Titrations of Acids and Bases: Lab Report

Categories: Chemistry

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Introduction

Acid-base titrations are widely used in chemistry to determine the unknown concentrations of acidic or basic substances through neutralization reactions. A neutralization reaction involves the combination of an acid and a base, resulting in the formation of water and a salt. The essential feature of this process is the combination of hydrogen ions (H+) with hydroxide ions (OH-) to produce water.

In a typical titration, an analyte (the solution with an unknown concentration) is reacted with a reagent (the solution with a known concentration).

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To prepare the analyte, the substance being studied is dissolved in a solution. An indicator, such as Phenolphthalein in this lab, is added to help indicate the endpoint, which is the point at which the color change occurs. The equivalence point is reached when the concentrations of H+ and OH- ions are equal, and the half-equivalence point is halfway to the equivalence point.

Titrations have significant real-world applications, including chemical analysis in laboratory medicine, pharmaceutical development, and the determination of water content, fat content, and vitamin concentrations in various products.

Chemical Reaction (Reaction 1):

( 8 4 4( ) + ( ) → 2 ( ) + 8 4 4 ( ))

Experimental Methods

The procedures and materials used in this experiment are detailed in the Principles of Chemistry II Laboratory Manual for Emmanuel College (2019) under "Titration of Acids and Bases." The primary modification made in the experiment was the addition of an extra trial, resulting in four trials instead of the usual three. The burette was refilled to 0 mL after each reading, maintaining the consistency of the procedure (Emmanuel College, 2019).

Results/Data

Table 1 – Recorded Masses of KHP and NaOH Solution Volumes

Trial	Mass of KHP (g)	Final Burette Reading (mL)	Volume of NaOH Used (mL)	Volume of NaOH Used (L)	Concentration of NaOH (M)
1	0.5132	24.81	24.81	0.02481	0.1013
2	0.5134	32.19	32.19	0.03219	0.07812
3	0.5136	33.21	33.21	0.03321	0.07574
4	0.5138	32.74	32.74	0.03274	0.07458

Average Concentration of NaOH: (0.0824 , M)

Calculation for Table 1:

Volume of NaOH Used (L) = Volume of NaOH Used (mL) * (frac{1 , mL}{1000 , mL})

Concentration of NaOH (M) = (frac{{text{mass of KHP (g)}} times frac{1 , text{mol KHP}}{{text{molar mass KHP (g/mol)}}} times frac{1 , text{mol NaOH}}{{1 , text{mol KHP}}} times frac{1 , text{L}}{{text{volume of NaOH used (L)}}})

Table 2 – Equivalence Point and Half-Equivalence Point Data

Trial	Equivalence Point Volume (mL)	Half-Equivalence Point Volume (mL)	pH = pKa at Half-Equivalence Point
1	17.02	8.51	4.01
2	17.40	8.70	4.12

Table 3 – Average Concentration of Unknown Acid #1

Trial	pH = pKa	Ka	Average Ka	Vol (mL) of Base at Eq Point	Concentration of NaOH (M = mol/L)	Volume of Acid #1	Molarity (M)
1	4.01	9.77 x 10^-5	8.68 x 10^-5	17.02	0.0836 M	25 mL	5.69 x 10^-2 M
2	4.12	7.59 x 10^-5	8.68 x 10^-5	17.40	0.0836 M	25 mL	5.82 x 10^-2 M

Discussion & Conclusion

The objective of this laboratory experiment was to determine the concentration of a sodium hydroxide (NaOH) solution through standardization and to find the pH and ionization constant (Ka) of an unknown weak acid (#1). The experiment yielded successful results, with the average concentration of NaOH determined to be (8.24 times 10^{-2} , M), the Ka value as (8.68 times 10^{-5}), and the average concentration of the unknown acid as (5.76 times 10^{-2} , M).

The hypothesis that the concentration of NaOH would be approximately (0.1 , M) was confirmed, as the average concentration obtained from the experiment was (0.0824 , M), which is in close proximity to the hypothesis.

While the experiment was conducted with precision, minor random errors may have occurred during the reading of the burette or when noting significant figures. To enhance precision in future experiments, careful observation of the meniscus at eye level and reporting measurements with four significant figures is recommended.

Systematic errors may have arisen from the visual detection of the endpoint, as the observable endpoint volume was greater than the equivalence point, resulting in a higher concentration value. Additionally, inaccuracies in the measurement scales for KHP could have affected the results. To mitigate systematic errors, multiple trials should be conducted to ensure accurate endpoint observation, and calibration of equipment is essential to maintain accuracy.