By Mikayla Messing 8/3/12 Section 623 Abstract To examine the effectiveness of buffers by titrating two sets of five different solutions using HCl and NaOH and monitoring the pH change of the various solutions. The data collected shows that the buffer systems made with sodium acetate and acetic acid were effect when titrated with the strong acid and the strong base. Comparison of all the solutions shows that the concepts of buffers holds true for the results from the experimentation. Introduction
The main objective of this lab was to test the ability of buffered and unbuffered solutions to resist changes in pH with the addition of strong acids and strong acid. This will be accomplished by making two sets of five different solutions. They will be made using water, a salt solution (sodium chloride), and various concentrations of a buffer. Once the solutions are made, one set of the five will be used to observe the changes in pH made by adding hydrochloric acid (HCl) drop by drop. The second set of the five solutions will be used to observe the changes in pH made by adding sodium hydroxide (NaOH) drop by drop.
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It is predicted that the first two solutions with water and the salt solution will experience immediate drastic changes in their pH. This will happen because neither solution has conjugates in their systems. The other three solutions will resist a change in pH until its buffer capacity is reached. The solution with the highest concentration of the buffer will be the most effective at resisting pH changes. This is because buffers become more effective as their concentration increases. Therefore, the buffer system with the greatest amount of sodium acetate will be the most effective buffer.
Method Five different buffer solutions were made twice to make ten buffer solutions; one set was be used to observe the changes in pH caused by adding a strong acid (HCl) and the other set was used to observe the changes in pH caused by adding a strong base (NaOH). The five solutions made were distilled water, . 10 M sodium chloride (NaCl), a solution of 1 gram of solid sodium acetate (CH3COONa) and 50 milliliters (mL) of acetic acid (CH3COOH), a solution of 5 grams of solid sodium acetate and 50 mL of acetic acid, and a solution of 10 grams of solid sodium acetate and 50 mL of acetic acid.
The specific contents of the ten beakers are as follows: beakers 1 and 6 contained 50 mL of distilled water. Beakers 2 and 7 contained 50 mL of a NaCl solution. Beakers 3 and 8 contained 1 gram of solid sodium acetate and 50 mL of acetic acid. Beakers 4 and 9 contained 5 grams of solid sodium acetate and 50 mL of acetic acid. Beakers 5 and 10 contained 10 grams of solid sodium acetate and 50 mL of acetic acid. A special probe was used to monitor the pH of one set of each solution as HCl was added drop by drop.
The results were recorded in a table found later on in the report. The probe was also used to monitor the second set of solutions as NaOH was added drop by drop. For more detailed instructions on the instructions for mixing the solutions and other experimental details, refer to Chemistry 110 lab manual by Kautz, Kinnan, and McLaughlin. Results The series of experiments were conducted and the pH was recorded with every drop of the strong acid or strong base. It was revealed through the experimentation that the distilled water and sodium chloride were not effective buffers.
The solutions of sodium acetate and acetic acid were good buffers; the effectiveness of the solution increased as the concentration of sodium acetate increased. The solution with 10 grams of sodium acetate was the most effective of the buffers, followed by the solution made with 5 grams of sodium acetate and the solution made with 1 gram. Effectiveness of these solutions were shown through the addition of both the HCl and the NaOH. Table 1. 1 below shows the change in pH when beakers one through five were titrated with 1. 0 M hydrochloric acid. The Effect of Acid on Solutions|
Beaker| 1| 2| 3| 4| 5| mL of HCl| pH| pH| pH| pH| pH| 0| 6. 82| 5. 23| 4. 46| 5. 58| 5. 85| 1| 1. 16| 1. 06| 4. 33| 5. 43| 5. 85| 2| 0. 966| 0. 885| 4. 24| 5. 28| 5. 73| 3| 0. 668| 0. 665| 4. 1| 5. 15| 5. 55| 4| 0. 658| 0. 585| 3. 78| 5. 04| 5. 53| 5| 0. 548| 0. 335| 3. 37| 4. 94| 5. 51| 6| 0. 538| 0. 295| 3. 04| 4. 85| 5. 49| 7| 0. 438| 0. 165| 1. 64| 4. 75| 5. 48| 8| 0. 358| 0. 125| 1. 06| 4. 72| 5. 44| 9| 0. 238| 0. 125| 0. 856| 4. 62| 5. 42| 10| 0. 128| 0. 115| 0. 686| 4. 53| 5. 42| Table 1. 1 shows the effect of adding HCl on the pH of the beakers. Table 1. below shows the change in pH when beakers six through ten were titrated with 1. 0M sodium hydroxide. Beaker| 6| 7| 8| 9| 10| mL of NaOH| pH| pH| pH| pH| pH| 0| 6. 34| 8. 173| 5. 128| 5. 762| 5. 73| 1| 11. 01| 11. 39| 5. 048| 5. 772| 5. 96| 2| 11. 68| 11. 76| 5. 298| 5. 912| 6. 17| 3| 11. 92| 12. 14| 5. 368| 6. 012| 6. 33| 4| 12. 06| 12. 33| 5. 488| 6. 142| 6. 53| 5| 12. 37| 12. 35| 5. 928| 6. 382| 6. 59| 6| 12. 41| 12. 41| 11. 07| 7. 642| 6. 78| 7| 12. 56| 12. 46| 12. 26| 12. 11| 7. 14| 8| 12. 62| 12. 55| 12. 37| 12. 45| 12. 63| 9| 12. 65| 12. 57| 12. 54| 12. 71| 13. 33| 10| 12. 71| 12. 67| 12. 2| 12. 78| 13. 58| Table 1. 2 shows the effect of adding NaOH on the pH of the beakers. Table 1. 1 and 1. 2 show that the pH of distilled water and sodium chloride changed dramatically with the addition of just one mL of the hydrochloric acid or the sodium hydroxide. When milliliters of either HCl or NaOH were added to the three solutions of various amounts of sodium acetate and acetic acide, the pH change was very small until its buffer capacity was breached, at which point the pH of the buffer solution changed drastically to a new number. These results can also be seen below in Graph 1. and 1. 4. These show the pH of all the solutions vs. volume of HCl and NaOH, respectively. Graph 1. 3 shows the pH of the solutions vs. volume of added HCl The average change in the pH of the solution with 1 gram of sodium acetate when HCl was added was . 377 per mL of HCl.. With the solution of 5 grams of sodium acetate, the average change in pH was . 104 per mL of HCl.. The solution with 10 grams of sodium acetate had an average change in pH of . 043 per mL of HCl. Graph 1. 4 shows pH vs. volume of added NaOH The average change in pH for the solution with one gram of sodium acetate was . 28 pH per mL of NaOH. The solution with five grams of sodium acetate had an average change in pH of . 711 per mL NaOH. The final solution with ten grams of sodium acetate had an average change in pH of . 785 per mL NaOH. Discussion Buffers are solutions that resist change in pH when an acid or base is added to it. The buffers in this experiment had a weak acid and its conjugate base, which is how most buffers are made. When HCl was added to the buffer solution, a reaction occurred and the strong acid was neutralized by CH3COO-. The buffer reacts in a similar way when NaOH was added.
The strong base reacted with the buffer and was neutralized by CH3COOH. As shown in the results of the experiment, the most effective buffer was the solution that contained ten grams of sodium acetate and fifty milliliters of acetic acid. It was the most effective buffer for both the NaOH titration and the HCl titration. The prediction at the beginning of this report was correct. While the solution with ten grams of sodium acetate had a greater average increase in pH over the solution with five grams of the material in the trial with NaOH, it was still a more effective buffer because its capacity was greater.
It took more milliliters of NaOH for the solution with ten grams to surpass its buffer capacity. As far as buffer effectiveness goes, the closer the concentrations of the acid and base in it are, the greater its buffer capacity. The closer that concentration is to one, the better the buffer will be. It can be calculated by using the Henderson-Hasselbalch Equation: pH = pKa + log([base]/[acid]) This equation can be used to find the Ka value, from which the acid concentration and pH of solution can be determined before any acid or base is added to the buffer.
The equation works under the assumption that x is small, if x is not small, this equation will not yield the correct concentration or Ka values that would come from the use of a different method. The concentration of the CH3COO- in beaker five is . 989 M. When this concentration is very close to the concentration of the other buffer component in a way that the ratio of the two is close to one, the buffer will be more effective. From all the data, it is clear to see that the solutions with ten grams of sodium acetate and fifty milliliters of acetic acid, beakers five and ten are the most effective buffers.
It is interesting to note that the three buffer systems are more effective at resisting pH changes when acid is added than when a base is added. It took more milliliters of HCl for the buffer capacity to be surpassed than it did for the NaOH. This is true for all three solutions made with varying amounts of sodium acetate. Distilled water and sodium chloride are not worth much discussion; when HCl or NaOH was added to these solutions; the pH of the mixtures changed drastically with the very first drop of the strong acid or strong base.
The solutions with the sodium acetate and acetic acid were significantly more effective buffers; they resisted changes in pH, and it took several milliliters of the acid or base to surpass the buffer capacity to make the pH change dramatically. A source of error in this experiment could be the molarity of the HCl or the NaOH. During the middle of experimentation, the amount of HCl had dwindled and needed to be replaced; considering that there was no 1. 0M HCl available in the lab that day, it is possible that the molarity of the acid changed during the middle of the experiment.
It is possible that the readings taken from the pH probe were not accurate due to constant fluctuation in its readings. The number recorded was the reading the probe landed on the most, not necessarily when the reading was stable. The sodium acetate may not have been completely dissolved in the acetic acid, so its concentration may not have been what it would come out as mathematically. Conclusion This lab reiterated the concepts behind buffers and how important it is to understand the concepts behind the math that comes with the chemistry.
Minimizing errors in the lab is crucial to being able to understand every bit of what is going on with the chemistry. An error can completely change the results and make something difficult to explain happen in the trends of the experiment. The lab was very beneficial in furthering the understanding of buffers and how they work outside of the mathematical equations. References Kautz, J. , D. Kinnan, and C. McLaughlin. Chemistry 110 Laboratory Manual Hayden-McNeil, 2010. Print. Tro, Nivaldo. Chemistry: A Molecular Approach. Upper Saddle River, NJ: Pearson Education, Inc. , 2008. Print.