Acid-Base Equilibria: Determination of Acid Ionization Constants
Introduction:
The purpose of this lab is to learn about acid ionization constants and buffer solutions. We will be determining the acid ionization constant by finding pH with pH meters. In part three we will prepare a buffer solution and then observe the change in pH as either an acid or base is added.
Data:
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Discussion:
The majority of this experiment was focused on determination of Ka, or the acid ionization constant. To find Ka we were given concentrations of each chemical in the solution and its concentration individually. We then determined the pH, and solved using an ICE table.
In part 1, we found the half equivalence point by titrating HA with NaOH, plotting the graph with excel, and then visually finding the halfway mark. The pH at this point was 4.26, which tells us that Ka = 5.50 x 10-5 by knowing pH = pKa and Ka = 10-pKa at the half equivalency point. This means that HA has very little dissociation in an aqueous solution, which is expected because it is a weak acid.
In part 2, we were yet again attempting to find the disassociation constant of acetic acid in different solutions by knowing volume, concentration and pH. We found three very different constants, an error which will be covered later. Actually completing these steps was very elementary and does not need to be discussed.
The Essay on Solution Preparation And Standardization
Generally, there are two ways in preparing a solution, one is by dissolving a weighed amount of solid in a required solvent and the other is by dilution of a concentrated solution into the desired concentration. In diluting concentrated solution, the concentration of the diluted solution can be determined by standardization. To standardize a solution, we will need to perform titration. In this ...
Part 3 was the second half of this lab. We created a buffer solution using 30mL each of 1.0M NaA and HA, and the pH was recorded. We then divvied this into two beakers. 2mL 1.0M HCl was added to one and 2mL 1.0M NaOH was added to the other. The pH level was recorded for each of these solutions. The original buffer solution had a pH of 4.49. Solution 1 with HCl was 4.30, and Solution 2 with NaOH was 4.55. Previous to this we created an unbuffered solution with DI water and added the same amounts of acid and base to the solution. For HCl, there was a negative 5.62pH change and positive 4.6 for NaOH. In the buffered solution, the changes were: HCl – (-0.19); NaOH – (+0.06).
This demonstrates the degree to which a buffer solution can resist change.
All Ka values determined are far off from the actual 1.76×10-5. I believe a very large factor in this is the measuring device itself. From the beginning there were numerous problems with the pH meters, both with working at all and reading incorrect values. To prove the extent to which the meters were not accurate, in part 2 we conveniently were required to determine the %error. In one solution there was a 61.4% error and in another it was 32.0%. This can very easily throw off all the calculations for the Ka constant. In part 1, either my partner or I made an error in adding 1.0M HA instead of 0.1M HA, this caused us to be unable to titrate to the equivalence point with even 50mL 0.1M NaOH.
