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Titrations stand as one of the cornerstone experiments within the realm of chemistry, playing a pivotal role in quantifying the concentration of an unknown solution through its reaction with a solution of known concentration. This report delves into the comprehensive findings derived from an extensive pH titration experiment. The experiment was meticulously conducted, centering on the utilization of hydrochloric acid (HCl) as the enigmatic solution and sodium hydroxide (NaOH) as the discernible solution of known concentration. Divided into two distinct yet interconnected parts, the experiment unfolded through the meticulous execution of titration methodologies employing both traditional indicators and sophisticated pH meters.
In this part of the experiment, a 0.25 molar NaOH solution was used to titrate the unknown HCl solution.
The following steps were followed:
This initial phase of the experiment commenced with the preparation of a precise 0.25 molar NaOH solution, meticulously filled into a 50-milliliter buret, with the initial volume meticulously recorded to ensure accuracy.
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Following this, a volume ranging between 20 and 40 milliliters of the mysterious HCl solution was precisely measured out, with every minuscule detail of the volume meticulously noted for subsequent calculations.
The measured volume of the HCl solution was then gently introduced into a meticulously cleaned 100-milliliter Erlenmeyer flask, ensuring no spillage or contamination occurred.
Subsequently, to visually monitor the progression of the titration, two drops of phenolphthalein indicator were judiciously added to the solution within the Erlenmeyer flask. This indicator, with its characteristic transition from colorless to a delicate shade of pink, served as a visual cue for the nearing of the endpoint of the titration process. With all components meticulously prepared, the titration process commenced, as the NaOH solution was gradually introduced into the HCl solution within the Erlenmeyer flask. The addition of NaOH was meticulously controlled, with the solution being stirred gently to ensure homogeneity and accuracy in the detection of the color change indicating the endpoint of the titration. Upon achieving the faint pink hue indicative of the endpoint, the final volume of the NaOH solution within the buret was meticulously recorded, marking the completion of this phase of the experiment.
In this part of the experiment, the same NaOH solution was used to titrate a different volume of the unknown HCl solution. The procedure was as follows:
| What Was Measured | Measurement |
|---|---|
| Volume of HCl used | 20.0 mL |
| Initial Volume of NaOH in buret | 50.0 mL |
| Final Volume of NaOH in buret | 34.0 mL |
The color in the flask changed from colorless to a light pink that remained constant even after several seconds of stirring.
| What Was Measured | Measurement |
|---|---|
| Volume of HCl used | 30.0 mL |
| Initial volume of NaOH in buret | 50.0 mL |
| Initial volume of NaOH in buret when end point was reached | 26.0 mL |
| Initial pH of HCl | 0.70 |
| pH after 5 mL NaOH added | 0.87 |
| pH after 10 mL (total) NaOH added | 1.06 |
| pH after 15 mL NaOH added | 1.30 |
| pH after 20 mL NaOH added | 1.70 |
| Volume at which pH = 7.0 | 24.0 mL NaOH added to the flask (26.0 mL left in buret) |
| pH after 25 mL NaOH added | 11.65 |
| pH after 30 mL NaOH added | 12.40 |
| pH after 35 mL NaOH added | 12.63 |
| pH after 40 mL NaOH added | 12.76 |
| pH after 45 mL NaOH added | 12.85 |
| pH after 50 mL NaOH added | 12.91 |
Following the meticulous titration process, the volume of NaOH utilized in the reaction can be precisely determined through simple arithmetic. By subtracting the initial volume of NaOH, as recorded prior to the titration, from the final volume observed after the titration, the exact volume of NaOH dispensed into the solution is elucidated.
Final Volume of NaOH - Initial Volume of NaOH = Volume of NaOH used
34.0 mL - 50.0 mL = 16.0 mL
To determine the molarity of NaOH:
Molarity = (moles of NaOH) / (volume of NaOH used in liters)
Molarity = (1 mol/L * 16 mL) / 1000 mL
Molarity ≈ 0.420 mol/L
Hence, the calculated molarity of the NaOH solution approximates 0.420 mol/L, signifying the concentration of the NaOH solution employed in the titration experiment.
The experiment confirmed that NaOH is a base and HCl is an acid. As the amount of base (NaOH) increases, the reaction becomes less acidic, leading to an increase in pH. This relationship was evident in the trend observed during titration, where the pH gradually increased as more NaOH was added. When the pH reached 7, the solution was considered neutral.
Stirring the solution in the flask was found to be crucial as it enhanced the rate of the reaction and ensured that the reaction was complete before any conclusions were drawn. Failure to remove bubbles from the buret could lead to inaccuracies in volume measurements, affecting the calculated molarity of the solutions. Other factors such as adding too much base or acid, impurities in the equipment, or mathematical errors could also impact the accuracy of the results and conclusions drawn from the experiment.
In conclusion, pH titration experiments are valuable tools in chemistry for determining unknown concentrations of solutions and understanding acid-base reactions. Proper technique and attention to detail are essential for obtaining accurate and reliable results.
pH Titration Experiment. (2024, Feb 26). Retrieved from https://studymoose.com/document/ph-titration-experiment
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