Standardization of NaOH and Analysis of Unknown Acid Sample.

Procedure:

1. Clean and rinse all glassware that will be used in the experiment. Rinse buret and pipet with distilled water.
2. Clean and rinse the 1-L bottle and stopper. Label the bottle “0.1 M NaOH”. Put about 500mL of distilled water into the bottle.
3. Weigh out 4g (0.1 mol) of NaOH pellets and transfer them into the 1-L bottle. Stopper and shake the bottle to dissolve the sodium hydroxide. Once the NaOH has dissolved, fill the rest of the bottle with another 500mL of distilled water.
4. Set up the buret with the clamp and fill the buret with the NaOH solution
5. Find 3 of the Erlenmeyer flasks, label them numbers 1 through 3, and fill each with the KHP, between 0.6 and 0.8g of the for each of the flasks.
6. Tear the scale when each flask is placed on the weighing pad, making sure it zeroes out. Record the amount of KHP in each flask to the nearest 0.1 milligram.
7. At the lab station, add 100mL of water to sample 1. Add 2-3 drops of phenolphthalein indicator solution, and swirl to dissolve the solution completely.
8. Record the initial reading of the NaOH solution in the buret to the nearest 0.02mL, remembering to read across the bottom of the curved solution surface (meniscus).
9. Begin adding NaOH solution from the buret to the flask, swirling the flask constantly during the addition. It should require at least 20mL of NaOH to properly titrate the solution. The streaks of pink will begin to persist for longer periods of time.
10. Once the solution is a steady, pale pink (one which does not fade upon swirling), the solution has been properly titrated.
11. Repeat the titration of the remaining two samples of KHP.
12. Calculate the number of moles in each sample, given that the molecular weight of KHP is 204.2g.
13. Calculate the concentration of NaOH given the volume of NaOH used, and the fact that the reaction between NaOH and KHP is of 1:1 stoichiometry. Average the values together to find the concentration of the NaOH mixture.

Data and Observations:

Observed variances:

The cork broke in our 1-L bottle of NaOH, possibly tainting our results (after sample 2 was titrated). Sample 2 was a darker pink relative to flask 1 and 3. Flask 1 and 3 were relatively the same shade.

Calculations:

Trial I:         .637g KHP x (1 mol KHP / 204.23g per mol) =  .003119 mol KHP

                .003119 mol KHP x (1/ .03354 L NaOH) = .09283 M NaOH

Trial II:        .600g KHP x (1 mol KHP / 204.23g per mol) =  .002937 mol KHP

                .002937 mol KHP x (1/ .03123 L NaOH) = .09398 M NaOH

Trial III:        .598g KHP x (1 mol KHP / 204.23g per mol) =  .002928 mol KHP

                .002928 mol KHP x (1/ .03115 L NaOH) = .09399 M NaOH

                Average Concentration: 0.09370 M

Part II: Identification of H(a), The Unknown Acid

Materials:

• Buret and clamp
• 5-mL pipet and safety bulb
• 1-L glass or plastic bottle with stopper (filled with NaOH solution from part I)
• 3 Erlenmeyer flasks
• Digital scale
• Chemical spatula
• The Unknown Acid, ‘H(a)’
• Phenolphthalein indicator solution
• Lab grip and safety goggles

Procedure:

1. Find 3 of the Erlenmeyer flasks, label them numbers 1 through 3, and fill each with the unknown acid H(a), between 0.6 and 0.8g of the for each of the flasks.
2. Tear the scale when each flask is placed on the weighing pad, making sure it zeroes out. Record the amount of H(a) in each flask to the nearest 0.1 milligram.
3. At the lab station, add 100mL of water to sample 1. Add 2-3 drops of phenolphthalein indicator solution, and swirl to dissolve the solution completely.
4. Record the initial reading of the NaOH solution in the buret to the nearest 0.02mL, remembering to read across the bottom of the curved solution surface (meniscus).
5. Begin adding NaOH solution from the buret to the flask, swirling the flask constantly during the addition. It should require at least 20mL of NaOH to properly titrate the solution. The streaks of pink will begin to persist for longer periods of time.
6. Once the solution is a steady, pale pink (one which does not fade upon swirling), the solution has been properly titrated.
7. Repeat the titration of the remaining two samples of H(a).
8. Calculate the number of moles in each sample and the molecular weight of H(a), given the concentration of the NaOH standardized in part I.

Data and Observations:

Observed variances:

Flask 1 was pinker than both 2 & 3. 2 & 3 were roughly equal in tint.

Calculations:

Trial I:                .02728 L NaOH x 0.0937 M NaOH = .00255 mol NaOH or H(a)

0.529g H(a) x (1/.00255 mol H(a) ) = 207.54g/mol H(a)

Trial II:        .03138 L NaOH x 0.0937 M NaOH = .00294 mol NaOH or H(a)

0.617g H(a) x (1/.00294 mol H(a) ) = 209.86g/mol H(a)

Trial III:        .03478 L NaOH x 0.0937 M NaOH = .00326 mol NaOH or H(a)

                0.670g H(a) x (1/.00326 mol H(a) ) = 205.59g/mol H(a)

Average molar mass: 207.66g/mol H(a)

Conclusion:

While there were many mistakes made during the execution of this lab experiment, overall, the results did not deviate extremely far from the norm, or expected result. We determined that the unknown acid, H(a), was actually the acid used during the first set of titrations in Part I, potassium hydrogen phthalate. We were off by 3.46 g/mol, probably due to poor titration skill. The group let the solution become too pink, or too saturated. Another possible reason that the error can be attributed to was the fact that we filled the buret up to the 50mL mark, as opposed to filling it up partway and taking an accurate reading. While we let out enough of the NaOH solution to make it appear as if it were exactly at the 50mL mark, this was not the appropriate way to set up the titration lab. The colors were not uniform, and the average did not provide an accurate representation of the real-world concentration of NaOH nor the actual concentration of KHP in both instances, when it was known and unknown. Finally, the fact that cork dropped into our NaOH solution could possibly the culprit of the inaccuracies.