Titration of an unknown concentration sample of HCl with a 0.15 M sample of NH4OH

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Chemistry 10 St. Augustine Preparatory School

Managua, Nicaragua May 20, 2016

Introduction Titration takes place when one solution, of which the concentration is known, is added to another solution, of which the volume is known but the concentration is not. This process occurs until the reaction neutralizes. However, one must be diligent because the measurements must be precise, otherwise the endpoint could quickly surpass the desired solution and its pH level. In this titration lab, a strong acid with an unknown concentration, hydrochloric acid, and a weak base with a concentration of 0.15 M ammonium hydroxide, will be provided. As the process transpires, the change in pH will be shown through an acid-base indicator. Common indicators are Phenolphthalein, Bromothymol blue, and Methyl red. For this titration, Methyl orange will be used. This is due to the fact that a strong acid and weak base are being used and Methyl orange will change color at a pH range of 3.1-4.4. Objective The objective of this lab is to find the unknown concentration of the hydrochloric acid, given a solution of ammonium hydroxide with a concentration of 0.15 M. Hypothesis If twice the amount of acid is added as base, then the solution should neutralize and turn into an orange color. Materials Safety glasses Lab apron Rubber gloves Ipad to take notes Pipette Erlenmeyer flask Eyedropper 250mL beaker 100mL beaker Ammonium hydroxyde Hydrochloric acid Methyl orange Burette Burette stand Funnel

Procedure 1. To ensure safety, put on a lab apron, safety glasses, and rubber gloves. 2. Clean the burette by adding 15 mL of NH4OH and turning the burette around so that it reaches all of its inner surface area. Place and tighten the burette on the burette stand. 3. Using the pipette, obtain 20 mL of hydrochloric acid and add it into the Erlenmeyer flask. 4. Using an eyedropper, obtain Methyl orange and add 4 drops to the hydrochloric acid in the Erlenmeyer flask. Stir cautiously. 5. Position the Erlenmeyer flask below the burette so that the tip of the burette is inside the flask. 6. Fill up the 100mL beaker with NH4OH. 7. Place a funnel on the upper part of the burette and pour the NH4OH until it reaches the 0mL mark. 8. Gently open the stopcock, so that a light stream is released into the Erlenmeyer flask. 9. Observe for any color changes in the solution. When noticed, close the stopcock to the point that small drops are released into the Erlenmeyer flask. Stop when the solution reaches the desired color or pH level. For this lab, the desired color is orange. 10. Record the amount of base added to the hydrochloric acid. 11. Pour the solution into the 250mL beaker which will serve as your waste container. 12. Repeat steps 3-5 and then 8-10 at least 3 times, in order to compare and achieve more accurate results.

Observations When the stopcock was slightly opened, a small stream of NH4OH was being added into the Erlenmeyer flask. For the first 10mL added, there were no color changes observed in the hydrochloric solution. Nevertheless, between the addition of 12-14mL a back and forth color change of red and orange was identified. By around 15.3mL, a pale orange color was observed in the erlenmeyer flask. However, the change was rapid and the color quickly became a strong orange. In the first trial, a few more drops were added when the hydrochloric solution was a strong orange and by around the addition of 17.7mL, it became yellow. These observations are shown in the pictures below:

Data, results, and calculations Trial 1 Trial 2 Trial 3

Final Burette reading(mL) 18.51 35.00 32.19

Initial Burette reading(mL) 0.2 18.51 14.81

Total volume added 18.31 16.49 17.39

Color of indicator Yellow Orange Orange Concentration calculation: HCl + NH4OH ---- > HOH + NH4Cl n = (0.002541mol)(1/1) | Concentration n = 0.002541mol | of NH4OH: 0.15 mol/L | C = n / v | n = (0.15mol/L)(average of good trials) C= 0.002541mol/0.02L | n = (0.15mol/L)(16.49 + 17.39 / 2 mL) Concentration of | n = (0.15mol/L)(16.94 mL) HCl: 0.12705 mol/L | n = (0.15mol/L)(0.01649 L) | n = 0.002541 mol Analysis In this titration lab, three trials were done, of which only two neutralized and maintained an orange color. From the endpoints of these two successful trials, an average of the volume of the base added was calculated. The data recorded, clearly shows that the 20 mL hydrochloric solution will neutralize with the addition of about 16.94 mL of ammonium hydroxide. This result proves the hypothesis to be false. The hypothesis predicted the need of much more weak base than strong acid in order to neutralize the solution. However, the data concludes that about only ¾ of the amount of acid is needed as base to neutralize the solution. A key factor that contributed to the inaccuracy of the hypothesis, is that it was solely based on the strength of the acid and base, rather than their concentrations. After calculating the concentration of hydrochloric acid, it is observed that it has a concentration of 0.12705 mol/L. On the other hand, ammonium hydroxide has a concentration of 0.15 mol/L. Therefore, since ammonium hydroxide has a higher concentration of moles per liter, it will need less volume than hydrochloric acid to reach an equivalence point. Nevertheless,this chemical equation saw a titration between a strong acid weak base resulting in the product of NH4Cl which is slightly acidic, hence why the endpoint of this titration is below a pH level of 7. Throughout this lab, there were many factors which could have caused certain inaccuracy and mismeasurement. First, students were provided with a time span of 50 minutes to complete this lab. This included the cleaning and organization of materials for each group. At the same time,

students were desired to complete at least 3 or 4 trials of the experiment. Although most groups were capable of reaching such goal, it required them to work with a certain speed. Under a longer time period, measurements could have been handled more carefully and the lab overall could have been more precise. Second, students were placed in groups of up to 5 people. This required all students within each group to be very exact when it came to the recording and transmission of data. Any misunderstanding of measurements between members could affect the whole group. Third, only three trials were achieved with this particular group. Of these three, only two reached the desired endpoint. For comparison and better accuracy of data, more trials should have been done, of which at least five reached the desired endpoint. Conclusion For this lab, students were provided with a solution of hydrochloric acid and their objective was to determine the concentration of this sample. To help reach this purpose, students were given the concentration of the weak base in the titration, ammonium hydroxyde. After three trials were completed, the first trial was identified as exceeding the desired endpoint with a yellow color, while the other two maintained the desired pH level and orange color. These two successful trials showed that less volume of the weak base was needed than that of the strong acid to reach an equivalence point. Therefore, the hypothesis was incorrect. This is due to the fact that the student did not take into account the importance of concentration in titration. With the use of the average of ammonium hydroxide volume added to the hydrochloric acid in the two successful trials and the concentration of ammonium hydroxyde provided, 0.15 mol/L, the concentration of hydrochloric acid was calculated to be 0.12705