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An acid-base titration is a procedure that can be conducted to determine the concentration of an unknown acid or base. In an acid-base titration, a certain amount of a titrant with a known concentration is added to completely neutralize the titrand— the unknown concentration, reaching the equivalence point. The equivalence point is reached when the moles of titrant added to the solution is stoichiometrically equal to the titrand in the solution. The purpose of the experiment was to first be able to determine if an unknown solution was a buffered, or an unbuffered solution. A buffer solution is compromised of a weak acid or a weak base and its salt, and can be shown in chemical equation form such as CH3 COOH(aq) + NaOH(aq)→CH3 COONa + H2O(l). The purpose in the second part of the experiment was to titrate an unknown solution with 0.1 M NaOH and use the titration curve formed to determine if the unknown acid in the solution was monoprotic or diprotic, as well as determine its pKa(s) and molecular weight.
Table 1: Initial and Final pH of Unknown Solutions
|Solution||Initial||HCl treated||NaOH Base treated|
Table 2: Calculated Molar Masses of Unknown Acids
|Iodic Acid||HIO3||175.9 g/mol|
|Oxalic Acid||H2C2O4||90.0 g/mol|
|Maleic Acid||H2C4H2O4||116.1 g/mol|
|Sulfurous Acid||H2SO3||82.1 g/mol|
|Malonic Acid||H2C3H2O4||104.1 g/mol|
|Mandelic Acid||HC8H7O3||152.2 g/mol|
Chart 1: Titration of Unknown Acid B with NaOH
In part one, ~3-mL samples of aqueous unknown 1 were added to two separate 10-mL graduated cylinders, and the initial pH was recorded by using a pH probe. 5 drops of a dilute strong acid (0.1 M HCl) were added to the first beaker, and 5 drops of a strong base (0.1 M NaOH) were added to the second beaker. The final pH’s were found and recorded, making it possible to determine that unknown solution 1 was buffered, since the pH barely changed from initial to final in both cylinders. In part two of the experiment, 0.7128 g of Unknown B weak acid was dissolved with water in a 100-mL volumetric flask, and 25.0-mL of that solution was pipetted into an Erlenmeyer flask. 0.1 M NaOH was slowly titrated into the unknown acid in 0.2 mL intervals for precision and accuracy. The pH was measured at each interval creating Chart 1, which displays the relationship between the added NaOH to the titrand and the pH of the solution.
The reaction between a weak acid and a strong base is shown in Chart 1, which quantitatively displays two equivalence points—points where the moles of added NaOH equal the moles of acid in the solution. The equivalence points, or first derivatives, are inflection points where the concavity of the titration curve flipped. The equivalence points were used to determine the pKa and the moles added of the titrant, which equaled the titrand, and made it possible to calculate the molecular mass of the unknown acid. The titration curve contains three regions with nearly flat gradually increasing slopes; the first two are called buffer regions, where the acid in the solution rapidly consumes the base—the titrand. Due to hydrolysis of the salt in the solution, the pH at the first equivalence point was still acidic with a pH less than 7. Since the titration curve displayed two equivalence points, the acid was diprotic. Since the moles of NaOH at the equivalence point is equal to the moles of acid in the solution, it was possible to calculate the molecular weight of the unknown, revealing that it was H2C4H2O4 – Maleic acid. The calculated mass of the unknown B was 118.8 g, which is close to Maleic acid—116.1 g, shown in Table 2.
The experimental goals were met. Solution 1a in Table 1 was a buffered solution, since the pH barely changed after adding a strong acid and a strong base to the pre-existing solution. The unknown acid in part 2 was found to be diprotic, and it was identified as Maleic acid after finding the pKas and calculating its molecular weight. The first source of error came from possibly misreading the molarity of the stock solution NaOH, which would make the calculations slightly less accurate. Another source of error came not precisely adding exactly 0.2 mL titrant each interval, which slightly flawed the titration curve. Although there was only 1% error, accuracy and precision could have been improved by precisely measuring correct amounts of solutions, and using the correct values for calculations.
The first part of the overall goal was achieved by concluding the unknown solution was buffered. The second part of the goal was achieved as well since the pKa(s) and molecular weight of the unknown acid were calculated, and the unknown was identified. To determine the unknown solution was buffered, the initial pH was recorded before adding a strong base and strong acid to a solution in two separate cylinders, and monitoring the slight pH change. The approach to the second part of the experiment was to titrate an unknown acid with a known concentration of the strong base, NaOH. Finding the equivalence point where the moles of NaOH was stoichiometrically equal to moles of unknown made it possible to calculate the molecular weight of the unknown acid. Referring to Table 2, the unknown acid B was concluded to be Maleic acid, H2C4H2O4.
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