Acid-Base Titration PDF

# Pharmaceutical Analysis: Acid Base Titration ## Acid Base Titration An acid-base titration is a quantitative analysis of acid and bases by which known concentration of an acid or base (i.e. titrant) neutralizes the unknown concentration of an acid or base (i.e. titrand). It is based on neutrali...

# Pharmaceutical Analysis: Acid Base Titration ## Acid Base Titration An acid-base titration is a quantitative analysis of acid and bases by which known concentration of an acid or base (i.e. titrant) neutralizes the unknown concentration of an acid or base (i.e. titrand). It is based on neutralization reaction between acid and bases that form water by combination of $H_2O^+$ ions, therefore known as water forming reaction. - The estimation of an alkaline solution by using a standard acidic solution is known as Acidimetry. - The estimation of an acidic solution by using a standard alkaline solution is known as Alkalimetry. The titration progress can be monitored by visual indicators, pH electrodes or both. ## Type of Titration & Neutralization Curve The mechanism of neutralization can be understood by studying the change in the hydrogen ion concentration during the course of appropriate titration. 1. The equivalence point of an acid-base reaction is the point at which the amounts of acid and of base are just sufficient to cause complete neutralization. 2. The pH of the solution at equivalence point is dependent on the strength of the acid and strength of the base used in the titration. - For strong acid-strong base titration, pH = 7 at equivalence point - For weak acid-strong base titration, pH > 7 at equivalence point - For strong acid-weak base titration, pH < 7 at equivalence point ## The Neutralization Curve The neutralization curve can be categorized into four classes, which are listed below: 1. Titration of a strong acid with strong base 2. Titration of a weak acid with a strong base 3. Titration of a strong acid with a weak base 4. Titration of a weak base with a weak acid ### Titration of a Strong Acid with Strong Base Suppose analyte is hydrochloric acid HCl (strong acid) and the titrant is sodium hydroxide NaOH (strong base). If we start plotting the pH of the Titration curve of a strong acid with a strong base, it would look like: - **Point 1:** At point 1, no NaOH is added yet, so the pH of the analyte is low (because it predominantly contains $H_3O^+$ from dissociation of HCl). - $HCl + H_2O \rightleftharpoons H_3O^+ + Cl^-$ - As NaOH is added dropwise, $H_3O^+$ slowly starts getting consumed by $OH^-$ produced by dissociation of NaOH. Analyte is still acidic due to predominance of $H_3O^+$ ions - **Point 2:** This is the pH recorded at a time point just before complete neutralization takes place. - **Equivalent point (Point 3):** This is the equivalence point where, moles of NaOH is equal to moles of HCl in the analyte. - At this point, $H_3O^+$ ions are completely neutralized by $OH^-$ ions. The solution only has salt (NaCl) and water and therefore the pH is neutral i.e. pH 7. - $HCI + NaOH \rightleftharpoons NaCl + H_2O$ - **Point 4:** Addition of NaOH continues, pH starts becoming basic because HCl has been completely neutralized and now an excess of $OH^-$ ions are present in the solution (from dissociation of NaOH). - After this NaOH react with indicator and produce color that is end point. - $NaOH \rightleftharpoons Na^+ + OH^-$ ### Titration of a Weak Acid with a Strong Base Let's assume our analyte is acetic acid $CH_3COOH$ (weak acid) and the titrant is sodium hydroxide NaOH (strong base). If we start plotting the pH of the analyte against the volume of NaOH that we are adding from the burette, we will get a titration curve as shown below. - **Point 1:** No NaOH added yet, so the pH of the analyte is low (it predominantly contains $H_3O^+$ from dissociation of $CH_3COOH$). - $CH_3COOH + H_2O \rightleftharpoons CH_3COO^- + H_3O^+$ - As NaOH is added dropwise, $H_3O^+$ slowly starts getting consumed by $OH^-$ produced by dissociation of NaOH. But analyte is still acidic due to predominance of $H_3O^+$ ions. - **Point 2:** This is the pH recorded at a time point just before complete neutralization takes place. - **Equivalent point (Point 3):** This is the equivalence point. At this point, moles of NaOH added is equal to moles of $CH_3COOH$ in the analyte. The $H_3O^+$ are completely neutralized by $OH^-$ ions. - The solution contains only $CH_3COONa$ salt and $H_2O$. - $CH_3COOH + NaOH \rightleftharpoons CH_3COONa + H_2O^+$ - In the case of a weak acid versus a strong base, the pH is not neutral at the equivalence point. The solution is basic (pH~9) at the equivalence point. - $CH_3COONa = CH_3COO^- + Na^+$ - $CH_3COOH + OH^-$ **Makes solution basic at equivalent point** - **Point 4:** Addition of NaOH continues, pH starts becoming basic because $CH_3COOH$ has been completely neutralized and now excess of $OH^-$ ions are present in the solution (from dissociation of NaOH). After this NaOH react with indicator and produce color that is end point. ### Titration of a Strong Acid with a Weak Base Suppose our analyte is hydrochloric acid HCl (strong acid) and the titrant is ammonia $NH_3$ (weak base). If we start plotting the pH of the analyte against the volume of $NH_3$ that we are adding from the burette, we will get a titration curve as shown below. - **Point 1:** No $NH_3$ added yet, so the pH of the analyte is low (it predominantly contains $H_3O^+$ from dissociation of HCl). - $HCl + H_2O \rightleftharpoons H_3O^+ + Cl^-$ - As $NH_3$ is added dropwise, $H_3O^+$ slowly starts getting consumed by $NH_3$. Analyte is still acidic due to predominance of $H_3O^+$ ions. - $NH_3 + H_2O \rightleftharpoons NH_4^+ + H_2O$ - **Point 2:** This is the pH recorded at a time point just before complete neutralization takes place. - **Equivalent point (Point 3): ** This is the equivalence point. At this point, moles of $NH_3$ added is equal to the moles of HCl in the analyte. The $H_3O^+$ ions are completely neutralized by $NH_3$. - In the case of a weak base versus a strong acid, the pH is not neutral at the equivalence point. - The solution is in fact acidic (pH~5.5) at the equivalence point. At the equivalence point, the solution only has ammonium ions $NH_4^+$ and chloride ions $Cl^-$. Where ammonium ion $NH_4^+$ is the conjugate acid of the weak base $NH_3$. So $NH_4^+$ is a relatively strong acid (weak base $NH_3$ has a strong conjugate acid), and thus $NH_4^+$ will react with $H_2O$ to produce hydronium ions making the solution acidic. - $NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+$ **Makes solution acidic at equivalent point** - **Point 4:** After the equivalence point, $NH_3$ addition continues and is in excess, so the pH increases. $NH_3$ is a weak base so the pH is above 7. - After this NaOH react with indicator and produce color that is end point. ### Titration of a Weak Base with a Weak Acid Suppose our analyte is $NH_3$ (weak base) and the titrant is acetic acid $CH_3COOH$ (weak acid). If we start plotting the pH of the analyte against the volume of acetic acid that we are adding from the burette, we will get a titration curve as shown below. There isn't any steep bit in this plot. There is just because of a 'point of inflexion' at the equivalence point. Lack of any steep change in pH throughout the titration renders. Titration of a weak base versus a weak acid is difficult, and not much information can be extracted from such a curve. The chief feature of the curve is that change of pH near equivalent point and during whole titration is very gradual hence the end point cannot be detected by ordinary indicator so mixed indicator is often used. Even it can be done by using non aqueous acid base titration. ## Choice of Indicator in Titration | **S.NO** | **TYPE OF TITRATION** | **PH** | **INDICATOR** | | -------- | -------- | -------- | -------- | | 1 | Strong acid verses strong base | 2-10 | Phenolphthaleine, Methyl orange, Methyl Red | | 2 | Strong acid verses weak base | 2-6 | Methyl orange, Methyl Red | | 3 | Weak acid verses strong base | 8-10 | Phenolphthaleine, Thymol Phthaleine | | 4 | Weak acid verses Weak base | - | Mixed indicators are used like neutral red methylene blue for dilute ammonia or ethanoic acid. | ## Indicator Properties | **S.No** | **INDIACATOR NAME** | **PH** | **COLOR OBSERVATION** | | -------- | -------- | -------- | -------- | | 1 | Thymol blue | 1.2-2.8 | Acid: Red, Base: Yellow | | 2 | Quinaldine Red | 1.4-3.2 | Acid: Colourless, Base: Red | | 3 | Methyl orange | 2.9-4.6 | Acid: Red, Base: Orange | | 4 | Methyl Red | 4.2-6.3 | Acid: Red, Base: Yellow | | 5 | Bromo Thymol blue | 6-7.6 | Acid: Yellow, Base: Blue | | 6 | Phenol Red | 6.8-8.6 | Acid: Yellow, Base: Red | | 7 | Phenolphthaleine | 8.3-10 | Acid: Colourless, Base: Pink | | 8 | Thymol Phthaleine | 9.5-10.5 | Acid: Colourless, Base: Blue | ## Mixed Indicator Mixture of two or more indicators are used to give a sharp end point over a narrow and selected range of pH. | **MIXTURE OF INDICATOR** | **COMPOSITION** | **PH** | **COLOR CHANGE** | | -------- | -------- | -------- | -------- | | Bromocresol Green : Methyl Green | 1 Part of 0.1% aqueous solution of sodium salt: 1 Part of 0.2% Water. | 4.3 | Orange blue green | | Bromocresol Green : Chlorophenol Red | 1 Part of 0.1% aqueous solution of sodium salt: 1 Part of 0.1% aqueous solution of sodium salt. | 6.1 | Pale green blue violet | | Bromomethyl blue : Neutral Red | 1 Part of 0.1% solution in ethanol : 1 Part of 0.1% solution in ethanol. | 7.2 | Rose Pink green | | Bromomethyl blue : Phenol Red | 1 Part of 0.1% aqueous solution of sodium salt : 1 Part of 0.1% aqueous solution of sodium salt. | 7.5 | Yellow violet | | Thymol blue: Cresol Red | 3 Part of 0.1% aqueous solution of sodium salt: 1 Part of 0.1% aqueous solution of sodium salt. | 8.3 | Yellow violet | ## Universal Indicator A universal indicator is made of a solution of several compounds like water, propane 1-ol, Phenolphthaleine, sodium salt, Thymol blue etc. that exhibit several color changes over a wide range of pH values to indicate the acidity or basicity of solutions. | **PH RANGE** | **DESCRIPTION** | **COLOUR** | | -------- | -------- | -------- | | <3 | Strong acid | Red | | 3-6 | Weak acid | Orange or yellow | | 7 | Neutral | Green | | 8-11 | Weak alkali | Blue | | >11 | Strong alkali | Violet or Indigo | ## Thank you...