Lab 6: Acid/base titration

Experiment 6: My First Acid - Base Titration

Learning Outcomes

Upon completion of this lab, the student will be able to:

1. Start preparing for emergencies in the chemical lab

2. Set-up a titration stand

3. Prepare primary standard solutions of Potassium Hydrogen Phthalate (KHP) solutions

4. Standardize a solution of Sodium Hydroxide (NaOH)

5. Perform a fast titration and a slow titration

 

Reagents and Supplies

Approximate 1.0 M aqueous sodium hydroxide (NaOH), potassium hydrogen phthalate (KHP), and a phenolphthalein solution.

Burette, materials from your drawer

(See posted SDS, Safety Data Sheets)

 

Introduction

Titration is an analytical quantitative technique used to determine the concentration of a solute; a pH-titration is used to determine the concentration of an acid or a base. Titrations play an important role in determining amount and purity in many manufacturing processes. These include food processing, textile, wood product manufacturing, petroleum, pharmaceuticals and chemical manufacturing. Titrations play a critical role in the manufacturing of biodiesel, water purification and waste water treatment and even in the production of various products in the dairy industry.

Here are some examples from pharmaceutical manufacturing. The active ingredient in many cough syrups, ephedrine, is tested for purity through titration with perchloric acid. Some drugs, such as the antifungal, clotrimazole, are end products of titration reactions. Titrations are even used to determine the amount of non-pharmacologically active ingredients, such as the binding substances that make up medicine tablets.

A titration involves two solutions: the titrant and the analyte. Typically, the titrant is slowly added from a burette to the flask containing the analyte until you reach the endpoint. At the endpoint, you know the initial volume of analyte (how much you added to our flask) and the volume of added titrant (by reading off the burette). The titrant is a solution of known concentration and you figure out the concentration of the analyte.

Consider the reaction between aqueous solutions of sulfuric acid and potassium hydroxide. The balanced chemical equation for the reaction is shown below:

\[\ce{H2SO4(aq) + 2 KOH(aq) -> K2SO4(aq) + 2 H2O(l)}\tag{1}\]

Assume that the molar concentration of \(\ce{KOH}\) is 0.100 M and that of \(\ce{H2SO4}\) is unknown. 

For this example, \(\ce{KOH}\) was slowly titrated to 10.00 mL of \(\ce{H2SO4}\). The volume of 13.75 mL of \(\ce{KOH}\) was needed to completely react with the \(\ce{H2SO4}\). The concentration of the sulfuric acid (in the flask, at the beginning of the titration) is calculated as follows:

For an interactive version with more explanation, follow this link or the QR code above.

Two takeaways from this calculation:

1. As seen from the above calculation, the stoichiometric ratio between the two reactants is the key to the determination of the concentration of the unknown solution. (You need a balanced equation).

2. Determining the unknown concentration requires meticulous determination of the volume of the titrant.

In order to conduct the above experiment, typically the \(\ce{H2SO4}\) is in an Erlenmeyer flask, and the \(\ce{KOH}\) is in a burette. The \(\ce{KOH}\) is added one drop at a time from the burette into the acid solution with constant stirring to ensure that the reagents combine and react.

The equivalence point of an acid/base reaction

The equivalence point is defined as that point in the titration when stoichiometrically equal amounts of acid and base are present. In the \(\ce{H2SO4}\)/\(\ce{KOH}\) example shown previously (eq 1), that would be when two moles of \(\ce{KOH}\) have been added to one mole of \(\ce{H2SO4}\). Therefore, the equivalence point depends on the reaction stoichiometry.

At the beginning of the titration, the solution in the Erlenmeyer flask is acidic. As the base is added, it completely reacts with the acid and the solution in the Erlenmeyer flask continues to be acidic. But, at the equivalence point, the acid has completely reacted with the base. If even one tiny drop of base is added beyond that needed to arrive at the equivalence point, the solution in the Erlenmeyer flask is basic. This difference in the acid/base property of the solution in the Erlenmeyer flask is used to visually determine the end of the titration.

Using color to visualize the equivalence point

An indicator is a chemical substance whose color depends on the acid/base property of the medium it is present in. Phenolphthalein is an indicator, which is colorless in an acidic medium and has a pink color in a basic medium.

In this titration, a few drops of phenolphthalein should be added to the acid in the Erlenmeyer flask. The solution will remain colorless until the equivalence point.

When the equivalence point has been crossed and the solution becomes basic, the phenolphthalein will take on a pink color. This is the reason to add the base drop by drop, so that even though the equivalence point will be crossed, the titration can be stopped at the appearance of the first permanent pale pink color.

This point in the titration when the indicator changes color is referred to as the End Point. Note, the End Point of a titration is slightly beyond the Equivalence Point. In the case of the phenolphthalein, the intensity of the pink color increases as the solution becomes more and more basic. Therefore, it is important to stop the titration at the appearance of a permanent pale pink color. In order to easily observe the color changes in the solution, it is a good idea to place a sheet of plain white paper beneath the flask.

My first titration

Today we will be learning the basics of acid-base titrations. We will be going over the following techniques:

How to

• Assemble a burette stand

• Properly condition the burette for use

• Accurately measure the volume of the burette

• Dispense titrant using a ‘stop-cock’

• Determine the end point of the titration

• Use and dispose of waste

• Analyze titration data

Today, we will determine the exact concentration of a solution of sodium hydroxide. In order to accomplish this, we will titrate the sodium hydroxide (NaOH) solution with standardized solution of KHP of known concentration.

Sodium hydroxide is an extremely hygroscopic solid. When solid sodium hydroxide is weighed using an electronic balance it tends to absorb moisture from the atmosphere; this leads to inaccuracies in the mass of the sodium hydroxide and thereby inaccuracies in its molarity calculations.

In order to circumvent the above problem, a sodium hydroxide solution of approximate concentration is prepared first. This solution is titrated with a primary acid standard. The concentration of this acid is known with greater accuracy. The data from this titration is then used to calculate a more accurate value for the concentration of the sodium hydroxide solution. A commonly used primary standard for titration with sodium hydroxide solution is the weak acid potassium hydrogen phthalate or KHP (\(\ce{C8H5O4K}\)).

The reaction between KHP and NaOH is given below:

\[\ce{HC8H4O4K + NaOH -> C8H4O4KNa + H2O}\]

____ mole(s) of NaOH must be added for every _____ mole(s) of \(\ce{C8H5O4K}\) in solution

The titration of NaOH with KHP involves adding NaOH from the burette to a known mass of KHP in solution. The data from the titration is then used to calculate the molarity of the NaOH.

Titrations as a technique

Today you will be performing two types of titrations:

1. FAST

2. SLOW

In a fast titration, the titrant (solution in the burette) is released quickly into the solution in the flask below, hence fast. The goal of this titration is to determine the approximate volume of titrant needed to induce the change of color (determine the end point). This titration is not quantitative; it will not give an accurate determination of the unknown concentration.

In contrast, a slow titration will give you an accurate determination of the unknown concentration; however, the titrant is released very slowly into the solution in the flask below. This is especially true as you near the end point. Remember, it is important to stop the titration at the appearance of a permanent pale pink color. This slow, methodical process will yield an accurate measurement of volume of titrant.

 

 

Procedure

Record all your calculations, measurements, and observations into your notebook, and then fill in your Summary Sheet.

Part 1:  Safety Video on Preparing for Emergencies

Please watch video #6 in the ACS lab safety series. 

As you are watching, think of which emergencies could have happened in previous sessions already (with the materials and substances we have been using), and how we prepared for them. When you are done, answer the questions on the handout. Today's lab involves acids and bases. Listen carefully as your instructor explains the hazards associated with them.

Part 2:  Standardization of the Sodium Hydroxide (NaOH) Solution 

Preparing KHP Solution (in Flask)

1. Calculate the mass of KHP needed to prepare 25.00 mL of 0.100 M KHP and write this calculation in your notebook.  Check your calculation with your instructor before moving to Step 2.

2. Weigh an amount of KHP as close as possible to the calculated mass in step 1 and record the exact mass of KHP measured.

3. Transfer the solid KHP into a 125-mL Erlenmeyer flask.

4. Add 25 mL of deionized water into the 125-mL Erlenmeyer flask containing the solid KHP.  

5. Add one drop of phenolphthalein solution into the Erlenmeyer flask and swirl the flask to mix.

Preparing NaOH Solution (in Burette)

6. Calculate the volume of approximately 1.0 M aqueous sodium hydroxide solution needed to prepare 250-mL of an approximately 0.10 M aqueous sodium hydroxide solution.  Write your calculations in your lab notebook.  Check your calculations with the lab instructor before proceeding to the next step.

7. Measure the volume of 1.0M aqueous sodium hydroxide solution calculated in step 3 using the 25-mL graduated cylinder located in the fume hood. Pour this volume into your Florence flask and dilute the solution with enough deionized water to make a total volume of 250-mL.  Remember this 0.1M is only an approximate concentration.

8. Label the flask with your name, date, lab section, and 0.1M aqueous sodium hydroxide (do not use chemical symbols) with a Sharpie. 

9. Clamp a burette to the burette stand.

10. Condition the burette by filling with water. Allow the water to drain. Water can be discarded in sink. Rinse the burette with a few mL of NaOH solution and then discard this into the waste container. Then fill the burette slightly above the "zero" mark with NaOH solution.

11. Drain some of the NaOH solution so that there are no air bubbles in the tip of the burette AND so that the level of NaOH falls between 0.00mL and 1.00mL.  Record the initial burette reading, which will NOT be at 0.00mL.   

Fast titration (perform once)

12. Titrate NaOH into the KHP solution until a faint but permanent pale pink color is obtained. Make sure to swirl the Erlenmeyer flask thoroughly to ensure mixing of the reagents. In case any NaOH solution falls on the side of the Erlenmeyer flask, rinse the sides of the flask with deionized water. Record the volume of NaOH used to reach the endpoint.

Slow titration (perform three times)

13. Set up another fresh KHP solution in a clean Erlenmeyer flask with indicator (see steps 1-5).

14. Titrate with NaOH slowly. If the level of NaOH is close to 50 mL (i.e. bottom of the scale), write down the value, fill the burette with more diluted NaOH, and write down the new value. This allows you to calculate the total amount dispensed from the burette.  Don't let the level of the NaOH go below the 50.00mL line on the burette or you won't be able to determine the volume dispensed and will need to start the titration over.

15. Remember to slow down to a drip at least 1-2 mL before the recorded volume you observed in the fast titration. Make sure to swirl the Erlenmeyer flask thoroughly to ensure mixing of the reagents. Stop when a sustained faint pink color is observed.

16. Repeat steps 13 – 15 to conduct three total slow titrations.

 

Part 3:  Calculate the concentration of NaOH

17. Using the average data from the slow titrations, calculate the concentration of NaOH in units of mol/L.

 

18. Cover the remaining NaOH solution with parafilm and store your labeled Florence flask on the side counter.

19. Check with your instructor before cleaning up.  Dispose all waste into appropriate waste disposal containers as instructed by your instructor.  Clean and return all glassware.  Place the burette upside down on the stand and turn the stop-cock to the open position to allow it to drain and dry.