Data Sheets (To be attached to your lab report. Recopy if messy):

Unknown Letter: Triprotic Molarity of NaOH (mol/L) _______

Volume of unknown acid (mL):

Run 1

Run 2

Run 3

Initial volume NaOH

0.00

8.80

17.45

Final Volume NaOH

8.80

17.45

26.20

NaOH volume added (subtract B - A)

8.80

8.65

8.75

Moles NaOH added

Analysis:

Molarity of unknown acid (assuming the acid is triprotic):

Run 2 _ Run 3 _ Average _______

Sample Calculation of Molarity of unknown acid:

Post Lab Questions:

1. Why is it important to record the endpoint of the titration at the permanent faint pink color as opposed to the permanent dark pink color?

2. What would happen to your calculated acid concentration if your final solution was too pink?

3. Calculate the molarity of the unknown acid if the acid were:

triprotic

4. Why is the stir bar used?

5. Would it be possible to do this experiment without using an indicator dye? If so, explain how.

6. Is the titration technique limited to acid/base reactions?

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Nelly StrackeLv2

17 Dec 2019

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I am unable to solve these questions due to a time constraint. I received help on part 1 of the unknown acid. Need help with the rest! I really appreciate it!

Lab: ACID-BASE TITRATION

Objective: To use titration to determine in Part 1, the concentration (molarity) of an unknown acid solution and in Part 2, the purity of a sample of KHP acid. Background: Titration is a volumetric technique used to determine the concentrations of solutions, molar masses of solids, and purity of samples. A titration inolves the addition of a titrant (a solution of known concentration) to an analyte ( a solution of unknown concentration), or vice versa, using a piece of glassware called a burette. The titration is carried out until it reaches an equivalence point (the exact point at which the reaction between the two solutions is complete). A chemical indicator is often used to aid in the identification of the equivalence point. An indicator changes color upon reaching the equivalence point of endpoint. Since we will be carrying out an acid-base titration, the indicator must change color upon reaching the endpoint at a specific pH. An example of such an indicator is phenolphthalein, which changes from colorless to pink near the equivalence point when pH approaches 7. It should be noted that the indicator selected is dependent upon a given titration, so that the observed color change is close to the ideal equivalence point. Once the titration is completed, we can use the volumes measured for each solution, as well as the concentration of the titrant, to determine the concentration, molar mass, or purity of the unknown analyte. Procedure Note: During lab, record all measurements and data in a clearly labeled table. Part 1: 1. Obtain ~80mL of ~0.1M of base (sodium hydroxide, NaOH) into a labeled beaker. Make sure to report the actual concentration from the bottle. 2. Obtain ~25mL of unknown acid into a second labeled beaker. Record whether your acid is monoprotic, diprotic, or triprotic. 3. Using a funnel, or carefully pouring, fill the burette with sodium hydroxide NaOH. Make sure a labeled waste beaker is under the buretter to catch any drippings. Rinse your burette by filling it with ~5-10mL of NaOH. Turn the knob to let ~2 mL out the tip into the waste beaker. Make sure your waste beaker is under the burette to catch any drips. Next, turn the Burette upside down and pour the remaining NaOH into your waste beaker while twisting the burette to coat the inside with NaOH. 4. Fill your burette with sodium hydroxide. Turn the knob to fill the tip of the burette with NaOH. Record the initial volume to the correct number of significant figures. 5. Using a volumentric pipette, transfer 10.00 mL of your unknown acid into the Erlenmeyer flask. 6. Add 3 drops of pH indicator and magnetic stirrer to the Erlenmeyer. 7. Under the burette, place the Erlenmeyer on a hot plate and turn on the stir knob. Heat should be off. 8. Titrate the unknown acid until the solution turns pink and remains pink. 9. Record the final volume on the burette to the correct number of significant figures. 10. Repeat (4) to (7). Pre-plan the whole lab and who will do which task. MY RESULTS: PART 1 The actual concentration of NAOH is 0.5939, 80 mL Initial Final 24.9 mL 44.5mL 26.1 mL 45.57mL PART 2 Impure KHP (SAMPLE B ) Only Paper: 0.280 g Impure KHP Mass: 0.526 g Initial volume: 22.30 mL + water +26.10 mL Final Volume: 26.80 mL QUESTIONS: 1. Calculate the concentration (molarity) of your unknown acid. 2. Calculate how many grams in your impure KHP sample was indeed KHP. 3. Calculate the percent (by mass) KHP in the sample. Part 2: 1. What would happen to the final result( state what the final result is ) if the tip of the burette was not filled when reading the initial volume? 2. What would happen to the final answer if you used Ca(OH)2 instead of NaOH and DID NOT know it? 3. What would happen to the final result if the Erlenmeyer flask was contaminated with other types of acid? 4. Recalculate the concentration of your unknown acid if it was given as a diprotic acid. 5. Draw a molecular level picture (showing Na+, H+, OH-, H2O, etc...) of what is in your beaker before, at half point, at equivalent point, and at end point of the titration. Part 3: 1. What would have happened to the final result if the impurity in the sample was also acidic? 2. What would have happend to your final result if you used ~25mL of water to dissolve the solid acid instead of 20mL? Summary data: Create a clearly labeled table of all your final data results for each run, and athe average of multiple runs. Part 1: Unknown #, Vol base, M base, Vol acid, M acid. Part 2: Sample #, % KHP (purity)

Lab Results 1.For your coarse titration record the following results. volume of NaOH solution in the burette at the start (mL) volume of NaOH solution in the burette at the end (mL) volume of NaOH solution dispensed (mL) volume of HCl solution in the flask (mL) Data Analysis 2.Calculate the molarity of the HCl concentration using your coarse titration results. Experiment 2: Perform a Fine Titration Lab Results 1.For your fine titration record the following results. volume of NaOH solution in the burette at the start (mL) volume of NaOH solution in the burette at the end (mL) volume of NaOH solution dispensed (mL) volume of HCl solution in the flask (mL) Data Analysis 2.Calculate the molarity of the HCl concentration using your fine titration results. Conclusions 1.What do you expect to happen if you forget to add phenolphthalein to the Erlenmeyer flask? 2.Suppose you had to titrate a solution of NaOH of unknown concentration with a solution of HCl of known concentration using phenolphthalein as an indicator. How will you determine the color change? 3.Suppose you titrated a sample of acetic acid (monoprotic acid) with a 0.125 M solution of NaOH. Given the data in the table below, what is the concentration of the acetic acid solution? volume of 0.125 M NaOH dispensed 22.40 mL volume of acetic acid solution used 15.00 mL volume of water added to the HCl solution 15.00 mL

Question 1 and 2