Acid-Base Titration: Standardization and Estimation of Sodium Bicarbonate in Antacid

Preparation of Standard Oxalic Acid Solution and Standardization of NaOH

• Objective and Overview:

  • Preparation of a $100\,\text{mL}$ standard solution of $0.1\,\text{M}$ oxalic acid ($\text{H}_2\text{C}_2\text{O}_4$) as a primary standard.

  • Standardization of sodium hydroxide ($\text{NaOH}$) solution using the prepared oxalic acid solution.

• Quantitative Weighing and Solution Preparation Procedure:

  • Turn on the analytical balance and wait until the digital display reads zero.

  • Place a piece of creased weighing paper on the balance pan and press the tare button to subtract the mass of the paper.

  • Slowly transfer oxalic acid salt onto the paper using a spatula until the target mass is reached.

  • Precaution: Do not add excess salt and attempt to return it to the original container, as this risks contaminating the bulk reagent and causing laboratory spills.

  • Quantitatively transfer the weighed oxalic acid salt into a $100\,\text{mL}$ volumetric flask ( standard measuring flask, SMF) using a funnel.

  • Wash down any residual salt adhering to the funnel and the inner neck of the volumetric flask using distilled water to guarantee complete transfer.

  • Dissolve the oxalic acid salt completely by swirling/shaking the flask.

  • Add distilled water up to the etched calibration mark on the neck of the flask.

  • Stopper the flask securely with a lid and invert it repeatedly to ensure complete concentration homogeneity throughout the solution.

  • Label the flask clearly with its determined concentration (e.g., "…. M Oxalic Acid").

• Molarity and Mass Formulas:

  • Molarity ($M$) equation:   $M = \frac{W \times 1000}{M' \times V_{\text{mL}}}$

  • Mass ($W$) calculation formula:   $W = \frac{M \times M' \times V_{\text{mL}}}{1000}$

  • Variable Definitions:

  • $M$ = molarity of the solution ($\text{mol\,dm}^{-3}$ or $\text{M}$)

  • $M'$ = molar mass of the solute ($\text{g\,mol}^{-1}$)

  • $W$ = mass/weight of the solute ($\text{g}$)

  • $V_{\text{mL}}$ = volume of the solution in milliliters ($\text{mL}$)

• Standardization Setup and Procedure for NaOH:

  • Setup Configuration:

  • Burette: Filled with sodium hydroxide ($\text{NaOH}$) solution of unknown concentration ($?\,\text{M}$).

  • Pipette: Used to measure exactly $10\,\text{mL}$ of the prepared $0.1\,\text{M}$ oxalic acid solution into a conical flask.

  • Indicator: Suitable acid-base indicator added to the conical flask.

  • Titration Method:

  • Titrate $10\,\text{mL}$ of oxalic acid solution against $\text{NaOH}$ from the burette until the endpoint color change is attained.

• Chemical Reaction:   $2\text{NaOH} + \text{H}_2\text{C}_2\text{O}_4 \rightarrow \text{Na}_2\text{C}_2\text{O}_4 + 2\text{H}_2\text{O}$

• Stoichiometry and Mole Relations ($n$\text{-factor}):

  • Molarity definition: $\text{Molarity} = \frac{\text{Number of Moles}}{\text{Volume in Liters}}$

  • Moles of $\text{NaOH}$:   $n_{\text{NaOH}} = M_{\text{NaOH}} \times V_{\text{NaOH}}$

  • Moles of Oxalic Acid ($\text{ox}$):   $n_{\text{ox}} = M_{\text{ox}} \times V_{\text{ox}}$

  • Stoichiometric mole ratio from balanced chemical equation:   $\frac{n_{\text{NaOH}}}{n_{\text{ox}}} = \frac{2}{1} = \frac{M_{\text{NaOH}} \times V_{\text{NaOH}}}{M_{\text{ox}} \times V_{\text{ox}}}$

  • Rearranging to find $\text{NaOH}$ molarity:   $M_{\text{NaOH}} = \frac{2 \times M_{\text{ox}} \times V_{\text{ox}}}{V_{\text{NaOH}}}$

Standardization of Hydrochloric Acid (HCl)

• Objective and Overview:

  • Accurate determination of the concentration of a dilute hydrochloric acid ($\text{HCl}$) solution using the standardized $\text{NaOH}$ solution.

• Experimental Setup and Materials:

  • Burette: $50\,\text{mL}$ burette filled with aqueous $\text{NaOH}$ up to the $0\,\text{mL}$ mark using a filter funnel.

  • Pipette: $10\,\text{mL}$ graduated pipette, pre-rinsed with distilled water and subsequently rinsed with the provided $\text{HCl}$ solution.

  • Reaction Vessel: Clean conical flask containing $10\,\text{mL}$ of dilute $\text{HCl}$ measured via pipette bulb.

  • Indicator: $2$ to $3$ drops of phenolphthalein indicator added directly to the $\text{HCl}$ solution.

• Titration Procedure:

  • Swirl the conical flask to thoroughly mix $\text{HCl}$ and phenolphthalein indicator.

  • Dispense $\text{NaOH}$ dropwise from the burette while continuously swirling the flask.

  • Terminate the titration immediately upon the first appearance of a persistent pale pink color throughout the solution.

  • Record the final burette volume reading.

  • Repeat the titration procedure to obtain concordant burette readings.

  • Dispose of the post-titration solution safely in the designated base waste container.

• Chemical Reaction and Equivalence Formula:

  • Neutralization Reaction:   $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

  • Equivalence relation ($1:1$ stoichiometric ratio):   $M_{\text{HCl}} \times V_{\text{HCl}} = M_{\text{NaOH}} \times V_{\text{NaOH}}$

  • Calculation of $\text{HCl}$ concentration:   $M_{\text{HCl}} = \frac{M_{\text{NaOH}} \times V_{\text{NaOH}}}{V_{\text{HCl}}}$

  • Note: The calculated value of $M_{\text{HCl}}$ is utilized as a fixed parameter in subsequent antacid analysis calculations.

Antacid Sample Analysis: Estimation of NaHCO3 via Back-Titration

• Theoretical Principle of Back-Titration:

  • Rationale: Sodium bicarbonate ($\text{NaHCO}_3$) is a weak base that reacts slowly with titrants, making direct titration to a sharp endpoint difficult.

  • Strategy: A known, precise excess of standard strong acid ($\text{HCl}$) is added to react completely with all $\text{NaHCO}_3$ in the sample.

  • Measurement: The leftover, unreacted excess $\text{HCl}$ is subsequently back-titrated using standard $\text{NaOH}$ solution.

• Chemical Reaction:   $\text{NaHCO}_3 + \text{HCl} \rightarrow \text{NaCl} + \text{H}_2\text{O} + \text{CO}_2\uparrow$

• Step-by-Step Laboratory Procedure:

  • Refill the $50\,\text{mL}$ burette with standard $\text{NaOH}$ solution up to the $0\,\text{mL}$ mark using a funnel.

  • Pipette exactly $25\,\text{mL}$ of prepared $0.2\,\text{M}\,\text{HCl}$ solution into a $250\,\text{mL}$ conical flask.

  • Add $10\,\text{mL}$ of the unknown antacid solution containing $\text{NaHCO}_3$ into the flask containing the acid.

  • Stir the solution with a glass rod until the $\text{NaHCO}_3$ powder is completely dissolved.

  • Allow the mixture to rest for $10\,\text{minutes}$ to ensure full reaction between $\text{NaHCO}_3$ and $\text{HCl}$.

  • Add $2$ drops of phenolphthalein indicator to the mixture.

  • Commence back-titration by delivering $\text{NaOH}$ solution from the burette (adding $1\,\text{mL}$ at a time initially, then dropwise near the endpoint).

  • Swirl the conical flask vigorously throughout the titration.

  • Stop the titration immediately when the solution turns from colorless to a permanent pale pink color.

  • Record the final burette volume of $\text{NaOH}$.

  • Repeat the complete determination procedure to obtain consistent, reproducible values.

• Mathematical Calculation Sequence:

1. Volume of Unreacted HCl ($x$\text{ in mL}):

   • Calculated from the Step III back-titration endpoint volume of $\text{NaOH}$:      $x = \frac{V_{\text{NaOH}} \times M_{\text{NaOH}}}{M_{\text{HCl}}}$

2. Volume of Reacted HCl ($y$\text{ in mL}):

   • Total volume of $\text{HCl}$ added minus leftover unreacted volume $x$:      $y = 25 - x$

3. Equivalents / Moles of Antacid ($z$\text{ in mol}):

   • Moles of antacid equal the moles of $\text{HCl}$ consumed during the primary reaction:      $z = \left(\frac{y}{1000}\right) \times M_{\text{HCl}}$

4. Mass of NaHCO3 in Sample ($\text{g}$\text{):

   • Multiply the mole quantity $z$ by the molar mass of $\text{NaHCO}_3$ ($84\,\text{g\,mol}^{-1}$):      $\text{Mass of } \text{NaHCO}_3 = z \times 84\,\text{g}$