Lecture 3: Acid-Base Titration (Elmansi). PDF

Lecture 3 Volumetric Analysis Lec. 3 Objectives of the Lecture After this lecture, you should know: Acid base Titration pH of buffers...

Lecture 3 Volumetric Analysis Lec. 3 Objectives of the Lecture After this lecture, you should know: Acid base Titration pH of buffers Applications indicators curve Lec. 3  Lec. 3 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.  Lec. 3 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.  Lec. 3 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 and pH = pOH = 7 At acidic medium, [H+] > 1x10-7 M and pH < 7 At basic medium, [H+] < 1x10-7 M and pH > 7  Lec. 3 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”  Lec. 3 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}  Lec. 3 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  Lec.3 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)] Lec. 3  Lec. 3 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. pH range of buffer = pKa ± 1 at pH = pKa the buffer has Examples: maximum capacity ◼ 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.  Lec. 3 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. Lec. 3  Lec.3 Acid-Base 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”. ◼ 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 classified into: Simple color indicators – Screened indicators – Mixed indicators – Universal indicators Lec. 3  Lec.3 (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  Lec.3 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 and the titrant. {strong or weak} (2) Concentration of the titrated solution and the titrant. { concentration  height of the vertical part } The titration curve helps in studying the titration reaction and selecting a suitable indicator.  Lec.3 How to predict the shape of the titration curve?  Lec.3 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.  Lec.3 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.  Lec.3 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)].  Lec.3 Titration of strong acid and strong base HCl + NaOH → NaCl + H2O e.g. At the equivalence point, pH = 7 (Neutral) The pH changes Titration curve very sharply between about 2 and 12 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.  Lec.3 Titration of weak acid against strong base CH3COOH + NaOH ⇌ CH3COONa + H2O At the equivalence point, pH e.g. sample titrant (Basic) >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.  Lec.3 Titration of weak base against strong acid NH4OH + HCl ⇌ NH4Cl + H2O At the equivalence point, pH e.g. sample titrant (Acidic) 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.  Lect.3 2 Titration of Na2CO3 against standard HCl ph.ph. M.O. Representative Diagram Lec. 3