Abstract: Examining The Equivalence Point In This Experiment ✓ Solved

Abstract In This Experiment Examining The Equivalence Point In A T

In this experiment, examining the equivalence point in a titration with NaOH identified an unknown diprotic acid. The molar mass of the unknown was found to be 100.78 g/mol with pKa values of 2.6 and 6.6. The closest diprotic acid to this molar mass is malonic acid with a percent error of 3.48%.

The purpose of the experiment was to determine the identity of an unknown diprotic acid. The equivalence and half-equivalence points on the titration curve give important information, which can then be used to calculate the molecular weight of the acid.

The equivalence point is the moment when there is an equal amount of acid and NaOH. Knowing the concentration and volume of added NaOH at that moment, the amount of moles of NaOH can be determined. The amount of moles of NaOH is then equivalent to the amount of acid present. Dividing the original mass of the acid by the moles present gave the molar mass of the acid. In this particular titration, there were two equivalence points as the acid is diprotic.

Consequently, the titration curve had two inflection points. The acid dissociated in a two-step process with the net reaction being: H2X + 2 NaOH → Na2X + 2 H2O. This was important to take into consideration when calculating the molar mass of the diprotic acid. If the first equivalence point was to be used, the ratio of acid to NaOH was 1:1. If the second equivalence point was used in the calculations, the ratio became 1:2 as now a second set of NaOH molecules reacted with the acid to dissociate the second hydrogen ion.

The titration curve also showed the pKa values of the acid. This happened at the half-equivalence point where half of the acid was dissociated to its conjugate base (twice, because of the diprotic property). The Henderson-Hasselbalch equation, pH = pKa + log([A-]/[HA]), indicates that at the half-equivalence point, the pH equals pKa. This was visually represented by the flattest part of the titration curve, where the pKa values could be deduced.

Sample Paper For Above instruction

The determination of the identity of an unknown diprotic acid through titration analysis presents intricate insights into acid-base chemistry, especially concerning equivalence points and pKa values. This study aimed to identify the unknown acid by analyzing a titration performed with sodium hydroxide (NaOH), employing precise volumetric and spectrophotometric measurements to elucidate the molecular characteristics of the acid sample.

The experiment’s fundamental principle revolved around titration, which involves slow addition of a titrant (NaOH) to the analyte (unknown diprotic acid) until the reaction reaches the equivalence point, where stoichiometrically equivalent amounts of acid and base are present. For a diprotic acid, such as the one tested, the titration curve typically exhibits two distinct inflection points corresponding to the two deprotonation steps, each with characteristic pKa values detectable near the half-equivalence points.

To commence, the unknown acid was accurately weighed and dissolved in deionized water to form a solution suitable for titration. The concentration of the NaOH titrant was known, allowing for calculation of the moles of NaOH added at each increment via their volume measurements. The titration was monitored using a pH meter, which provided a continuous curve displaying the pH change as a function of titrant volume. The double inflection points on the titration curve confirmed the diprotic nature of the acid, with the first and second equivalence points signifying the complete deprotonation of each acidic hydrogen.

The critical aspect of the analysis involved calculating the molecular weight of the acid. This process depended on accurately determining the moles of NaOH at the equivalence points. At the first equivalence point, half of the acid had been deprotonated, with the pKa observable from the pH value. Using the Henderson-Hasselbalch equation, the pKa at the half-equivalence point was identical to the pH, allowing an estimation of the acid’s dissociation constant, which in turn informed about its strength and structure.

Analysis of the titration curve yielded a molar mass estimate of approximately 100.78 g/mol for the unknown acid. The calculated pKa values of 2.6 and 6.6 suggest a molecule akin to malonic acid, which bears similar dissociation constants. The percent error in the molar mass determination, 3.48%, reinforced the reliability of the experimental procedures and measurements.

Potential sources of error were acknowledged, notably the partial loss of solid acid during transfer, which may have led to an approximate underestimation of the actual mass. To enhance measurement accuracy, the use of a spectrophotometer to observe the precise wavelength shifts corresponding to the pKa values was suggested for future experiments, offering more definitive detection over traditional colorimetric methods.

The experiment's success hinged on meticulous volumetric titrations, proper calibration of the pH meter, and correct identification of the inflection points on the titration curve. The close match between the calculated molar mass and that of malonic acid validated the approach and highlighted the importance of combined quantitative and qualitative data in chemical identification.

In conclusion, titration-based analysis proved effective in elucidating the identity of the unknown diprotic acid, demonstrating the interconnection between titration data, pKa values, and molecular structure. Such approaches showcase the vital role of analytical chemistry techniques in qualitative and quantitative analysis, fostering accurate chemical characterization of unknown substances.

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