Pharmacy Analysis PDF

Analysis Unit 1: Fundamentals of volumetric analysis and titrametry methods [4 Hrs] After the completion of the course, students will be able to a Discuss Mole concept b Discuss different methods of expressing concentration c Define and list primary and secondary standards d Discuss preparation, st...

Analysis Unit 1: Fundamentals of volumetric analysis and titrametry methods [4 Hrs] After the completion of the course, students will be able to a Discuss Mole concept b Discuss different methods of expressing concentration c Define and list primary and secondary standards d Discuss preparation, standardisation and storage of various volumetric solutions.  Volumetric analysis Volumetric analysis, also known as titration, is a widely used technique in analytical chemistry to determine the concentration of a substance in a solution This method relies on the reaction between a solution of known concentration (the titrant) and the analyte (the substance being analyzed) to determine the analyte's concentration. This method called titration  Titration methods are essential in various scientific and industrial applications, including pharmaceuticals, environmental analysis, food chemistry, and more. Here are the fundamentals of volumetric analysis and common titrametry methods:  Standard Solution: In titration, you have a solution of known concentration called the standard solution or titrant. This solution is used to react with the analyte (the substance of unknown concentration) to determine its concentration accurately. Let's say you want to determine the concentration of hydrochloric acid (HCl) in a sample. To do this, you prepare a standard solution of sodium hydroxide (NaOH) of known concentration. Sodium hydroxide can react with hydrochloric acid in a one- to-one ratio according to the balanced chemical equation: NaOH + HCl → NaCl + H2O  Preparation of Standard Solution: You weigh out a precise amount of sodium hydroxide (e.g., 0.1 moles) and dissolve it in a known volume of distilled water (e.g., 1000 mL). This solution is now your standard solution of sodium hydroxide. Titration: You take a sample of the hydrochloric acid solution of unknown concentration and add a few drops of an indicator, such as phenolphthalein, to it. Using a burette, you slowly add the standard solution of sodium hydroxide to the hydrochloric acid solution while stirring. The sodium hydroxide will react with the hydrochloric acid, and the indicator will change color when all the acid has reacted. The point at which the indicator changes color is called the endpoint of the titration.  Equivalence Point: The equivalence point is a critical concept in titration. It's the point at which the stoichiometrically equivalent amounts of the titrant and analyte have reacted. At this point, you have completely neutralized the analyte or reached another predetermined endpoint, which could be a color change or a change in a physical property.  The endpoint is the The equivalence point at which an point is the point at indicator changes which the Definition color, signaling that stoichiometrically the reaction has equivalent amounts nearly reached of the reactants have completion. reacted completely. No indicator is An indicator is used to necessarily used at detect the endpoint the equivalence Indicator by changing color point; it is determined when the reaction is by the stoichiometry nearing completion. of the reaction.  The endpoint is The equivalence visually observable point is not typically when the indicator observed visually but Visual Observation changes color, is calculated based allowing for manual on data collected detection. during titration.  Moles A mole is defined as the amount of substance that contains exactly Avogadro's number of elementary entities (usually atoms or molecules). This number is approximately 6.02214076 × 10^23  For example, the molar mass of water (H2O) is approximately 18.015 g/mol, so one mole of water weighs approximately 18.015 grams. Moles (n) = Mass (m) / Molar Mass (M) Where: Moles (n) is the amount of substance in moles. Mass (m) is the mass of the substance in grams. Molar Mass (M) is the molar mass of the substance in grams per mole (g/mol).  Calculate the number of moles of water (H2O) in 36 grams of water. Moles of H2O = Mass of H2O / Molar Mass of H2O Moles of H2O = 36 g / 18 g/mol Moles of H2O = 2 moles  Molar mass? the mass of one mole of a substance, expressed in grams per mole (g/mol). Calculate the molar mass of carbon (C): The atomic mass of carbon is approximately 12.011 u,so Molar Mass of Carbon (C) = 12.011 g/mol Example: Calculate the molar mass of water (H2O): The molar mass of hydrogen (H) is approximately 1.008 g/mol. The molar mass of oxygen (O) is approximately 15.999 g/mol. Molar Mass of H2O = (2 × Molar Mass of H) + Molar Mass of O Molar Mass of H2O = (2 × 1.008 g/mol) + 15.999 g/mol  Molar Mass of H2O = 2.016 g/mol + 15.999 g/mol Molar Mass of H2O = 18.015 g/mol  Measuring Amounts of Substances: Moles are used to quantify the amount of a substance in a sample. By knowing the number of moles, chemists can determine how many atoms, molecules, or ions are present in a given quantity of a substance. This is essential for conducting chemical reactions and making precise measurements. Balancing Chemical Equations: In chemical reactions, it is crucial to ensure that the number of atoms of each element on the reactant side is equal to the number of atoms on the product side. Moles play a vital role in balancing chemical equations because they allow chemists to relate the quantities of reactants and products. Calculating Molar Mass: Molar mass is the mass of one mole of a substance and is expressed in grams per mole (g/mol). Calculating the molar mass of a compound is essential for stoichiometry (the study of the quantitative relationships in chemical reactions) and for determining the mass of a given quantity of a substance.  Determining Empirical and Molecular Formulas: Moles are used to find the empirical and molecular formulas of compounds. By analyzing the mole ratios of elements in a compound, chemists can determine its chemical formula. Stoichiometric Calculations: Moles are essential for performing stoichiometric calculations, such as determining the limiting reactant (the reactant that is completely consumed in a reaction) and calculating the theoretical yield (the maximum amount of product that can be obtained). Gas Law Calculations: In gas chemistry, moles are used in the ideal gas law (PV = nRT) to relate the pressure (P), volume (V), temperature (T), and number of moles (n) of a gas. This equation is crucial for understanding the behavior of gases. Concentration Calculations: Moles are used to calculate the concentration of solutions in terms of molarity (moles of solute per liter of solution). This is essential for analytical chemistry, titrations, and preparing solutions with precise concentrations. Determining Percent Composition: Moles are used to find the percent composition of elements in a compound, which is important for understanding the composition of substances.  Concentration In chemistry refers to the measure of the amount of a substance (the solute) that is dissolved or suspended in a given volume of another substance (the solvent).  There are several methods for expressing concentration Molarity (M): Molarity is defined as the number of moles of solute per liter of solution. It is expressed in moles per liter (mol/L or M). The formula for molarity is: Molarity (M) = moles of solute (n) / volume of solution (V, in liters)  Mass Concentration: This is also known as mass/volume percent, weight/volume percent, or mass/volume ratio. It expresses the mass of the solute (in grams) per unit volume of the solution (in milliliters or liters). It is often denoted as a percentage (%). The formula is:Mass Concentration (%) = (mass of solute / volume of solution) × 100%  Molality (m): Molality is defined as the number of moles of solute per kilogram of solvent. It is expressed in moles of solute per kilogram (mol/kg). The formula for molality is: Molality (m) = moles of solute (n) / mass of solvent (in kg)  Mole Fraction (X): Mole fraction is the ratio of the moles of a particular component (solute or solvent) to the total moles in the solution. It is a dimensionless quantity and is often expressed as a decimal. The formula for mole fraction is: Mole Fraction (X) = moles of component / total moles in solution  Volume Percent: This method expresses the volume of solute (in milliliters) per 100 milliliters of the solution. It is used primarily for liquid-liquid solutions. The formula is: Volume Percent = (volume of solute / volume of solution) × 100%  Mole Ratio: In some cases, the concentration may be expressed as a simple ratio of moles of solute to moles of solvent or moles of solute to moles of the total solution. This is a straightforward way to express concentration without specifying volume.  Normality (N) is a measure of the concentration of a solute in a solution and is commonly used in acid-base chemistry and redox reactions. It represents the number of equivalents of a solute per liter of solution.  For Acids and Bases: Normality (N) = (Number of moles of solute) / (Volume of solution in liters) × (Equivalent factor) The "number of moles of solute" refers to the moles of the acid or base in the solution. The "volume of solution" is in liters. The "equivalent factor" is a number that represents the number of acidic or basic equivalents in one mole of the solute. For a monoprotic acid (e.g., HCl), the equivalent factor is 1. For a diprotic acid (e.g., H2SO4), the equivalent factor is 2 because it can donate two moles of H+ ions per mole.  For Redox Reactions: Normality (N) = (Number of moles of oxidizing or reducing agent) / (Volume of solution in liters) × (Equivalent factor) The "number of moles of oxidizing or reducing agent" refers to the moles of the substance involved in the redox reaction. The "volume of solution" is in liters. The "equivalent factor" for redox reactions is determined by the number of electrons transferred per mole of the substance. For example, in the reaction 2Fe²⁺ + Cl₂ → 2Fe³⁺ + 2Cl⁻, the equivalent factor for Fe²⁺ is 2 because it can donate 2 moles of electrons. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) In this reaction, one mole of HCl reacts with one mole of NaOH to produce one mole of NaCl and one mole of water (H2O). The reaction is a 1:1 stoichiometry for HCl and NaOH.  Suppose you have 0.5 liters of a hydrochloric acid solution, and you want to find its normality. The solution contains 0.2 moles of HCl. Calculate the normality of the HCl solution. SOLUTION Normality (N) = (Number of moles of solute) / (Volume of solution in liters) × (Equivalent factor) Number of moles of HCl = 0.2 moles Volume of solution = 0.5 liters HCl is a monoprotic acid, so its equivalent factor is 1. Now, plug these values into the formula: Normality (N) = (0.2 moles) / (0.5 liters) × (1) Normality (N) = 0.4 N So, the normality of the hydrochloric acid (HCl) solution is 0.4 Normal (N).  Indicator: Indicators are substances that change color near the endpoint of the titration. They help signal that you are close to the equivalence point. Common indicators include phenolphthalein (colorless to pink in acidic to basic solutions) and bromothymol blue (yellow to blue). Stoichiometry: Knowledge of the balanced chemical equation for the reaction between the titrant and analyte is crucial. It tells you the molar ratio between the two substances and allows you to calculate the moles of analyte based on the moles of titrant used. Buret: A buret is a laboratory instrument used for precise dispensing of the titrant solution. The volume of titrant delivered from the buret is carefully measured and used in calculations.  Titration Curve: A titration curve is a graph of the pH (or other relevant property) of the solution being titrated against the volume of titrant added. It helps visualize the progress of the titration, including the buffering region before the equivalence point and the sharp pH change at the equivalence point. Primary Standard: A primary standard is a highly pure substance that can be used to prepare a standard solution. It should have a known, accurate molar mass and be stable in the air. Common primary standards include potassium hydrogen phthalate (KHP) and sodium carbonate (Na2CO3).  Dilution: Dilution is often required to reduce the concentration of the analyte or titrant to a more manageable range or to increase the precision of the analysis. The principles of dilution and the dilution equation (C1V1 = C2V2) are crucial in titration experiments. C1​V1​=C2​V2​ Where:  1C1​ = Initial concentration of the solution (before dilution)  1V1​ = Initial volume of the solution (before dilution)  2C2​ = Final concentration of the solution (after dilution)  2V2​ = Final volume of the solution (after dilution) Here's an example of a numerical problem involving dilution: Problem: You have a solution of hydrochloric acid (HCl) with a concentration of 0.4 M (moles per liter). You need to prepare 250 mL of a less concentrated HCl solution with a concentration of 0.1 M. How much of the 0.4 M solution should you use, and how much water should you add? Solution: First, write down the formula for dilution: C1​V1​=C2​V2​.  Primary Standard A primary standard is a highly pure and stable substance that can be used to prepare a standard solution with a precisely known concentration. Primary standards are typically used in titration and other analytical techniques to determine the concentration of an unknown substance.  Examples Sodium carbonate (Na2CO3) for acid-base titrations. Potassium dichromate (K2Cr2O7) for redox titrations. Anhydrous sodium sulfate (Na2SO4) for water content determination.  Secondary Standard A secondary standard, also known as a working standard, is a substance or solution of known concentration that is prepared using a primary standard. Secondary standards are used for routine calibration and analytical measurements, but they are not as pure or stable as primary standards.  Commercially available solutions of sodium hydroxide (NaOH) for titrations where the exact concentration of NaOH is not critical. Potassium permanganate (KMnO4) solutions prepared from primary standard KMnO4 for use in redox titrations.  A volumetric solution, often referred to as a standard solution or titrant, is a prepared solution of a precisely known concentration that is used in analytical chemistry to determine the concentration of another substance through a chemical reaction  How to prepare volumetric solution? Materials Needed: standard solution (a highly pure and stable substance with a known chemical composition). Volumetric flask (a glass container with a precise volume). Deionized or distilled water. Analytical balance. Glassware (pipettes, burettes, beakers). Stirring rod.  To study the steps involved in volumetric analysis MATERIALS REQUIRED Spatula, Funnel, Burette, Conical flask, Burette stand, Dropper, Volumetric pipette,Volumetric flask, Glass rod, Beaker, weighing balance.  THEORY Terms involved in volumetric analysis: Titrant: Solution whose concentration is known. Titrand: Solution whose concentration is unknown. Stoichiometric / End point: It shows that reaction between titrant and titrand is complete. Standard solution: Solution whose exact concentration is known. Titration: Reaction between titrand and titrant.  Steps involved in volumetric analysis: Method selection: For analysis of base (NaOH) → acid is used (HCl): Acid – Base titration. Sampling: Small amount of chemical is taken as sample. Solution preparation: Using appropriate formula, weight of chemical is calculated, weighed and dissolved in suitable solvent. Removing interferences: Calibration, blank titration, parallel determination is done. Observation: Volume of Titrant used for end point is observed. Calculation: Using equivalent factor, concentration of sample solution is calculated. Result analysis: Sample pass or fail as per pharmacopoeial standards Volumetric Preparation Solution Type Purpose Standard Steps Sodium Weigh a known Standard Acid Acid-base amount of carbonate Solution titrations Na2CO3. (Na2CO3) Dissolve it in Potassium deionized or hydrogen distilled water. phthalate Dilute to the Hydrochloric desired volume acid (HCl) in a volumetric flask, mixing thoroughly.  Potassium Weigh a known Standard Acid-base amount of KHP. hydrogen Base Solution titrations phthalate Sodium Dissolve it in hydroxide deionized or distilled water. (NaOH) Dilute to the Sodium desired volume carbonate in a volumetric (Na2CO3) flask, mixing thoroughly.  Weigh a known Standard Sodium Acidimetry (Acid Sodium hydroxide amount of NaOH. Hydroxide (NaOH) content) (NaOH) Dissolve it in deionized or distilled water. Dilute to the desired volume in a volumetric flask, mixing thoroughly.  Measure a known volume Standard of Alkalimetry Hydrochloric Hydrochloric concentrated (Base content) acid (HCl) Acid (HCl) HCl..Dilute it to the desired concentration with deionized or distilled water in a volumetric flask.  Standardization of a volumetric solution It is the process of determining its exact concentration or molarity through a carefully controlled chemical reaction or titration with primary standard.  To prepare and standardize 100 ml of 0.1 N NaOH solution using Oxalic Acid as primary standard.  MATERIALS REQUIRED Apparatus Required: Funnel, volumetric flasks, conical flask, volumetric pipette, beaker, burette, burette stand, weighing balance. Chemicals Required: NaOH, oxalic acid, phenolphthalein solution.  PROCEDURE 1. Preparation of NaOH solution …….. g of NaOH was weighed, dissolved in distilled water & volume made upto 100 ml. 2. Preparation of 0.1 N Oxalic acid solution. ……..g of Oxalic Acid was weighed, dissolved in distilled water & volume made upto 100ml 3. Standardization of NaOH solution 10 ml of Oxalic Acid solution & 1-2 drops of phenolphthalein solution were taken in a conical flask. To that solution NaOH solution was added until the solution became just pink. The titration was performed 3 times Step Procedure Choose a highly pure and stable primary 1. Selection of Primary Standard standard substance. Prepare a secondary standard solution 2. Preparation of Secondary Standard from the primary standard. Titrate the secondary standard solution 3. Titrating the Secondary Standard with the primary standard solution. Calculate the concentration of the 4. Calculating the Concentration secondary standard solution based on titration results. Conduct multiple titrations for precision 5. Repeat and Average and calculate the average concentration. Record the determined concentration 6. Record and Adjust and make any necessary adjustments if required.  Preparation of a 0.1 M Sodium Hydroxide Preparation of Volumetric Solutions: (NaOH) Solution Selection of Primary Standard: Na2CO3 and Weight it (e.g., 4.0 grams of Na2CO3) Dissolving the Primary Standard: deionized or distilled water Transferring to a Volumetric Flask Dilution with Water: Slowly add deionized or distilled water to the volumetric flask, filling it up to the calibration mark (e.g., 1 liter). Mixing: The solution is now a 0.1 M NaOH standard solution. Na2CO3 + H2O → 2NaOH + CO2  Standardization of Volumetric Solutions: Preparation of the HCl Solution: standardized HCl solution of known concentration (e.g., 0.1 M Preparing the Burette: Fill the burette with the standardized HCl solution Preparing the NaOH Solution: Pipette a precise volume (e.g., 25.00 mL) of the prepared NaOH solution into a clean and dry Erlenmeyer flask. Adding Indicator: drops of phenolphthalein indicator to the Erlenmeyer flask containing the NaOH solution. Titration: Place the Erlenmeyer flask under the burette. (Slowly add the standardized HCl solution drop by drop to the NaOH solution while swirling the flask continuously. The pink color will disappear initially but reappear as the reaction progresses) Endpoint Detection Recording the Volume: Record the final volume of HCl solution used from the burette Calculating the Concentration  Storage of volumetric solution ? type of Volumetric Solution Storage Considerations - Store in glass bottles with tight-sealing Acid Solutions (e.g., HCl) caps. - Keep away from strong bases and reactive chemicals. - Protect from light if sensitive to photodegradation. - Store in glass bottles with tight-sealing Base Solutions (e.g., NaOH) caps. - Keep away from strong acids and reactive chemicals. - Protect from light if sensitive to photodegradation.  - Store in glass bottles with tight-sealing Salt Solutions (e.g., NaCl) caps. - Protect from contamination by other salts. - Maintain proper labeling to prevent mix-ups. - Store in amber glass bottles to protect Standard Solutions (e.g., KHP) from light. - Keep the bottles tightly sealed to prevent evaporation or contamination. - Check the concentration periodically to ensure stability. - Store in glass bottles with tight-sealing Indicator Solutions caps. - Protect from contamination and evaporation. - Label clearly for easy identification. Complexometric Titration Solutions - Store in glass bottles with tight- (e.g., EDTA) sealing caps. - Keep away from heavy metal ions that can interfere. - Label for concentration and purpose.  Titration Titration is a laboratory technique used in analytical chemistry to determine the concentration of a specific substance (the analyte) in a solution. It involves the controlled addition of a known solution (the titrant) to the solution containing the analyte until the chemical reaction between the two solutions is complete Titration Method Description Key Components - Titrant (standardized acid Neutralization of acids and or base) - Indicator (e.g., bases to determine Acid-Base Titration phenolphthalein, concentration of one or bromothymol blue) - Buret the other. - Analyte (acid or base) Involves transfer of - Titrant (oxidizing or electrons in a redox reducing agent) - Indicator reaction to determine the Redox Titration (if needed) - Buret - concentration of a Analyte (reducing or reducing or oxidizing oxidizing agent) agent. - Titrant (complexing agent Formation of stable like EDTA) - Indicator (if complexes with metal ions Complexometric Titration needed) - Buret - Analyte to determine metal ion (metal-containing concentration. solution)  Formation of an insoluble - Titrant (precipitating agent) precipitate; used to - Indicator (if needed) - Buret Precipitation Titration determine concentration of - Analyte (solution containing one of the ions involved. the ion of interest) Titrations performed in non- aqueous solvents (typically - Titrant - Indicator (if Non-Aqueous Titration organic solvents) when the needed) - Buret - Analyte (in analyte is insoluble or reacts a non-aqueous solvent) with water. Excess of titrant is added, - First titration components and the remaining titrant is (titrant, indicator, buret, Back Titration titrated with a second analyte) - Second titration reagent; used for difficult-to- components (second titrate analytes. reagent, indicator, buret) Measurement of gases - Titrant (gas-phase reagent) produced during a chemical - Appropriate gas-measuring Gas-Phase Titration reaction; commonly used for equipment - Analyte (gas- oxygen determination. containing sample)  Indicator Indicator is a substance or a component that is used to detect and determine the presence or absence of a specific chemical or the endpoint of a chemical reaction. Indicators are often employed in various analytical techniques to signal a change in the properties of a solution or a reaction, such as a change in color, pH, or electrical conductivity, that can be easily observed or measured.  Theory of indicators Colour change theory Structural theory  Imagine indicators as color-changing molecules in a solution. These molecules can exist in two forms, like having two outfits: one for acidic situations and one for basic situations. 1. Outfits and Colors: In an acidic environment (lots of H⁺ ions around), the indicator molecule wears one outfit and appears as one color (let's say red). In a basic environment (lots of OH⁻ ions around), it switches to a different outfit and changes color (let's say yellow). 2. The Change Process: When you add the indicator to a solution, it starts in one outfit (color) or the other based on the pH. As you change the pH of the solution (make it more acidic or more basic), the indicator can switch outfits and change color. 4. The Color Change Point: There's a special point in between (when the ratio of outfits is roughly equal) where the indicator is in the process of changing colors. This point is called the "end point." At the end point, you see the most noticeable color change because the indicator is almost wearing both outfits equally.  The key idea is that the color of the indicator depends on the ratio of its acidic form (HIn) to its basic form (In⁻), which varies with pH. When the solution is acidic, there are more H⁺ ions, and the acidic form of the indicator (HIn) predominates, giving a certain color. When the solution is basic, there are more OH⁻ ions, and the basic form of the indicator (In⁻) predominates, giving a different color. The point at which the ratio of HIn to In⁻ is approximately 1:1 is called the indicator's "end point," and this is where the most noticeable color change occurs.  Structural theory The structural theory of indicators study into the molecular structure of the indicator compounds. It explains. In this theory, the color change is attributed to alterations in the conjugated double bond systems within the indicator molecule. The change in electron distribution affects the absorption of light by the compound, leading to a change in color.  Common examples of indicators based on these theories include: Phenolphthalein: This indicator is colorless in acidic solutions (HIn form) and pink in basic solutions (In⁻ form). Methyl orange: It is red in acidic solutions (HIn form) and yellow in basic solutions (In⁻ form).  Acid-Base Titrations In acid-base titrations, indicators are substances that change color in response to variations in the pH (acidity or basicity) of the solution. They help determine the endpoint of the titration, which is when the reaction between the acid and base is stoichiometrically complete.  Acid-Base Titrations Phenolphthalein: It changes from colorless (acidic) to pink (basic). This indicator is commonly used in acid-base titrations. Methyl Orange: It changes from red (acidic) to yellow (basic) and is suitable for acid-base titrations, especially when the pH transition range is around 3 to 4. Bromothymol Blue: It shifts from yellow (acidic) to blue (basic) and is used in titrations involving strong acids and bases. Litmus: Litmus paper turns red in acidic solutions and blue in basic solutions, making it a simple indicator for acid-base titrations. Universal Indicator: This indicator provides a range of colors across the pH scale, making it versatile for a wide range of acid-base titrations.  Redox Titrations In redox titrations, indicators are used to detect the endpoint of the titration, which is when the redox reaction is stoichiometrically complete. These indicators change color when there is a specific change in the oxidation state of the analyte or titrant.  S  Redox Titrations Potassium Permanganate: It changes from purple to colorless as it undergoes reduction and is often used in redox titrations. Starch-Iodine Complex: The formation of a blue-black complex when iodine is generated is indicative of the endpoint in redox titrations.  Complexometric Titrations Complexometric titration is a type of titration used to determine the concentration of metal ions in a solution by forming a complex. The most common complexometric titration involves the use of ethylenediaminetetraacetic acid (EDTA) as the titrant. one EDTA molecule per metal ion. The metal ion solution is titrated with a standard solution of EDTA.  Complexometric Titrations Eriochrome Black T: It shifts from wine red to blue at the endpoint and is commonly used in complexometric titrations, especially those involving the determination of metal ions. Calcein: Used for the titration of calcium and other metal ions, calcein forms complexes that change color during titration.  Complexometric Titrations Complexometric titration is a type of titration used to determine the concentration of metal ions in a solution by forming a complex. The most common complexometric titration involves the use of ethylenediaminetetraacetic acid (EDTA) as the titrant. one EDTA molecule per metal ion. metal ion solution is titrated with a standard solution of EDTA.  Moa Add the indicator to the Metal ion solution, which turns red due to the Ca-indicator complex. Titrate with EDTA until the color changes from red to blue, indicating all Ca²⁺ ions are complexed with EDTA.  Easy way Step 1 : indicator in free state ( pink) Step 2: metal sample added + indicator Step 3: indicator is bound with metal ion ( indicator is binded , not in free state ) colourless Step 4: titrant added Step 5: titrant start attracting the meta ion from indicator leading to indicator in free state once all the titrant or edta bind with metal ion ,thus leading to free state of idicator ( pink )  Precipitation titration is a type of titration in which the formation of a precipitate (insoluble solid) indicates the endpoint of the reaction. In precipitation titrations, indicators are often not used in the same way as in acid-base or redox titrations, where indicators change color. Instead, the endpoint is determined by observing the formation of a visible precipitate.  Precipitation titration Adsorption Indicators: These are substances that are adsorbed onto the surface of the precipitate, causing a change in turbidity or appearance. Common adsorption indicators include: Ferric Ammonium Sulfate: Used in the Mohr method for chloride determination, where it forms a reddish-brown precipitate with silver ions. pH Indicators for Complexometric Precipitation: In some complexometric titrations, pH indicators may be used to detect the formation of metal complexes and the subsequent precipitation. Examples include: Eriochrome Black T: It changes from wine red to blue at the endpoint during complexometric titrations, indicating the formation of a metal-indicator complex.  Nonaquous titration Nonaqueous titrations involve titration reactions that take place in solvents other than water, typically organic solvents. In nonaqueous titrations, the choice of indicator is crucial, as the indicators must be soluble in the nonaqueous solvent and exhibit a sharp color change at the endpoint.  Phenolphthalein: when the solvent is an alcohol or an organic solvent. It changes from colorless to pink at the endpoint in the presence of a strong base. Bromothymol Blue: Bromothymol blue is soluble in organic solvents and can be used as an indicator in nonaqueous titrations, especially in the presence of strong acids and bases. It shifts from yellow (acidic) to blue (basic).  Mechanism of Types of Indicator Type Action Color Change Titrations Acidic form Colorless (acidic) Acid-Base Phenolphthalein pH Indicator (HIn) vs. Basic to Pink (basic) Titration form (In⁻) Acidic form Red (acidic) to Acid-Base Methyl Orange pH Indicator (HIn) vs. Basic Yellow (basic) Titration form (In⁻) Acidic form Bromothymol Yellow (acidic) Acid-Base pH Indicator (HIn) vs. Basic Blue to Blue (basic) Titration form (In⁻) Acidic form Red (acidic) to Acid-Base Litmus pH Indicator (HIn) vs. Basic Blue (basic) Titration form (In⁻)  Multiple color Red (acidic) to Universal Acid-Base pH Indicator changes across Purple (neutral) Indicator Titration pH range to Blue (basic) Reduction of Potassium Purple to Redox Indicator MnO₄⁻ ions to Redox Titration Permanganate Colorless Mn²⁺ ions Formation of Colorless to Starch-Iodine Redox Indicator blue-black Redox Titration Blue-Black complex with I₂ Formation of Wine Red to Eriochrome Complexometri Complexometri metal-indicator Blue (at Black T c Indicator c Titration complex endpoint)