# 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. Set-up a titration stand

2. Perform a fast titration and a slow titration

3. Monitor the pH of a titration reaction

# 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 a solution of known concentration and is slowly added to the analyte. The analyte has a known volume but its concentration is unknown.

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. Which is the titrant? Which is the analyte?

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

# Experimental Design

Today, we will determine the concentration of a solution of sodium hydroxide. In order to accomplish this, we will titrate the sodium hydroxide solution with standardized solution of HCl of known concentration.

## Titrations as a technique

Today you will be performing two types of titrations:

1. FAST

2. SLOW

In a fast titration, the titrant is released quickly into the analyte, 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 analyte. 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.

# Reagents and Supplies

Approximate 1.0 M aqueous sodium hydroxide, standardized aqueous hydrochloric acid, and a phenolphthalein solution.

Burette, pH paper, materials from your drawer

(See posted SDS, Safety Data Sheets)

# Procedure

## Analyte

1. Add 25 mL of hydrochloric acid solution into a 125 mL Erlenmeyer flask.

2. Add one drop of phenolphthalein solution into the 125 mL Erlenmeyer flask.

## Titrant

1. Calculate the volume of 1.0 M aqueous sodium hydroxide solution needed to prepare 250-mL of an approximately 0.10 M aqueous sodium hydroxide solution. Check your calculations with the lab instructor before proceeding to the next step.
2. 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.

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

4. Clamp a burette to the burette stand.

5. Condition the burette by filling with water. Allow the water to drain. Water can be discarded in sink. Then fill burette to mark with NaOH solution.

6. Record the initial burette reading. For some burettes, the fill line is marked “zero”.

## Using universal pH indicator paper to monitor changes in pH

1. Grab 5 strips of pH paper from the back cabinet. While performing the first titration, stop after adding five specific volumes of NaOH to the flask and measure the pH of the reaction mixture (Erlenmeyer flask). Use these values to fill out the chart in your worksheet.

## Fast titration (perform once)

1. Titrate NaOH into the HCl 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)

1. Set up another fresh HCl solution in a clean Erlenmeyer flask with indicator (see steps 1 and 2).

2. 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.

3. 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.

4. Repeat steps 11 – 13 until you have three successful slow titrations.

5. Dispose all waste into appropriate waste disposal containers as instructed by your instructor.

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

## Determine the concentration of NaOH

Using the data from the slow titrations, determine the concentration of NaOH in units of mol/L for the individual titrations, and their average. Calculate the standard deviation. Use this data to fill in your Workhseet.

Lab 6: Acid/base titration is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.