (NaOH) have the same concentration, then E.P2 = 2 E.P1 91 Acid-Base Titrations Determination of Borax (Na2B4O7) (Continued) Notes: The concentration of borax can be calculated from either E.P1 or E.P2 Borax is considered a sodium salt of boric acid and the reaction between HCl & borax is an example of displacement (substitution) titration where HCl [strong acid] displaces boric acid [very weak acid] in its salts. 92 Acid-Base Titrations Determination of Borax (Na2B4O7) (Continued) Question: If the original borax sample is directly treated with excess glycerol, will the resulting solution be acidic or basic? Answer: 93 Acid-Base Titrations Kjeldahl Method {‫}كيلدال‬ Kjeldahl method is the most common method for the determination of nitrogen and protein in organic materials such as food. It includes three main steps, the third step depends on an acid-base titration. The general procedure includes 3 steps: 1. Digestion of the sample with conc. H2SO4 and a catalyst to convert nitrogen in the sample into ammonium salt. 2. Neutralization of excess H2SO4 by NaOH & Distillation of NH3 produced from boiling NaOH with ammonium salt. NH3 is distilled in boric acid solution to form ammonium borate (basic salt). 3. Acid-Base Titration of ammonium borate against standard HCl. The E.P. is equivalent to the nitrogen content in the original sample. 94 Acid-Base Titrations Kjeldahl Method {‫}كيلدال‬ (Continued) 95 Acid-Base Titrations Biphasic Titration Water Biphasic = 2 Phases Organic solvent It is a type of acid-base titrations that is used for determination of salts that are soluble in water but their acids are insoluble in water and soluble in organic solvents like ether, chloroform …etc. Examples of salts determined by biphasic titration: 1. Sodium salicylate. 2. Sodium benzoate. 3. Ammonium salicylate. Main Principle: Salt (in aqueous phase) + stand. HCl Parent acid (in organic phase) Titrant 96 Acid-Base Titrations Determination of Sodium Salicylate By Biphasic Titration Titrant (basic) The aqueous solution of sod. salicylate is mixed with an HCl (titrant) organic solvent (ether) in a separating funnel, then the aqueous layer is titrated with standard HCl using bromophenol blue as indicator (color at E.P. is pale green). When HCl is added to the aqueous layer, salicylic acid is liberated and is extracted into the ether layer. If salicylic acid is not removed into the organic layer, it will interfere with the color of the indicator. 97 Dear Students, It has been my pleasure to teach you these few hours. I hope you feel the same and you all will take what you have learned this year, whether scientifically or ethically and build on it as you continue your studies and after you start your future career. All my best wishes for a bright future for all of you, full of success, joy and happiness. Yours Sincerely, Dr. Mohamed El-Awady

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