Experiment 3 Potentiometric Titration Using a Electrode Introduction Titrations are most commonly performed either to find out how much analyte is present or to measure equilibrium constants of the analyte. We can obtain the information necessary for both purposes by monitoring the of the solution as the titration is performed. In this experiment you will use a electrode to follow the course of an acid-base titration. You will observe how changes slowly during most of the reaction and rapidly near the equivalence point. You will compute the first derivatives of the titration curve to locate the end point. From the mass of unknown acid or base and the moles of titrant, you can calculate the molecular mass of the unknown. Reagents 1. Standardized 0.1 M HCl 2. Standardized 0.1 M NaOH 3. Unknown Procedure 1. Your instructor will recommend an appropriate mass of unknown (5~8 mmol) for you to weigh accurately and dissolve in distilled water in a 100-mL volumetric flask. Dilute to the mark and mix well. 15
2. Make a rough test of the of your unknown solution to ascertain whether it is an acid or a base. If your unknown is a base, follow Step 3 below. Otherwise, follow Step 4 below. 3. Preparation and standardization of 0.1 M HCl follow Procedure A and B of Experiment 2. 4. Preparation and standardization of 0.1 M NaOH follow Procedure C and D of Experiment 2. 5. The first titration is intended to be rough, so that you will know the approximate end point in the next titration. For the rough titration, pipet 20.0 ml of unknown into a 250-mL Erlenmeyer flask. If you are titrating an unknown base, add 3 drops of bromocresol green indicator and titrate with standard 0.1 M HCl to the green end point, using a 50-mL buret. If you are titrating an unknown acid, add 2 drops of phenolphthalein indicator and titrate with standard 0.1 M NaOH to the pink end point. 6. Following instructions for your particular meter, calibrate a meter and glass electrode, using buffers with values near 7 and 4. Rinse the electrodes well with distilled water and blot them dry with a tissue before immersing in any new solution. 7. Now comes the careful titration, pipet 40.0 ml of unknown solution into a 100-mL beaker containing a magnetic stirring bar. Position the electrode in the liquid so that the stirring bar will not strike the electrode. If a combination electrode is used, the small hole near the bottom on the side must be immersed in 16
the solution. This hole is the salt bridge to the reference electrode. Allow the electrode to equilibrate for 1 min with stirring and record the. 8. Add 1~2 drop of indicator and begin the titration. The equivalence volume will be about two times greater than it was in Step 5 (2V 5 ). Add ~2.0 ml aliquots of titrant and record the exact volume, the, and the 30 s after each addition. When you are within 1.0 ml of the indicator end point ( 2V 5-1), add titrant in 0.1 ml increments. Continue with 0.1 ml increments until you are 1.0 ml past the indicator end point (2V 5 +1). The equivalence point has the most rapid change in. Then, add five more 2.0 ml aliquots of titrant and record the after each. Data analysis 1. Construct a graph of versus titrant volume. Make on your graph where the indicator change was observed. 2. Compute the first derivative (the slope, Δ/ΔV) for each data point within ± 0.5 ml of the indicator end point volume (V i of step 8). From your graph, estimate the equivalence volume (V e ) as accurately as you can, as shown below. 17
3. Compare the indicator end point volume (2V 5 ) to the equivalence volume (V e ) estimated from the first derivative. 4. From the equivalence volume (V e ) and the mass of unknown, calculate the molecular mass of the unknown. 5. From the value at halfway the equivalence volume (1/2V e ), find the approximate pk a value of the unknown acid (or the conjugate acid of the unknown base). References 1. Skoog, West, Holler and Crouch, Fundamentals of Analytical Chemistry, 8 th ed., Chap. 14 & 15 2. D.C. Harris, Quantitative Chemical Analysis, 7 th ed., Chap. 11 18
Data 1. Sampling Unknown No. Weight of sample g 3. Standardization of HCl against Na 2 CO 3 Titration I II Weight of Na 2 CO 3 g g Initial buret reading ml ml Final buret reading ml ml Volume of HCl ml ml Volume of HCl used in blank titration ml ml Volume of HCl used in this titration ml ml Molarity of HCl M M Average Molarity of HCl M Calculation Weight of Na2CO3 Molar mass of Na2CO 3 M HCl V ml, HCl 1 2000 19
4. Standardization of NaOH against potassium hydrogen phthalate (KHP) Titration I II Weight of KHP g g Initial buret reading ml ml Final buret reading ml ml Volume of NaOH ml ml Volume of NaOH used in blank titration ml ml Volume of NaOH used in this titration ml ml Molarity of NaOH M M Average Molarity of NaOH M Calculation Weight of KHP M Molar mass of KHP NaOH V ml, NaOH 1 1000 20
5. Rough titration Initial buret reading Final buret reading Volume of NaOH or HCl (V 5 ) ml ml ml 8. Careful titration Estimated volume of end point for 40 ml unknown sol n (2V 5 ) ml 0 2.00 4.00 6.00 volume of indicator end point (V i of step 8) ml 21
Data analysis 1. Construct a graph of (y-axis) versus titrant volume(x-axis). Make on your graph where the indicator change was observed. 22
2. Compute the first derivative (the slope, Δ/ΔV) for each data point within ± 0.5 ml of the indicator end point volume (V i of step 8). From your graph, estimate the equivalence volume (V e ) as accurately as you can. / V (1/mL) V avg (ml) Graph of / V versus V avg : 23
3. Compare the indicator end point volume (2V 5 ) to the equivalence volume (V e ) estimated from the first derivative. 4. From the equivalence volume (V e ) and the mass of unknown, calculate the molecular mass of the unknown. 5. From the value at halfway the equivalence volume (1/2V e ), find the approximate pk a value of the unknown acid (or the conjugate acid of the unknown base). Discussion 24