# 2: Standardization of Acids and Bases (Experiment)


### Introduction

This experiment affects several experiments that follow. You should strive for maximum precision in this experiment since inaccuracies here will result in inaccuracies later.

Safety Precautions

Be especially careful when using the strong acid or strong base solutions as they can cause severe burns. Since concentrated HCl has a pungent odor it would be wise to dispense this solution in the hood. All waste solutions may be disposed of by rinsing them down the drain.

## Part I. Preliminary Steps

1. Obtain approximately 4 grams of potassium acid phthalate (KHP), which is in the laboratory. Transfer the contents of the vial to a clean, dry weighing bottle. Using a folded paper strip or crucible tongs put the weighing bottle into a 50 mL beaker that has your initials on it. With tongs, lay the stopper crosswise on the mouth of the weighing bottle and then put the beaker and its contents into a 110°C oven for at least one hour.
2. Obtain approximately 4 grams of unknown KHP from your TA. Record the unknown number in your lab notebook. Transfer the contents of the vial to a clean, dry weighing bottle. Using a folded paper strip or crucible tongs put the weighing bottle into a 50 mL beaker that has your initials on it. With tongs, lay the stopper crosswise on the mouth of the weighing bottle and then put the beaker and its contents into a 110°C oven for at least one hour.
3. Obtain approximately one gram of pure sodium carbonate, which is in the laboratory. Place the solid in a clean, dry weighing bottle. Using a folded paper strip or crucible tongs put the weighing bottle into a 50 mL beaker that has your initials on it. With tongs lay the stopper crosswise on the mouth of the weighing bottle and then put the beaker and its contents into a 110°C oven for at least one hour.

Question $$\PageIndex{1}$$

Why do we dry these chemicals before we weigh them? (Please do these questions as they appear and before you do the next step in the procedure. They are designed to help ensure you understand the procedures and avoid costly errors.)

1. Thoroughly wash two 1-liter glass bottles and their caps. Rinse everything well with deionized water. Dry the outside of each bottle but not the inside. Label one bottle HCl, the other NaOH.

## Part II. Preparation of Solutions

1. An approximately 0.1 M NaOH solution will be prepared by dilution of a 50 weight-percent NaOH solution. This solution has a density of 1.53 g/mL. Compute the volume of this solution that is required to prepare 1000 mL of 0.1 M NaOH. If you are unsure of your answer confirm it with your TA. Obtain 1000 mL volumetric flask, add around 500 ml (halfway) of ultra pure water from the proper carboy in the laboratory. Using the 50% by mass stock solution, measure the calculated volume of NaOH solution into a clean, dry 10 mL graduated cylinder. Pour the contents of the graduated cylinder into your flask. Fill the graduated cylinder with five successive 10 mL portions of ultra pure water, each time emptying it into the flask. Swirl then add more ultra pure water until the final volume of the system is 1000 mL. Pour the contents into the 1 L storage bottle and cap the bottle. Label the bottle with chemical name and concentration. Allow the solution to return to room temperature. Sodium hydroxide solutions should never be left uncapped. Base solutions will absorb carbon dioxide from the atmosphere, thereby changing the titer of the solution.

Question $$\PageIndex{2}$$

Give the balanced chemical equation for the formation of carbonic acid and for the reaction of carbonic acid with sodium hydroxide.

1. An approximately 0.1 M HCl solution will be prepared by dilution of concentrated HCl, which is about 12 M. Compute the volume of this solution that is required to prepare 1000 mL of 0.1 M HCl. If you are unsure of your answer confirm it with your TA. DO NOT REMOVE THE CONCENTRATED HCl BOTTLE FROM THE FUME HOOD! Obtain 1000 mL volumetric flask, add around 500ml (halfway) of ultra pure water from the proper carboy in the laboratory. Measure the calculated volume of concentrated HCl solution into a clean, dry 10 mL graduated cylinder. Pour the contents of the graduated cylinder into your flask. Always remember to add acid to water and never the reverse. Swirl then add more ultra pure water until the final volume of the system is 1000 mL. Pour the all the contents into the 1L storage bottle and cap the bottle. Label the bottle with chemical name and concentration. Allow the solution to return to room temperature.

Question $$\PageIndex{3}$$

Why do we always add acid to water and never the reverse?

## Part III. Standardization of the NaOH Solution

1. After the KHP has dried for at least 1 hour, remove the beaker from the oven using tongs, and set it on a pad of paper on the bench top. Allow it to cool there until you can put the inside of your forearm against the beaker without noticeable discomfort. Then put the beaker into a desiccator and allow it to cool for at least another half-hour. The weighing bottle should remain uncapped throughout this operation.
2. By difference, weigh (to the nearest 0.1 mg) triplicate 0.7000-0.9000 gram samples of KHP into three separate 250 or 300 mL Erlenmeyer flasks. Since the absolute accuracy of the balances can vary slightly, it is important that you use the same balance throughout an experiment. In this way, any systematic error in the balance is removed.

To do such a weighing, remove the beaker from the desiccator and, using tongs, seat the top of the weighing bottle. Immediately replace the desiccator lid to avoid breakage. With a folded paper strip transfer the stoppered bottle to the balance pan and weigh. Remove the bottle with the paper strip, remove the stopper with tongs, tap out some of the contents into the Erlenmeyer flask, replace the stopper, and re-weigh the bottle. Continue this operation until the desired sample size has been obtained. Weighing the second and third samples will be faster because you can match the pile size of the first sample.

1. To each of the Erlenmeyer flasks, add 50-75 mL of deionized water and 2 drops of phenolphthalein indicator solution. Against a white background, titrate each flask with your NaOH solution to the first appearance of a faint pink color that persists for 30 seconds. You will have to learn to "split" drops to get a very faint pink endpoint. In addition, you should use a black-lined background card to accurately read the meniscus. Wash down the sides of the flask with deionized water using your water bottle to be sure all the reactants are in the solution. Given that the molecular weight of KHP is 204.23 and that it is a monoprotic acid, calculate the molarity of NaOH for each of your samples. Your triplicate molarities should agree to within three parts per thousand. Calculate the average, standard deviation, 95% confidence limit, and relative deviation. This is easily done on a computer spreadsheet.

Question $$\PageIndex{4}$$

Why does it not matter how much water you add when dissolving the acid or when carrying out the titration?

1. Carefully cap and store the solution for use in later experiments.
2. Obtain approximately 4 grams of unknown soda ash from your TA. Record the unknown number in your lab notebook. Transfer the contents of the vial to a clean, dry weighing bottle. Using a folded paper strip or crucible tongs put the weighing bottle into a 50 mL beaker that has your initials on it. With tongs, lay the stopper crosswise on the mouth of the weighing bottle and then put the beaker and its contents into a 110°C oven for at least one hour.

## Part IV. Unknown KHP Analysis

You will be given an unknown sample of KHP (KHC8H4O4). Analyze the sample in the same way that you standardized your NaOH solution. A brief outline of the procedure is described below.

1. Weigh three samples of your dry, cooled unknown KHP within the mass range of 0.9000 - 1.1000 gram.
2. Dissolve in deionized water and titrate with standard NaOH.
3. Calculate the mass percent of KHP in each sample. Report an average, standard deviation, 95% confidence limit, and relative deviation.

Question $$\PageIndex{5}$$

Write a balanced chemical equation for this reaction.

## Part V. Standardization of HCl Solution

This is a more complicated procedure than the one you have just accomplished. The reason for this is that there are very few primary base standards. One of the best is sodium carbonate. However, the use of this standard is complicated by the fact that carbonate is a weak acid and that a final product of the titration is a gas. The two acid/base reactions are:

$\ce{H^{+} + CO3^{2-} <=> HCO3^{-}}$

$\ce{H^{+} + HCO3^{-} <=> H2CO3 -> H2O + CO2(g)}$

These two reactions combine to from a solution which is called a buffer and which is described in Chapter 10. We will boil the solution near the end of the titration to drive out the carbon dioxide and thus destroy this buffer and get a much sharper endpoint.

### Procedure

1. After the $$\ce{Na2CO3}$$ has dried for at least 1 hour, remove the beaker from the oven using tongs, and set it on a pad of paper on the bench top. Allow it to cool there until you can put the inside of your forearm against the beaker without discomfort. Then put the beaker into a desiccator and allow it to cool for at least another half-hour. The weighing bottle should remain uncapped throughout this operation.
2. By difference, weigh (to the nearest 0.1 mg) triplicate 0.2000-0.2500 gram samples of Na2CO3 into three separate 250 or 300 mL Erlenmeyer flasks. Since the absolute accuracy of the balances can vary slightly, it is important that you use the same balance throughout an experiment. This way, any systematic error in the balance is removed.
3. Dissolve the first sample of sodium carbonate in 50 mL of deionized water and add 20 drops of bromocresol green indicator. Titrate the solution with HCl until the solution just changes from blue to green. Wash down the sides of the flask with deionized water using your water bottle to be sure all the reactants are in the solution. Boil the solution for 2 to 3 minutes, during this time the color should revert back to blue. If it does not, then a measured amount of base must be added to change the color to blue and this volume of added base must be included in the calculations. Cool to room temperature, and complete the titration to the final green endpoint.

Question $$\PageIndex{6}$$

What is the reason for boiling the solution? Why is it important?

1. Compute the HCl molarity. Report an average, standard deviation, 95% confidence limit, and relative deviation.

Question $$\PageIndex{7}$$

Write balanced equations for the reactions involved in this standardization.

1. Carefully store and cap the solution for use in later experiments.

## Part VI. Unknown Soda Ash Analysis

Obtain an unknown sample of soda ash from the stockroom. Soda ash is a mixture of sodium carbonate and an unreactive substance. Analyze the sample in the same way that you standardized your HCl solution. A brief outline of the procedure is described below.

1. Weigh three samples within the mass range of 0.5600 to 0.6000 gram.
2. Dissolve in deionized water and titrate with standard HCl.
3. Calculate the mass percent of sodium carbonate in each sample. Report an average, standard deviation, 95% confidence limit, and relative deviation.

Clean-up

After the experiment is completed, drain any remaining solution from the burette. Rinse each burette with deionized water. Then, fill the burette with deionized water and carefully place it in the burette rack.