Acid-Base Titration Lecture Notes PDF

Acid-Base Titration PharmD Clinical Pharmaceutical Analytical Chemistry-I Lecture 4,5,6 Dr. rer. nat. Nahla Abdelshafi  Course Outline Fall 2024 Lesson 1. Volumetric Analysis Titration Lesson 2. Acid-Base Theories/pH Calculation...

Acid-Base Titration PharmD Clinical Pharmaceutical Analytical Chemistry-I Lecture 4,5,6 Dr. rer. nat. Nahla Abdelshafi  Course Outline Fall 2024 Lesson 1. Volumetric Analysis Titration Lesson 2. Acid-Base Theories/pH Calculations Lesson 3. Buffers/Indicators Lesson 4. Titration Curves Lesson 5. Applications of Acid-base Titration 2 Volumetric Analysis Titration 3  Types of Quantitative Analysis Fall 2024 a) Volumetric analysis Titration Measuring volume of standard solution(titrant) used for complete reaction with the sample. b) Gravimetric analysis Isolating and weighing of the final product with known pure, stable and definite form. 4  Types of reactions used in volumetric analysis Fall 2024 Volumetric analysis Titration Reaction between standard solution(titrant) and sample. Types of reactions used in volumetric analysis I-Ionic combination reactions Formation of weakly ionizable products A-Neutralization reaction: Formation of water: H+ + OH- → H2O B-Precipitation reaction: Formation of precipitate: Ag+ + Cl- → AgCl↓ C-Complexation reaction: Formation of complex: Ag++ 2 CN-→ [Ag(CN)2]- II-Electron transfer reactions: Redox Titration Electron transfer from one reactant to another 5  Requirements for volumetric reactions Fall 2024 Requirements for volumetric reactions 1-Simple reaction 2-Single reaction 3-Instantaneous reaction (Rapid) 4-Suitable standard must be available as titrant. 5-The end point of the reaction should be easily detected Suitable indicator which changes the solution color at the end point. Change in physical or chemical properties of the solution at the endpoint. TEACH A COURSE 6  Volumetric Titrations Fall 2024 Standard solutions Solutions of exactly known concentration Expressions of standard concentration 1-Molar standard solution 1 liter of it contains 1 mole (Molecular weight)of the substance Eg. HCl (Mwt= 36.5 g) 1M HCl= 36.5 g/L 0.5M HCl= 36.5/2g/L 2M HCl= 2x36.5 g/L Eg. H2SO4(Mwt= 98) 0.5M H2SO4= 98/2 0.1M H2SO4 = 98/10 1M H2SO4= 98 g/L g/L g/L TEACH A COURSE 7  Volumetric Titrations Fall 2024 Standard solutions 2- Normal standard solution 1 liter of it contains 1 equivalent weight of the substance Equivalent weight = Molecular weight /n For Acids: Eq.wt= M.wt/ number of replaceable H+ For Bases: Eq.wt= M.wt/ number of replaceable OH- For Salts: Eq.wt= M.wt/number of cations x its valence Eq.wt= M.wt/ number of anions x its valence Eq.wt of HCl = M.wt /1 Eq.wt of NaOH = M.wt/1 Eq.wt of H2SO4 = M.wt/ 2 Eq.wt of Ba(OH)3 = M.wt/3 Eq.wt of Na2CO3 = M.wt/2 x1 or = M.wt/1 x2 TEACH A COURSE 8  Volumetric Titrations Fall 2024 Standard solutions 2- Normal standard solution 1 liter of it contains 1 equivalent weight of the substance 0.1M H2SO4= M.wt/10 g/L → 98/10 g/L How to prepare 0.1 M H2SO4 and 0.1 N H2SO4? 0.1N H2SO4= Eq.wt/10 g/L , Eq.wt = 98/2 → 98/2x10 g/L Standard solutions having the same normality react by equal volumes Eg. 10 mL of 1N NaOH neutralized with 10 mL of 1N HCl Eg. 10 mL of 1N NaOH neutralized with 10 mL of 1N H2SO4 3-Empirical standard solution 1 mL of it reacts with a definite quantity of the sample No relation between different empirical solutions TEACH A COURSE 9  Standard solutions Fall 2024 1-Direct method For Primary Standard substances Characteristics of primary standard:- 1-Pure 2-Easily tested for impurities. 3-Stable(not absorb H2O or CO2 from air or volatile) 4-React stoichiometrically (quantitatively) 5-Readily soluble. 6-Have high molecular weight to minimize weighing error. Examples of primary standards Potassium acid phthalate KHC8H4O4 Benzoic acid Constant boiling point HCl Anhydrous Na2CO3 Potassium bicarbonate KHCO3 Mercuric Oxide HgO TEACH A COURSE 10  Standard solutions Fall 2024 2-Indirect method When substance is Not primary standard Any substance lose one property of primary standard is considered as secondary standard. Examples of secondary standards Sodium hydroxide (NaOH) Potassium permanganate (KMnO4) First, prepare solution of approximate concentration which is then standardized against primary standard to calculate the standardization factor (f). Standardization factor (f) Number which is multiplied by volume of approximate normality (unknown) to obtain the volume of exact normality (Known). volume of known normality f= volume of unkown normality (0.95–1.05) Equivalent factor (F) Number which represents how many grams of sample react with 1 ml of the primary standard (or corrected secondary standard) TEACH A COURSE 11  Acid Base Theories Fall 2024 1-Arrhenius Theory Acid: Substances which ionize to give H+ E.g. HCl Base: Substances which ionize to give OH- E.g. NaOH 2-Bronsted-Lowry Theory Acid: Substances which donate proton. HCl + H 2O → Cl- + H3O+ Acid Base Conjugated base Conjugated acid Base: Substances which accept proton. NH3 + H2 O → NH4+ + OH- Base Acid conjugated acid conjugated base TEACH A COURSE 12  Acid Base Theories Fall 2024 2-Bronsted-Lowry Theory Amphoteric electrolyte H2O + H2 O → H3O+ + OH- Conjugated acid Conjugated base Acid: Substances which donate proton, Base: Substances which accept proton. Find which is the acid and which is the base in the following conjugate acid-base pairs. Acid Base NH3, NH4+ NH4+ NH3 CH3COOH, CH3COO- CH3COOH CH3COO- HPO42-, H2PO4- H2PO4- HPO42- What is the conjugate acid and conjugate base of HSO4- H2SO4← HSO4- → SO42- Conjugated acid Conjugated base TEACH A COURSE 13  Acid Base Theories Fall 2024 3-Lewis Theory Acid: Substances which accept lone pair of electrons, E.g. BF3, AlCl3 Base: Substances which donate lone pair of electrons, E.g. NH3 Lewis base Lewis acid Lewis base Lewis acid TEACH A COURSE 14  Acid-base titration in aqueous medium Fall 2024 Electrolytic dissociation theory Molecule → Cation+ + Anion- Electrolytes: Which dissociate (ionize) and conduct electricity. Non-electrolytes: Which doesn't ionize and doesn't conduct electricity. Degree of dissociation (α) Number of dissociated molecules α= Total number of molecules before dissociation For strong electrolytes; α is near unity HC l→ H++ Cl- α= 0.92 ≈ One NaOH → Na++ OH- α= 0.91 ≈ One For weak electrolytes; α is far from unity Equilibrium CH3COOH H+ + CH3COO- α= 0.013 H3BO3 H+ + H2BO3- α= 0.001 [Products] K = [Reactants] NH4OH NH4+ + OH- α= 0.013 TEACH A COURSE 15  Acid-base titration in aqueous medium Fall 2024 Electrolytic dissociation theory Acid -base equilibrium in water Equilibrium always exists in solutions of weak electrolytes “ weak acids & weak bases” CH3COOH H+ + CH3COO- NH4OH NH4+ + OH- Ka= [H+][Ac-]/[HAc] Kb = [NH4+][OH-]/[NH4OH] Ka: dissociation constant of acid Kb: dissociation constant of base Dissociation of water H2O H+ + OH- K= [H+][OH-]/[H2O] Kw= [H+][OH-]= 10-14 at 25°C Kw: ionic product of water TEACH A COURSE 16  Acid-base titration in aqueous medium Fall 2024 Dissociation of water K= [H+][OH-]/[H2O] Kw= [H+][OH-]= 10-14 at 25°C For any aqueous solution : Kw= [H+][OH-]= 10-14 at 25°C HCl (Aqueous) Water (H2O) NaOH (Aqueous) Acidic Solution Neutral Solution Basic Solution [H+] > [OH-] [H+] = [OH-] [H+] < [OH-] [H+] > 10-7 [H+] < 10-7 [H+] = 10-3 [H+] = [OH-] = 10-7 [H+] = 10-11 [OH-] = 10-11 [OH-] = 10-3 TEACH A COURSE 17  Acid-base titration in aqueous medium Fall 2024 Dissociation of water Kw= [H+][OH-]= 10-14 at 25°C Hydrogen Ion Exponent (pH) [H+] = 10-3 pH = - log[H+] pH = - log[H+] pH = - log 10-3 pH = -(-3) log 10 pH = 3 [H+] = 10-3 → pH = 3 [H+] = 10-5 → pH = 5 [H+] decreases→ pH value increases [H+] = 10-7 → pH = 7 [H+] = 10-9 → pH = 9 TEACH A COURSE 18  Acid-base titration in aqueous medium Fall 2024 Dissociation of water Kw= [H+][OH-]= 10-14 at 25°C For any aqueous solution : pKw = pH + pOH = 14 Complete the following table by calculating the missing entries, indicating whether the solution is acidic or basic Acidic or pH pOH [H+] [OH-] Basic 5.25 8.75 5.6 x 10-6 1.8 x 10-9 Acidic 12.93 1.07 1.2 x 10-13 8.5 x 10-2 M Basic pH= -log [H+] [H+] = 10- pH [H+] = 10- 5.25 TEACH A COURSE 19  pH of acids and bases Fall 2024 pH of acids and bases pH of strong acids HCl → H+ + Cl- Strong acids are strong electrolytes, so they are completely ionized [H+] = Ca pH = pCa Eg. 0.1 N HCl pH = - log 0.1 = -log 10-1= 1 pH of strong bases NaOH → OH- + Na+ Strong bases are strong electrolytes, so they are completely ionized [OH-] = Cb pOH= pCb pH = pKw – pOH Eg. 0.1 N NaOH pOH= -log 10-1= 1 pH = 14 –1 = 13 TEACH A COURSE 20  pH of acids and bases Fall 2024 pH of acids and bases pH of weak acids pH = ½ pCa + ½ pKa CH3COOH H+ + CH3COO- pH of weak bases pH = pKw - ½ pCb - ½ pKb NH4OH NH4+ + OH- pH of salt solutions Salt of strong acid and strong base, pH= 7 Salt of strong acid and weak base pH= ½ pKw - ½ pKb + ½ pCs Salt of weak acid and strong base pH= ½ pKw + ½ pKa - ½ pCs Salt of weak acid and weak base pH= ½ pKw + ½ pKa - ½pKb TEACH A COURSE 21  pH of acids and bases Fall 2024 pH of acids and bases Find the pH of 0.050 M NH3. (Kb= 1.75 x10–5) pCb= -log 0.05 =1.3 pKb= -log 1.75 x 10-5= 4.76 pH= 14 - (1.3/2) - (4.76/2) = 10.97 Find the pH of 0.020 M CH3COONa. (pKb= 9.24) pCs= -log 0.02 = 1.70 pKa= 14 - 9.24 = 4.76 pH= 7 + (4.76/2) - (1.7/2) = 8.53 Find pH of 0.50 M CH3COONa. pCs = -log 0.5 = 0.30 pH = 7 + (4.76/2) - (0.3/2) = 9.23 TEACH A COURSE 22  Buffer solutions: Fall 2024 Solutions which resist changes in pH upon addition of small amount of acid or base They consist of weak acid and its salt (conjugate base) or weak base and its salt (Conjugate acid) Weak acid & its salt OR Weak base & its salt CH3COOH CH3COO-Na+ NH4OH NH4+Cl- (conjugate base) (conjugate acid) OH H+ OH - H+ - CH3COO- CH3COOH NH4+ NH4OH + H2O + H 2O TEACH A COURSE 23  Buffer solutions: Fall 2024 Solutions which resist changes in pH upon addition of small amount of acid or base They consist of weak acid and its salt (conjugate base) or weak base and its salt (Conjugate acid) TEACH A COURSE 24  pH of acids and bases Fall 2024 Buffer solutions: weak acid-base conjugate pair Henderson-Hasselbalch equation [base] pH = pKa+ log [acid] Calculate the pH of a buffer containing 0.1 M acetic acid and 0.1 M sodium acetate (pKa=4.76) pH= 4.76 + log 0.1/ 0.1= 4.76 Calculate the pH of a buffer containing 0.07 M NH3 and 0.28 M NH4Cl (pKb= 4.74) pKw = pKa + pKb pH= 9.26 + log 0.07/0.28= 8.65 pKa = pKw – pKb pKa= 14 – 4.74 = 9.26 TEACH A COURSE 25  pH of acids and bases Fall 2024 Buffer Capacity: It is the magnitude of the resistance of buffer to change in pH. Maximum buffer capacity is obtained when 1- High concentration of [acid] & [base] If we add base to this buffer (OH-) HAC + OH- → AC- + H2O 2- [acid] = [base]. At this case, pH = pKa 10 1 H OH- 1 10 + 10 10 Effective pH range HAC AC- HAC AC- HAC AC- pH = pKa ± 1 pH = pKa-1 pH = pKa pH = pKa+1 TEACH A COURSE 26  Buffer Solutions Fall 2024 Calculate the pH of a buffer that is 0.015 M in H2PO4- and 0.020 M in HPO42- [base] pH = pKa + log H3PO4 H+ + H2PO4- pKa = 2.1 [acid] H2PO4- H+ + HPO42- pKa = 7.2 HPO42- H+ + PO43- pKa = 12.4 [0.02] pH = 7.2 + log [0.015] pH = 7.33 Calculate pH of 0.1 N HCl pOH 0.1N NH4OH (pKb=4.76) pH of 0.1N acetic acid (pKa=4.76) pH of 0.1 M NaOH pH of 0.1N sodium acetate (pKa=4.76) pH of 0.1N ammonium chloride (pKb=4.76) pH of a solution containing 0.1M acetic acid and 0.2 M sodium acetate (pKa=4.76) TEACH A COURSE 27  Indicators Fall 2024 1- Color Indicators: indicator changes its color by changing the pH, sample should be colorless and clear. 2-Turbidity Indicators For colored solutions Indicators that form flocculent precipitate or turbidity at the end point. 3-Fluorescence Indicators For colored or turbid solutions Compounds emit visible radiation when exposed to UV light. This property may stop or intensify at end point. E.g. Umbelliferone TEACH A COURSE 28  1-Color Indicators Fall 2024 Theories of color indicators Ostwald-Arrhenious Theory (Ionization): “ The indicator is a weak organic acid or base that have the color of the dissociated form differ from un-dissociated form” Eg. Acidic indicator (H-Ind) → Phenolphthalein (Ph.Ph.) HInd H+ + Ind- Eg. Basic indicator (Ind-OH) → Methyl Orange(M.O.) IndOH OH- + Ind+ Objections to ostwald theory 1-Ph.Ph. becomes colorless upon addition of more alkali. 2-The change in color is gradual while ionic reactions are instantaneous. 3-The color of the indicator changes when titration is done in non aqueous medium where no ionization takes place. 29  1-Color Indicators Fall 2024 Chromophoric Theory: “ The color change is not due to ionization only but due to intramolecular rearrangement which change the structure of the indicator and change its color.” -OH (-H2O) Ph.Ph. Ph.Ph. (Acidic medium) (Alkaline medium) TEACH A COURSE 30  1-Color Indicators Fall 2024 Chromophoric Theory: Ph.Ph. (Phenolphthalein) Benzenoid (Colorless) Quinonoid (Pink) Tribasic Salt (Colorless) Below pH 7 pH 8-10 above pH 12 TEACH A COURSE 31  1-Color Indicators Fall 2024 Chromophoric Theory: The indicator molecule contains Chromophoric group: -C=C-C=C- NO2 -N=N- Unsaturated groups responsible for light absorption in the visible region therefore they are colored groups. If they change or rearrange the color of the indicator will changes. Auxochrome group -OH, NH2 Saturated groups “not colored” but aid in color formation due to lone pair of electrons which extend conjugation. TEACH A COURSE 32  1-Color Indicators Fall 2024 Screened Indicator: Indicator + dye → Sharp color change Methyl orange + indigo carmine Color change: Yellowish green (Alkaline) Violet(Acid) Mixed Indicator: Mixture of 2 indicators (Thymol blue + Cresol red) having the same pH range but showing contrasting color to give sharper color change Universal Indicator: Mixture of more than one indicator having different pH range For approximate pH determination TEACH A COURSE 33  1-Color Indicators Fall 2024 pH calculation of indicators [𝑩𝒂𝒔𝒊𝒄 𝒄𝒐𝒍𝒐𝒓] Henderson equation pH = pKin + log [𝑨𝒄𝒊𝒅𝒊𝒄 𝒄𝒐𝒍𝒐𝒓] When [basic color] = [acidic color] pH = pKin Effective Range of indicator pH = pKin ± 1 pH 2 3 4 5 6 7 8 9 10 11 PhenolPhthalein (8 – 10) Methyl Orange (3.1 – 4.4) Methyl Red (4 – 6) TEACH A COURSE 34  Neutralization indicators Fall 2024 Titration curves 1- Strong acid – strong base titration 2- Weak acid – strong base titration 3- Weak base – strong acid titration 4- Weak base – weak acid titration pH 12 11 10 9 PhenolPhthalein 8 7 6 5 4 Methyl Orange 3 2 1 0 Volume of Standard (mL) TEACH A COURSE 35  Titration curves Fall 2024 1- Strong acid – strong base titration 100 mL 0.1N HCl # 0.1N NaOH 0.1N NaOH pH before titration: 0.1N HCl H+ OH- H pH = pCa, pH = – log10-1 = 1 H+ H+ + pH during titration: HCl↓ pH HCl+ NaOH→ Na+Cl-+ H2O 12 11 10 pH rises slowly till 99.9% of HCl is neutralized (pH 4) 9 8 7 pH at end point: NaCl 6 5 4 All HCl is neutralized (pH7) 3 2 1 pH after end point: NaOH (pH 11) 0 Volume of 0.1 N NaOH (mL) pH =pKw – pCb All indicators are suitable TEACH A COURSE 36  Titration curves Fall 2024 2- Weak acid - strong base titration 100 mL 0.1N CH3COOH # 0.1N NaOH (pKa = 4.74) 0.1N NaOH pH before titration: 0.1N CH3COOH AC- OH- AC- HAC HAC AC- pH = ½ pKa + ½ pCa, pH during titration: HAC/AC- pH 12 Buffer solution 11 10 9 pH = pKa + log(base/acid) 8 7 6 pH at end point: AC- 5 4 3 pH = ½ pKw + ½ Pka – ½ pCs 2 1 0 pH after end point: NaOH Volume of 0.1 N NaOH (mL) pH =pKw – pCb Ph. Ph. Is the suitable indicator. TEACH A COURSE 37  Titration curves Fall 2024 3 - Weak base strong acid titration 100 mL 0.1N NH4OH # 0.1N HCl pKb = 4.74 0.1N HCl pH before titration: 0.1N NH4OH NH4+ H+ NH + 4 NH3 NH3 NH4+ pH = pKw –½ pkb –½ pCb pH during titration: NH3/NH4+ pH Buffer solution 12 11 10 pH =(pKw - pKb ) + log (base/acid) 9 8 7 pH at end point: NH4+ 6 5 4 pH = ½ pKw - ½ Pkb + ½ pCs 3 2 1 pH after end point: HCl 0 Volume of 0.1 N HCl (mL) pH = pCa M.O. Is the suitable indicator. TEACH A COURSE 38  Titration curves Fall 2024 4 Weak base weak acid titration 0.1N CH3COOH Smooth titration curve with no abrupt change in pH NH3 NH3 NH3 No sharp End Point NH3 NH3 No End Point can be detected pH 12 11 10 9 8 7 6 5 4 3 2 1 0 Volume of 0.1 N HAc (mL) TEACH A COURSE 39  Applications of acid-base titration Fall 2024 I-Direct titration methods Determination of Acids 1) Strong acid # Strong Std base (Ph.Ph. Or M.O.) 2) Weak acid ka > 10-7 # Strong Std base (Ph.Ph.) Determination of Bases 1) Strong base # Strong Std acid (Ph.Ph. Or M.O.) 2) Weak base kb> 10-7 # Strong Std acid (M.O.) TEACH A COURSE 40  Applications of acid-base titration Fall 2024 Determination of Bases 1) Strong base # Strong Std acid (Ph.Ph. Or M.O.) 2) Weak base kb> 10-7 # Strong Std acid (M.O.) 3) Very Weak acid ka < 10-7 Boric Acid (H3BO3) Potentiated by polyhydroxy compound # Strong Std base (Ph.Ph.) Glycerol Glycerol Using neutralized glycerol TEACH A COURSE 41  Applications of acid-base titration Fall 2024 I-Direct titration methods Determination of Acids 4) Insoluble acid: Dissolve then titrate Benzoic and Salicylic acids Dissolved in neutral alcohol # Strong Std base (Ph.Ph.) Free fatty acids Dissolved in Ether-Ethanol (1:1) # Strong Std base (Ph.Ph.) II-Double indicator titration Determination of : → Mixture of two monobasic acids → Di-proticacids & Poly-proticacids Applied when: 1.Difference in pKa is at least 4 2.No one of them has Ka less than 10-7 TEACH A COURSE 42  Applications of acid-base titration Fall 2024 II-Double indicator titration Determination of : → Mixture of two monobasic acids Mixture of HCl (strong acid) & CH3COOH (weak acid) pH 12 11 HCl → H+ + Cl- 10 9 8 7 HAC HA- + H+ Ka= 5.5x10-5 6 5 4 3 2 HCl suppresses the ionization of HAC 1 0 Volume of 0.1 N NaOH (mL) NaOH neutralize HCl first M.O. Then NaOH neutralize HAC Ph.Ph. TEACH A COURSE 43  Applications of acid-base titration Fall 2024 III-Displacement titration Determination of → Hydrolysable salts Salts of strong base + weak acids (KCN, Na2B4O7, Na2CO3) KCN + H2O → HCN + K+ + OH- KCN + HCl → HCN + KCl # Std acid (M.O.) Salts of strong acids + weak base FeCl3, Al2(SO4)3 AlCl3 + 3H2O → Al(OH)3 ↓ + 3HCl AlCl3 + 3NaOH → Al(OH)3 ↓ + 3NaCl # Std base (Ph.Ph.) TEACH A COURSE 44  Applications of acid-base titration Fall 2024 II-Double indicator titration Determination of : → Di-protic acid & poly-protic acids ❑ Difference in pKa is at least 4 ❑ Ka of each is not less than 10-7 H2A HA- + H+ ka1 HA- A2- + H+ Ka2 TEACH A COURSE 45  Applications of acid-base titration Fall 2024 II-Double indicator titration Determination of : → Di-protic acid & poly-protic acids H2CO3 ❑Difference X in pKa is at least 4 ❑ X Ka of each is not less than 10-7 Titrated as monobasic acid in one step H2CO3 HCO3- + H+ Ka1 = 4.3x10-7 pKa1 = 6.37 HCO3- CO32- + H+ Ka2 = 5.6 x10-11 pKa2 = 10.25 TEACH A COURSE 46  Applications of acid-base titration Fall 2024 II-Double indicator titration Determination of : → Di-protic acid & poly-protic acids H2S ❑Difference √ in pKa is at least 4 ❑ X Ka of each is not less than 10-7 Titrated as Dibasic acid in one step using Cu2+ ions. H2S HS- + H+ Ka1 = 9.1x10-8 pKa1 = 7.04 HS- S- + H + Ka2 = 1.2 x10-15 pKa2 = 14.92 Indirect determination by precipitating copper sulphide and titrate the released H+ H2S + Cu2+ → CuS ↓ + 2H+ TEACH A COURSE 47  Applications of acid-base titration Fall 2024 II-Double indicator titration Determination of : → Di-protic acid & poly-protic acids H3PO4 ❑Difference √ in pKa is at least 4 ❑ √ Ka of each is not less than 10-7 pH1 = (pK1 + pK2)/2 = 4.66 H3PO4 H2PO4- + H+ pKa1 = 2.12 H2PO4- HPO42- + H+ pKa2 = 7.21 pH2 = (pK2 + pK3)/2 = 9.94 HPO42- PO43- + H+ pKa3 = 12.67 pH3 = 7 + ½ pKa – ½ pCs = 12.8 2H3PO4 + 3 CaCl2 → Ca(PO4)2 + 6HCl TEACH A COURSE 48  Applications of acid-base titration Fall 2024 II-Double indicator titration Determination of : → Di-protic acid & poly-protic acids H3PO4 ❑Difference √ in pKa is at least 4 ❑ Ka of each is not less than 10-7 M.O. First Ionization Ph.Ph. Second Ionization CaCl2 Total protons 2H3PO4 + 3 CaCl2 → Ca(PO4)2 + 6HCl Volume of 0.1 N NaOH (mL) TEACH A COURSE 49  Applications of acid-base titration Fall 2024 III-Displacement titration Determination of → Hydrolysable salts Salts of strong base + weak acids (Na2B4O7 (Borax)) Na2B4O7 + 7H2O → 4H3BO3 + 2NaOH Borax 2 4 H+ OH- Glycerol # Std acid (M.O.) 4 H+ 2X # Std base (Ph.Ph.) TEACH A COURSE 50  Applications of acid-base titration Fall 2024 III-Displacement titration Determination of → Hydrolysable salts Salts of strong base + weak acids (Na2B4O7 (Borax)) Na2B4O7 + 7H2O → 4H3BO3 + 2NaOH Boric Borax Boric /Borax Mixture 2 ? H+ 4 H+ OH- Glycerol # Std acid (M.O.) ? H+ 4HH++ 2X +? H+ # Std base (Ph.Ph.) TEACH A COURSE 51  Applications of acid-base titration Fall 2024 III-Displacement titration Determination of → Hydrolysable salts Salts of strong base + weak acids Na2CO3 # Std HCl CO32- HCO3- CO2 + H2O Half CO32- (Ph.Ph.) (half neutralization) * All CO32- (M.O.) Half neutralization Complete neutralization pH 11.7 8.3 3.8 Ph.Ph. E.P. M.O. E.P. E.P. How to prevent escape of CO2? 1. Cooling 2. Dilution Na2CO3 + 2HCl → 2NaCl + CO2 + H2O 3. Stirring Na2CO3 + CO2 + H2O → 2NaHCO3 4. Dipping (burette nozzle in 2Na2CO3 + 2HCl → 2NaHCO3 + 2NaCl solution) TEACH A COURSE 52  Applications of acid-base titration Fall 2024 III-Displacement titration Determination of → Hydrolysable salts # Std HCl CO32-/ HCO3- CO32- HCO3- (Ph.Ph.) (M.O.) Total (M.O.) mL1 Half CO32- (Ph.Ph.) mL2 CO32- = 2x mL2 HCO3- = mL1 – (2x mL2) TEACH A COURSE 53  Applications of acid-base titration Fall 2024 IV- Biphasic titration Two immiscible solvents Determination of salts whose acids are water insoluble Na benzoate, Na salicylate # Std acid (Ph.Ph.) Pink (alkaline)→ Colorless (Acidic) COONa COOH OH OH + HCl → + NaCl Change indicator color Removed by extraction in Ether TEACH A COURSE 54  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration Why Back? 1- Insoluble sample ZnO, CaO, CaCO3, BaCO3 2- Reactions required heat CaCl2, BaCl2, SrCl2 3- Volatile sample NH3, Formic acid Sample # Standard Known xss Standard TEACH A COURSE 55  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration A- Determination of insoluble oxides and carbonates ZnO, CaO, CaCO3, BaCO3 ZnO Back with addition of NH4Cl ZnO + 2HCl → ZnCl2 + H2O Std NH4Cl X NaOH ZnO ZnCl Known 2 # Std NaOH (M.O.) xss Std HCl Zn(OH)2 ↓ TEACH A COURSE 56  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration A- Determination of insoluble oxides and carbonates ZnO, CaO, CaCO3, BaCO3 ZnO Back with addition of NH4Cl CaO, CaCO3, BaCO3 Back without addition of NH4Cl CaO Direct with addition of n-sucrose and n-alcohol CaO CaO + Sucrose → Ca-saccharate n-Sucrose Ca-sacc Soluble # Std HCl (Ph.Ph.) ► n-sucrose To avoid acidity due to bacterial growth. ► n-alcohol Added to prevent lumpiness of oxide. TEACH A COURSE 57  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration B- CaCl2, BaCl2, SrCl2 Neutral Salts # Std HCl (M.O.) ►Reactions required heat CaCl2 + Na2CO3 → CaCO3↓ + 2NaCl CaCl2 CaCO Known3 xss Std Na2CO3 Why Ph.Ph.? Ph.Ph. will get the end point at half neutralization, so less acid will be consumed in the titration and no attack to CaCO3 ppt TEACH A COURSE 58  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration C- Determination of inorganic ammonium salts NH4Cl NH4Cl + NaOH → NaCl + NH3 + H2O Indirect (Back) Boil the sample with Std NaOH NH4Cl # Std HCl (M.R.) NH3 xss Std NaOH Known Std NaOH NH3 NH4Cl NH4Cl NH4Cl 59 TEACH A COURSE 59  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration D- Determination of nitrogen in organic compounds Kjeldalh’s Method For Organic compound Organic Conc. H2SO4 compound CO2 + NH3 + SO2 K2SO4 + CuSO4 + H2SO4 K2SO4 or Na2SO4 : Raise Boiling Point (NH4)2SO4 CuSO4 or HgO: Catalyst Add NaOH to the sample (NH4)2SO4 → NH3 Distill the produced ammonia into known excess Std HCl NH3 HCl HCl Titrate excess un-reacted HCl with Std NaOH HCl NH3 NH 60 3 HCl HCl TEACH A COURSE 60  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration E- Other indirect titrations Determination of Ammonium Salts and Amino Acids Formol Titration Ammonium salts 4NH4Cl + 6HCHO → (CH2)6N4 + 4HCl + 6H2O Hexamine # Std NaOH (Ph.Ph.) Amino Acids HCHO + NH2-CH2-COOH → H2C=N-CH2-COOH + H2O # Std NaOH (Ph.Ph.) TEACH A COURSE 61  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration E- Other indirect titrations Determination of Aldehydes & Ketones By adding Hydroxylamine hydrochloride RCHO + H2N-OH.HCl → RCH=N-OH + HCl + H2O R2CO + H2N-OH.HCl → R2C=N-OH + HCl + H2O # Std NaOH (M.O.) TEACH A COURSE 62  Applications of acid-base titration Fall 2024 V- Indirect titration Back or Residual Titration E- Other indirect titrations Determination of formaldehyde HCHO + H2O2 → HCOOH + H2O Formic Acid Volatile! Back Titration The reaction is done in presence of known excess NaOH HCOOH + NaOH → HCOO-Na+ + H2O Excess base is back titrated against standard acid using Ph.Ph. as indicator. TEACH A COURSE 63 References: 1. https://gtu.ge/Agro- Lib/Vogels_TEXTBOOK_OF_QUANTITATIVE_CHEMICAL_ANALYSIS_5th_ed__- _G_H_Jeffery.MsuCity.pdf 2. https://gtu.ge/Agro- Lib/Harvey%20D.%20Modern%20analytical%20chemistry%20(MGH,%202000) (816s).pdf Thank You! Room 225 Office hours: Tuesday: 11-1 PM