L ABORATORY INVESTIGATION TITRATION OF AN UNKNOWN ACID

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Finding the Size of a Molecule:

LL ABORATORY INVESTIGATION TITRATION OF AN UNKNOWN ACID L ABORATORY INVESTIGATION TITRATION OF AN UNKNOWN ACID aboratory Investigation:

Titration of an Unknown Acid





Purpose:

To prepare a solution of NaOH, and to standardize that solution.


To use titration to determine the concentration of a solution of hydrochloric acid of unknown concentration.


Introduction:

In acid-base titration, an acid and base are reacted in order to determine the concentration of one or the other. In this lab we will react a solution of NaOH that we prepare with solid oxalic acid, H2C2O4, and with acid of an unknown concentration. The titration with the solid oxalic acid will be used to determine our NaOH solution concentration. Once we have standardized the NaOH solution, we will then determine the unknown acid concentration.


Materials:


Solid NaOH Oxalic Acid Dihydrate(H2C2O4. 2H2O) Buret

Volumetric Flask Erlenmeyer flask Wash Bottle Indicator Solution Hydrochloric Acid of unknown concentration


Procedure:


A. Preparation of a Solution of NaOH


1. Obtain a 100 mL volumetric flask and fill it half way with deionized water.


2. Calculate the mass of NaOH needed to make 100.0 mL of a 0.250 M solution of NaOH. Check your answer with your instructor before proceeding.


3. Mass out as close to this value as you can using a piece of weighing paper for the solid. Record the exact mass of NaOH used on your data sheet. You will use your this value as your “true value” to determine the percent error of your NaOH solution.


4. Immediately add your solid NaOH to the water in the volumetric flask. Put the top on the flask and mix as directed until the base has dissolved.


5. Use your wash bottle and add deionized water until the meniscus of the solution reaches the line on the flask.


6. Put the top on the flask and mix the solution again.





B. Standardization of the NaOH


1. Mass out as close to 0.20 g of oxalic acid dihydrate (H2C2O4·2H2O) as you can using a piece of weighing paper. Record the exact value used on your data sheet. Do not exceed .25 g of solid acid!


2. Add the oxalic acid dihydrate to an Erlenmeyer flask. Add approximately 25 mL of deionized water to the flask, and dissolve the acid by swirling gently. If not all of it dissolves, it will dissolve as the titration proceeds.


3. Rinse the buret on the “B” side with deionized water. Add the NaOH solution (from Part A) to the buret. Fill the buret to as close to the 0.00 mark as you can. Open the stopcock and let the tip fill with NaOH solution. Read the buret to the closest 0.01 (remember to read the buret directly from the top down) and record the reading on your data sheet.


4. Add 3 DROPS (too much indicator will cause cloudy solutions!) of the indicator to the flask, and titrate the solid acid with your NaOH. Try to obtain the lightest permanent shade of pink in the flask that you can. Use your wash bottle to rinse the sides of the Erlenmeyer. Remember to record starting and final volumes for the base on your data sheet.


5. Repeat the titration using a second sample of oxalic acid dihydrate. Again, record the mass of the solid, and starting and initial volumes of the base on your data sheet. If these two trials are off by more than 10%, repeat the titration, and then average the two closest trials.


C. Determination of the Molarity of an Unknown


1. Obtain a solution of HCl with an unknown molarity from your instructor. Note the letter for the unknown on your data sheet.


2. Rinse the buret on the “A” side with deionized water. Add the HCl solution to this buret until it is about half full. Again, open the stopcock to fill the tip with acid solution. Read the buret to the closest 0.01. Read the level on your base buret.


3. Put between 10 and 11 mL of HCl into the Erlenmeyer.


4. Add 3 drops of the indicator to the flask, and titrate HCl with your NaOH. Try to obtain the lightest permanent shade of pink in the flask that you can.


5. If your solution turns dark pink, add acid dropwise until it turns clear again. Then, titrate with base until you achieve a light pink color.


6. Remember to record starting and final volumes for the acid and base on your data sheet. Repeat the titration until you obtain at least two good endpoints.







Calculations:


Part B.


1. From the mass of the solid oxalic acid dihydrate, calculate the number of moles of acid used in each titration. Hint: since this is a dihydrate, remember to include the mass of water!


2. Using the balanced equation for the neutralization (below), calculate the number of moles of base used in each titration.


H2C2O4 + NaOH ?


3. Knowing the volume of the NaOH used, calculate the molarity of the NaOH for each trial.


4. Average the two trials to find the average experimental value for your NaOH.


5. Calculate the value for the molarity of the NaOH using the mass of NaOH you measured. Calculate percent error using this as your true value.


Part C.


1. For each trial, calculate the volumes of acid and base used.


2. Use these volumes and the molarity of your base obtained in the calculations above to calculate the molarity of the unknown acid.























DATA SHEET


Part A.


Mass of NaOH used in solution




Part B.



Trial #1

Trial #2

Trial #3

Mass of Oxalic Acid




Initial Buret reading




Final Buret reading




Volume of NaOH required





Part C.


Unknown Acid________________________________________________

Trial #1

Trial #2

Initial Buret reading—acid



Final Buret reading—acid



Initial Buret reading—base



Final Buret reading--base







Lab Writeup in Lab Notebook:


1. Data Tables, including units.

2. Calculations—neatly written out, including proper sig figs and units.

3. The answer for the molarity of your unknown acid—circle this number on bottom of page of calculations. There will be an accuracy grade for this part of the lab!

4. Error Analysis

5. Conclusion


4



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12 ĆWICZENIE LABORATORYJNE NR 4 (W24) 4 022010


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