pH Titration Curve Equivalence Point

Acid and Base Organic Chemistry Example Problems

I have most of the answers on the attachment. I just looking to see if I am on the right track. I have added everything including back ground info that was posted on my lab assignment. I have also included all of my lab work which is completed and assumed correct on an attachment below.

1. What is the equivalence volume of the second titration? That is, how many mL of NaOH were added to the solution to cause the solution to turn pink? Note that you must calculate this volume; it is not the burette reading!

2. What is the pH at the equivalence point?

3. Calculate the molarity of the HCl from the volumes of acid and base at the equivalence point and the molarity of the NaOH. Use the equation given in the Background section.

4. Generate a titration curve by plotting the pH on the y-axis and the number of mL of base added to the flask on the x-axis. How do you determine the equivalence volume of the titration from this curve? What is the equivalence volume according to the titration curve? Calculate the molarity of the acid using this equivalence volume and the equation provided in the Background section of this laboratory. Why use a titration curve when the raw data are available in your data table?

5. How is the number of moles of NaOH affected at the equivalence point of the solution? Explain.

6. Compare the molarity of the acid determined by different individuals in the class. Are all these values same? Why, or why not?

Back Ground and procedures:

Titration is one of the most useful techniques for determining the concentration of an unknown solution, and it is often used with acidic and basic solutions. An apparatus called a burette is used to add very small controlled amounts of substance to a solution to the point at which the characteristics of the solution reach a specific point of change. The end point of titration is reached when a stoichiometric number of moles of an acid or base are added to the unknown solution to neutralize it.

PLEASE NOTE: The technique of titrating an acid with a base involves several steps in order to obtain exact results. Thus it is important to follow precisely the laboratory procedure.

The determination of the end point of titration can be detected by several methods:

- indicator - a substance is added to the solution that changes color when the end point of titration is reached.

- pH meter - a potentiometric change in the solution as indicated by a pH meter as the solution transitions from acidic to basic or vice versa.

- spectrophotometer - a change in the amount of light absorbed by the solution as indicated by a spectrophotometer as the solution transitions in relative concentrations of substances.

- conductivity meter - a change in the conductivity of electricity of the solution as it transitions in relative concentrations of substances.

In this lab an indicator, phenolphthalein, will be used to determine the end point of titration to determine the concentration of a substance in a solution. Phenolphthalein causes the solution color to change from clear to pink at the endpoint of titration.

A graph called a "titration curve" is used to determine the precise endpoint of titration. The titration curve is generated by plotting the the pH of the solution vs. the amount of base (in mL) added to the solution. The number of mL of base added to an acid in order to reach the endpoint or equivalence point is called the equivalence volume.

When titrating a monoprotic acid, such as HCl, with sodium hydroxide (NaOH), the molar ratio for the reaction is 1:1 - that is, one mole of NaOH will neutralize one mole of HCl - and all of the following are true at the equivalence point:

(a) Moles of acid in flask = moles of base added from the buret

(b) (molarity of acid) x (volume of acid) = (molarity of the base) x (volume of added base)

Equation (b) takes the familiar form M1*V1 = M2*V2, from which the molarity of the acid, M2, is determined.

1. Take a clean Erlenmeyer flask from the Glassware shelf and place it on the workbench.

2. Add 25 mL HCl (of unknown concentration) to the flask.

3. Add 2 drops of phenolphthalein, an indicator, to the flask.

4.Take a burette from the Glassware shelf and place it on the workbench.

5. Fill the burette with 50 mL of 0.5M NaOH solution. Record this initial volume.

6. Drag the flask to the lower half of the burette such that the burette can deliver NaOH to the flask. Make sure the burette and the flask are connected.

7. Open the Data window and click on the flask. Click the pushpin icon in the Data window to lock its display to the flask.

8. Take a pH meter from the Tools shelf and set it on the flask. Record the initial pH of the solution.

9. Open the Properties window and click on the burette. Enter "1" in the amount window to add NaOH to the flask in 1 mL increments. One mL of NaOH will be delivered to the flask each time the stopcock button is clicked.

10. Set up a data table on a piece of paper with the following column headings:
mL of base (burette reading)
pH
moles of NaOH

11. Click the stopcock button on the Properties window to deliver 1 mL of NaOH to the flask. Continue to add NaOH in 1 mL increments. Record the mL of base (burette reading), the pH of the solution in the flask, and the moles of NaOH in the flask for each mL of base added. Note that the initial mL of base is 50 mL, which is equal to zero mL of base added to the flask.

12. The solution in the flask will change in color from clear to pink due to the phenolphthalein indicator when the endpoint is either reached or crossed. Mark the point in the data table at which the color change occurs.

13. Add 5 more 1 mL increments of NaOH to the flask recording the mL of base, the pH, and the moles of NaOH at each increment.

14. Detach the pH meter and drag the burette and flask to the recycling chute.

15. Take a new flask from the Glassware shelf and place it on the workbench.

16. Add 25 mL of HCl and 2 drops of phenolphthalein to the flask.

17. Set a new burette on the workbench and fill it with 50 mL of NaOH.

18. Place the flask under the burette and attach the pH meter.

19. Based on the results of the previous titration, you may initially add a relatively large amount of NaOH solution all at once to the flask to get to 1 mL BEFORE the endpoint. For example, if 20 mL of NaOH caused a color change in your first titration, you know experimentally that you can add 19 mL of NaOH to the second titration without causing a color change. You can enter 19 in the Properties window and click the stopcock button to deliver 19 mL of NaOH to the flask at once.

20. AFTER the initial delivery of NaOH to the flask, change the amount of NaOH to be delivered to 1 in the Properties window. For this titration, use the DROP-WISE button to deliver very small (0.05 mL) of NaOH per click of the drop-wise button. In this way, you are measuring the endpoint of titration more precisely. Record the mL of base (burette reading), the pH, and the moles of NaOH in the flask with each click of the drop-wise button until the endpoint is reached. Mark this endpoint on your data sheet. Add 5 more drop-wise increments of NaOH noting the mL of base, pH, and moles of NaOH.

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