Experiments on neutralization titration and determination of the ionization constant of acetic acid

Summary

This experiment is from the official website of Southern Medical University

Operation method

Experiments on neutralization titration and determination of the ionization constant of acetic acid

Principle

HAc is a weak acid in solution with the following equilibrium: [H+], [Ac-] are both concentrations at equilibrium, and Ki is the ionization equilibrium constant, which is a constant at a given temperature. Based on that, the degree of ionization a (c is the total concentration of HAc) can be found. The ionization equilibrium constant, Ki, and the degree of ionization, a, can be found by determining the pH of acetic acid at a known concentration.In this experiment, a known concentration of NaOH solution was used to determine the total concentration of the HAc solution, and a pH meter was used to determine the pH of the HAc solution. General acids and bases often cannot be directly formulated into solutions of precise concentration because they contain impurities or are inherently unstable (e.g., HCl is volatile and NaOH is prone to water and CO2 absorption). In order to find out their exact concentration, they need to be calibrated with reference substances. In this experiment, potassium hydrogen phthalate was used as a reference substance to calibrate NaOH solution. The ionization constant of phthalic acid K1=1.1×10-3 and K2=3.9×10-8.

Materials and Instruments

Potassium hydrogen phthalate reference NaOH solution HAc solution NaAc solution Phenolphthalein solution Potassium hydrogen phthalate standard buffer
Analytical balance Bench scale Weighing flask Conical flask Alkaline buret titration table Volumetric flask Beaker Pipette Graduated pipette Wash bottle pH meter

Move

I. Calibration of NaOH solution

Accurately weigh the potassium hydrogen phthalate 0.40-0.45g three, respectively, placed in a conical flask, add water to dissolve, and then add phenolphthalein solution 2 drops, with the NaOH solution to be calibrated to the end of the drop until slightly red. Note down the volume of NaOH solution used and calculate the concentration of NaOH solution (take four significant figures) and the relative average deviation of the titration.

II. Determination of the concentration of HAc solution with NaOH standard solution

Pipette 25 mL (or 20 mL) of HAc solution of unknown concentration into a conical flask, add phenolphthalein, and titrate to the end point with NaOH standard solution. Calculate the concentration of HAc solution (take four significant digits).

Titrate three times in parallel. Calculate the average concentration of HAc solution and the relative average deviation of the titration.

III. Preparation of HAc solutions of different concentrations

Pipette 25mL, 5mL and 2.5mL of HAc solution which has been measured accurately, respectively, in three 50mL volumetric flasks, dilute to the scale with distilled water and shake well to make 0.05 ^mol-L-1, 0.01 ^mol-L-1 and 0.005 ^mol-L-1 HAc solutions.

IV. Determination of pH of 0.1 ^mol-L-1, 0.05 ^mol-L-1, 0.01 ^mol-L-1 and 0.005 ^mol-L-1 HAc solutions

Using four beakers, take a certain amount of each of the above four concentrations of HAc solutions and determine their pH values from dilute to concentrated using a pH meter and record the temperature (room temperature).

The measured data and calculations are presented in the table below.

HAc solution	c	pH	[ H+ ]	Ki	a
1
2
3
4
5

v. homoionic effects

Pipette 25 mL of 0.1 ^mol-L-1 HAc solution in a 50 mL volumetric flask, add 5 mL of 0.1 ^mol-L-1 NaAc solution, and dilute to the scale for solution No. 5. Measure its pH and calculate the ionization constant and degree of ionization of HAc according to the above experimental method, fill in the above table with the experimental results, and explain the experimental results by comparing it with solution No. 2.

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