Laboratory Report on Acid – Base Titration Joenar R. Alontaga An acid-base titration is an experimental procedure used to determine the unknown concentration of an acid or base by precisely neutralizing it with an acid or base of known concentration. Objective/s: to determine the total Alkaline Strength of unknown sample to perform acid-base titration Materials: concentrated HCl, solid NaOH, primary standard grade Na2CO3, methyl red indicator, unknown soda ash sample Equipment: stirring rods, beakers (50 mL, 1000 mL), Erlenmeyer flask, burette, glass bottle, plastic bottle, weighing boats, funnel, graduated cylinder, watch glass, burette holder, digital weighing scale Procedure: Prepare all the necessary materials. Wash all needed materials and rinse thoroughly with distilled water. Materials was obtained by the student from a person lab. The 3 samples were prepared and weighed by the student. A. Preparation of Solutions: 1. Prepare 1 L of 0.1 M HCl from concentrated HCl solution (12M). 2. Prepare 500 mL of 0.1 M NaOH from solid NaOH. 3. Use distilled water which has been recently freed from CO2 by boiling for the preparation of the solutions. 4. Show all necessary computations needed for the above preparations in your journal. B. Volume Ratio Determination: 1. Rinse a previously cleaned 50-mL acid burette at least four times with 10 mL- portions of the prepared HCl, allowing some of the liquid to run through the tip of the burette. Discard each portion; do not return any of the rinsing to the stock solution in the bottle. 2. Then fill the burette with the acid solution. In all these steps, pour directly from bottle to burette (using a funnel to avoid spillage) and place the burette in a burette holder. 3. Using the prepared NaOH solution, carry out the same procedure with the previously cleaned base burette. A strong alkaline solution should not be allowed to remain for a long time in a glass-stoppered buret because it tends to cement the stopcock to the cavity. If desired, a plain buret or a Teflon stopcock buret maybe used for the NaOH solution. 4. When the burets are ready for use and any air bubbles have been displaced from the tips, run out each solution to a point where the bottom of the meniscus is just below the zero mark. Allow the liquids to stand for a half-minute, remove and discard any hanging drops and read the exact positions of the meniscus of the liquid in each buret. Record the readings in the journal. 5. Run from the buret into a 250-mL beaker (or flask) about 35 mL of the base and add about 40 mL of the distilled water (which has been recently freed from CO2 by boiling and cooled to room temperature). 6. During titration carry it out in a well-lighted place with a white background (e.g. a towel or sheet of paper) below and in the back of the solution being titrated, in order to precisely ascertain the end point. If a beaker is used, have the tip of the buret close to the surface of the liquid in the beaker and use a stirring rod to stir gently the liquid without spattering during titration. If a flask is used, do not use a stirring rod, but gently rotate the flask during titration to give a swirling motion to the liquid. 7. Add 2-3 drops of methyl red (MR) indicator to the base solution to give a distinct yellow color. 8. Cautiously run the HCl from the buret to the base until the yellow color of the indicator changes to pink. 9. Wash down the sides of the flask or beaker with a little amount of water from a wash bottle. Add base from the buret to restore the yellow color, then run in HCl from the buret to the base solution, up to the point where the solution is of an orange shade, intermediate between pink and yellow. 10.In reaching this point, any hanging drops from the burets should have been transferred to the solution by means of a stirring rod or by contact with the inside of the flask. 11.Wait 1⁄2 minute for the liquids on the inside of the burets to drain, and then take the final readings for both acid and base. Record the values in your journal. 12.Refill the burets and repeat the titration. 13.For each trials, calculate the volume ratio of the HCl to that of the NaOH solution; also calculate the volume ratio of the alkali to that of the acid. 14.Check results of the three titrations should not vary by more than two parts in one thousand (2 ppt). If the variation in results is greater than this, repeat the titration until satisfactory values are obtained. Calculate and report the mean value. C. Standardization of HCl Solution by the Na2CO3 Method: 1. Weigh (to the nearest 0.1 mg), individual triplicate portions of about 0.15– to 0.20-g samples of pure Na2CO3 into 250-mL beakers or flasks. Record these weights in your journal and label the beakers or flasks to distinguish among these weights. 2. Pour about 75 mL of previously boiled distilled water (which has been cooled to room temperature) into the carbonate samples in each container, stir and warm gently (if necessary) until dissolution is complete. 3. Cool the sol’n. Add 2-3 drops of MR indicator to give a faint but distinct yellow color to the solution. 4. Fill the burets with the prepared HCl and NaOH solutions using the customary precautions. 5. Take and record the initial readings in your journal. 6. Titrate the carbonate samples with the acid until the solution just begins to change from yellow to a red-orange color. Stop the titration. Boil the solution for 1-2 min and cool to room temperature (Note 1). Continue the titration (Note 2) until a sharp yellow to an intermediate shade of orange color is observed. Record the final readings. 7. Calculate the net volume of HCl required and from this value calculate the molarity, M of the acid. Check that results agree to within 2 ppt before reporting the mean M of the acid. D. Analysis of the Unknown: 1. From a weighing bottle, weigh (to the nearest 0.1-mg) triplicate samples of about 0.20- to 0.25-g of the unknown soda ash, into beakers or flasks of about 250-mL capacity. Record these weights and label the beakers or flasks. 2. Repeat procedures C.# 2 – 6. Guide Question/s: 1. Explain the effects of: a) using unboiled distilled water in the preparation of titrant solutions. Distilled water contains very little dissolved substances that would interfere with a titration. b) boiling the titrated solution in the detection of the end-point using methyl red indicator. Standardization Sample 1 Final Volume (mL) Initial Volume (mL) Total (mL) HCl 32.50 0.80 31.7 NaOH 3.00 0.50 2.50 HCl 32.50 0.50 32.00 NaOH 4.30 3.00 1.30 HCl 35.30 0.20 35.1 NaOH 5.40 4.30 1.10 Sample 2 Final Volume (mL) Initial Volume (mL) Total (mL) Sample 3 Final Volume (mL) Initial Volume (mL) Total (mL) Analysis of Unknown Sample 1 2 3 Mass (g) 0.2277 0.2333 0.2431 Trial 1 Final Volume (mL) Initial Volume (mL) Total (mL) HCl 9.60 0.20 9.40 NaOH 5.00 0.50 4.50 HCl NaOH Trial 2 Final Volume (mL) Initial Volume (mL) Total (mL) 31.00 9.60 21.40 15.50 5.00 10.50 HCl 46.10 31.00 15.10 NaOH 31.20 15.50 15.70 Trial 3 Final Volume (mL) Initial Volume (mL) Total (mL) Calculations Standardization of HCl using 𝐿𝑖2 𝐶𝑂3 𝐿𝑖2 𝐶𝑂3 + 2HCl 2LiCl + 𝐶𝑂2 + 𝐻2 O Trial 1 Mass of 𝐿𝑖2 𝐶𝑂3 Total volume HCl used Total volume NaOH used 0.1692 g 31.70 mL 2.50 mL M 𝑚𝑜𝑙 𝐿 𝑚𝑜𝑙 𝐿 = = 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 2 𝑚𝑜𝑙 𝐻𝐶𝑙 )( ) 73.89 𝑔 𝐿𝑖2 𝐶𝑂3 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 1.15 𝑚𝐿 𝐻𝐶𝑙 1𝐿 − (2.50 𝑚𝐿 𝑁𝑎𝑂𝐻)]( 𝑚𝐿 𝑁𝑎𝑂𝐻 )(1000𝑚𝐿) 0.1692 g 𝐿𝑖2 𝐶𝑂3 ( [31.70 𝑚𝑙 𝐻𝐶𝑙 0.0046 2.6983 = 0.0017 mol/L Trial 2 Mass of 𝐿𝑖2 𝐶𝑂3 Total volume HCl used Total volume NaOH used 𝑚𝑜𝑙 𝐿 𝑚𝑜𝑙 𝐿 = 0.1777 g 3.50 mL 2.00 mL 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 2 𝑚𝑜𝑙 𝐻𝐶𝑙 0.1777 g 𝐿𝑖2𝐶𝑂3 ( )( ) 73.89 𝑔 𝐿𝑖2 𝐶𝑂3 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 𝑚𝐿 𝐻𝐶𝑙 1𝐿 3.50 𝑚𝑙 𝐻𝐶𝑙 − (2.00 𝑚𝐿 𝑁𝑎𝑂𝐻)(1.15 𝑚𝐿 𝑁𝑎𝑂𝐻 )(1000𝑚𝐿) 0.0048 = 3.4977 = 0.0014 mol/L Trial 3 Mass of 𝐿𝑖2 𝐶𝑂3 Total volume HCl used Total volume NaOH used 𝑚𝑜𝑙 𝐿 = 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 2 𝑚𝑜𝑙 𝐻𝐶𝑙 )( ) 73.89 𝑔 𝐿𝑖2 𝐶𝑂3 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 1.15 𝑚𝐿 𝐻𝐶𝑙 1𝐿 3.50 𝑚𝑙 𝐻𝐶𝑙 − (1.10 𝑚𝐿 𝑁𝑎𝑂𝐻)( 𝑚𝐿 𝑁𝑎𝑂𝐻 )(1000𝑚𝐿) 0.1764 g 𝐿𝑖2𝐶𝑂3 ( 0.1764 g 3.50 mL 1.10 mL 𝑚𝑜𝑙 𝐿 0.0048 = 3.4987 = 0.0014 PPT 1, 2 = = 𝑀1 − 𝑀2 𝑀1 + 𝑀2 2 x 1000 0.0017 mol/L − 0.0014 mol/L 0.0017 mol/L +0.0014 mol/L 2 x 1000 0.0003 𝑚𝑜𝑙/𝐿 = 0.0016 𝑚𝑜𝑙/𝐿 x 1000 = 187.50 PPT 1, 3 = = = 𝑀1 − 𝑀2 𝑀1 + 𝑀2 2 x 1000 0.0017 mol/L − 0.0014 mol/L 0.0017 mol/L +0.0014 mol/L 2 0.0003 𝑚𝑜𝑙/𝐿 0.0016 𝑚𝑜𝑙/𝐿 x 1000 x 1000 = 187.50 PPT 2, 3 = = 𝑀1 − 𝑀2 𝑀1 + 𝑀2 2 x 1000 0.0014 mol/L − 0.0014 mol/L 0.0014 mol/L +0.0014 mol/L 2 x 1000 0 = 0.0014 𝑚𝑜𝑙/𝐿 x 1000 =0 𝐻𝐶𝑙 𝑀𝑎𝑣𝑒𝑟𝑎𝑔𝑒 = 0.0014 + 0.0014 2 = 0.0014 Analysis of Unknown Sample 1 Mass of 𝑢𝑛𝑘𝑛𝑜𝑤𝑛 Total volume HCl used Total volume NaOH used % mass 𝐿𝑖2 𝐶𝑂3 = = 0.0281 0.2277 [9.40 mL HCl − (4.50 mL NaOH)( 0.2277 g 9.40 mL 4.50 mL 1.15 𝑚𝐿 𝑁𝑎𝑂𝐻 1𝐿 .18 𝑚𝑜𝑙 𝐻𝐶𝑙 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 73.89 𝑔 𝐿𝑖2 𝐶𝑂3 )]( )( )( )( ) 𝑚𝐿 𝐻𝐶𝑙 1000 𝑚𝐿 𝐿 2 𝑚𝑜𝑙 HCl 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 0.2277 g x 100 = 12.34 Sample 2 Mass of 𝑢𝑛𝑘𝑛𝑜𝑤𝑛 Total volume HCl used Total volume NaOH used 0.2333 21.40 mL 10.50 mL x 100 % mass 𝐿𝑖2 𝐶𝑂3 = [21.40 mL HCl − (10.50 mL NaOH)( 1.15 𝑚𝐿 𝑁𝑎𝑂𝐻 1𝐿 .18 𝑚𝑜𝑙 𝐻𝐶𝑙 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 73.89 𝑔 𝐿𝑖2 𝐶𝑂3 )]( )( )( )( ) 𝑚𝐿 𝐻𝐶𝑙 1000 𝑚𝐿 𝐿 2 𝑚𝑜𝑙 HCl 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 0.2333 g x 100 0.0620 = 0.2333 x 100 = 26.50 Sample 3 Mass of 𝑢𝑛𝑘𝑛𝑜𝑤𝑛 Total volume HCl used Total volume NaOH used % mass 𝐿𝑖2 𝐶𝑂3 = [15.10 mL HCl − (15.70 mL NaOH)( 0.2333 15.10 mL 15.70 mL 1.15 𝑚𝐿 𝑁𝑎𝑂𝐻 1𝐿 .18 𝑚𝑜𝑙 𝐻𝐶𝑙 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 73.89 𝑔 𝐿𝑖2 𝐶𝑂3 )]( )( )( )( ) 𝑚𝐿 𝐻𝐶𝑙 1000 𝑚𝐿 𝐿 2 𝑚𝑜𝑙 HCl 1 𝑚𝑜𝑙 𝐿𝑖2 𝐶𝑂3 0.2431g x 100 0.0196 = 0.2431 x 100 = 8.06 PPT 1, 2 = = %1 − %2 %1 + %2 2 12.34 % − 26.50 % 12.34 % + 26.50 % 2 x 1000 x 1000 14.19 % = 19.42 % x 1000 = 730.69 PPT 1, 3 = = %1 − %3 %1 + %3 2 12.34 % − 8.06 % x 1000 x 1000 12.34 % + 8.06 % 2 4.28 % = 10.85 % x 1000 = 39.45 % PPT 2, 3 = = %2 − %3 %2 + %3 2 26.50 % - 8.06 % 26.50 % + 8.06 % 2 x 1000 x 1000 18.44 = 17.28 x 1000 = 17280 average % = 12.34 % + 26.50% + 8.06% 3 = 15.63 % Conclusion: The technique of titrating is important in determining the concentration of an unknown. In the process of titration, a basic solution is gradually added to the acidic solution until complete neutralization is obtained. The end point of the titration is detected with the help of an indicator (methyl red in our case) as color of the solution changes upon neutralization. In general, in performing the acid base titration you need to record systematically and read precisely all the results to accurately determine the concentration of the unknown. My calculated result was only 15.63 % which is far behind with the accepted value of average percentage which ranges from 50-55% only. Thus, I failed to determine the concentration of my unknown sample, but I gained substantial knowledge that will help me on what to do and what to avoid for next laboratory experiment.