Chem 1B Lab Saul Ruiz Experiment 8 Acid Base Equilibria During this lab we determined the dissociation constant of acetic acid by using a pH meter compared it to the actual Ka value which gave us a percentage of error In the last part of the experiment we also tested the resistance that buffer solutions had to small changes in pH CH3COOH aq H aq CH3COO aq NaCH3COO aq CH3COOH aq NaCH3COO aq CH3COOH aq HCl aq H aq Cl aq NaCH3COO aq Na aq CH3COO aq CH3COO aq H aq HCH3COO aq OH aq H aq H2O l In order to determine the dissociation constant we had to utilize a pH meter and constantly jot down data roughly taking readings after about every 1 mL As soon as the pH began to spike we decreased the increments to about 0 2 mL and wrote down the readings Through this method we were able to take our data and create a graphical representation with volume on the x axis and pH on the y axis This allowed us to determine the equivalence point as well as the half equivalence point which was useful for determining the pKa This is possible because in the Henderson Hasselbalch equation A pH p K a log At the half equivalence point the concentration of A is equal to that of HA and as such you are left with pH pKa With pKa we were able to obtain the Ka since pKa log Ka In parts two and three we were able to determine how buffered solutions compared to unbuffered solutions By making two solutions one with DI water and another with a buffer and adding NaOH and HCl to them it was a clear representation of how buffers are resistant to small changes in pH in comparison to their unbuffered counterparts The Ka value obtained in Part One was obtained by finding the equivalence point of NaOH by using the graph Once we found the volume relative to the equivalence point we halved it in order to order to get the half equivalence point With that said the pH at the half equivalence point is equal to the pKa and from there we determined the Ka value which turned out to be higher than the actual Ka value meaning that our solution was slightly more acidic In Part Two our Ka values were 1 01 x 10 4 and 3 14 x 10 5 The values stated were higher than the actual values which basically means that there was more error present For our pH values our results showed a 4 9 and 3 5 compared to 5 1 and 4 0 respectively Our first value was a little less than the calculated value but our second value was off by a bit more Various sources of error can take the blame in this experiment One of them would be the fact that we began by stirring the solution in a slow setting which may not have allowed the solutions to mix properly Another source of error could have possibly been the pH meter which acted up a few times and gave us false readings and then needed to be re calibrated This may have altered our data and may explain why some of our measured data is far off from the calculated data that is given