Acid and Base Concentrations

Source: Smaa Koraym at Johns Hopkins University, MD, USA

1. Preparation of ~0.1 M NaOHExpand

   In the first part of the lab, you will use a 50% w/w solution of NaOH to prepare 500 mL of ~0.1 M. The 50% w/w NaOH is indicative of its weight ratio. For example, if the instructor prepared 150 mL of the 50% w/w NaOH solution, then 150 g of NaOH was dissolved in 150 g of water, and the total weight of the solution is 300 g.

   • To begin, put on the appropriate personal protective equipment, including gloves, chemical splash goggles, and a lab coat.
   • Calculate the molarity of 50% w/w NaOH, which is M₁ in the dilution formula. Note: The density of 50% w/w NaOH is 1.53 g/mL, and the molar mass of NaOH is 39.998 g/mole.

     Table 1. Preparation of 0.1 M NaOH from 50% w/w NaOH

     Density of 50% w/w stock solution	1.53 g/mL
     Molar massNaOH	39.998 g/mol
     Mass of NaOH in 50% w/w stock solution (mg)
     Total mass of 50% w/w stock solution
     Volume of 50% w/w stock solution (mL)
     Moles of NaOH in 50% w/w solution (mol)
     Molarity of 50% w/w stock solution (M1)
     Volume of 50% w/w solution needed (V1)

     Click Here to download Table 1
   • Use the dilution formula to determine the volume of 50% w/w NaOH needed to prepare 500 mL of ~0.1 M NaOH.
   • Label the 500 mL polyethylene bottle as ‘~0.1 M NaOH’.
   • Adjust the volume on a 1-mL pipette to the value calculated and use it to transfer the 50% w/w NaOH to the polyethylene bottle.
   • Use a 100-mL graduated cylinder to measure the amount of water and pour it into the bottle containing the NaOH. Cap the bottle and invert it several times to mix the solution. Note: You will need 500 mL minus the calculated volume of NaOH.
2. Standardization of 0.1 M NaOHExpand

   After you have prepared 0.1 M NaOH, determine its exact concentration, or standardize it, using the acid-base titration method. In this technique, a base like NaOH is slowly added to an acid like potassium hydrogen phthalate (KHP). The chemical reaction that takes place in the flask is a neutralization reaction. In this neutralization reaction, one mole of base neutralizes one mole of acid, resulting in salt and water. This reaction is performed in the presence of the indicator phenolphthalein, which is colorless at the start of the reaction when the pH is acidic. The indicator turns pink as soon as enough NaOH is added to the flask to make the pH basic.

   • Label the three Erlenmeyer flasks as ‘A’, ‘B’, and ‘C’.
   • Weigh 0.5 – 0.7 g of KHP for each of the flasks and record the mass for each. Note: Try to measure the same mass of KHP for all three flasks.

     Table 2. Standardization of NaOH

     Molar massKHP = 204.23 g/mol	Flask A	Flask B	Flask C
     Mass of KHP (g)
     Volumeinitial NaOH(mL)
     Volumefinal NaOH (mL)
     VolumeNaOH (mL)
     Moles of KHP
     Moles of NaOH
     Molarity of NaOH
     Average molarity
     Standard deviation

     Click Here to download Table 2
   • Add 50 mL of deionized water to each flask, and use a glass stirring rod to stir the solutions until they appear homogeneous.
   • Add 2-3 drops of phenolphthalein to each of the three flasks.
   • Set up the titration apparatus. Attach the burette clamp to the ring stand, and securely clamp the 50-mL glass burette to it. Make sure the valve of the burette is closed.
   • Label a 400-mL beaker as ‘waste’ and place it under the burette. Rinse the burette by pouring about 5 mL of 0.1 M NaOH into the burette. Open the burette valve to allow the NaOH to flow into the beaker.
   • Close the valve and fill the burette with slightly more than 50 mL of NaOH. Release any air bubbles present in the tip of the burette by opening and closing the burette valve. Record the starting volume of NaOH.
   • Place flask A below the tip of the burette, and titrate the solution using 1 mL volumes of NaOH. Swirl the solution after each addition. Continue adding 1 mL volumes to the flask until the pink color persists. This is considered the endpoint. Record the volume of 0.1 M NaOH added to achieve the endpoint.
   • Repeat the titration for flasks B and C. This data will be used to calculate the actual concentration of NaOH.
3. Titration of phosphoric acidExpand

   In this experiment, we will determine two of the three pKa values for the triprotic acid, phosphoric acid, using acid-base titration. In this neutralization reaction, phosphoric acid reacts with NaOH to form water and the salt, sodium phosphate.

   • Attach the drop counter to the ring stand, below the burette clamp. Secure the plastic burette so that its tip is just above the drop counter.
   • Connect the drop counter to the data acquisition system, and make sure the two valves of the plastic burette are both in the closed position.
   • Place the waste container below the burette and pour a few mL of 0.1 M NaOH into the burette. Open both valves to drain the NaOH into the waste beaker. Then close the valves.
   • Fill the plastic burette with 25 mL of 0.1 M NaOH. Drain about 5 mL into the waste beaker—enough so that NaOH fills the burette tip. Make sure there are no air bubbles, then close the valves.
   • Calibrate the drop counter. Replace the waste beaker beneath the burette with a 10-mL graduated cylinder. Then, open the bottom valve on the burette, while keeping the top valve closed. Turn on the data acquisition system and set it to ‘drop counting mode’.
   • Slowly open the top valve to very slowly release drops, ideally at one drop every 2 s. Allow the drops to empty from the burette until there is 9-10 mL of 0.1 M NaOH in the graduated cylinder.
   • Close the bottom valve and leave the top valve as is. Read the volume of NaOH in the graduated cylinder to the first decimal place and enter this value in the data acquisition system. Record the value for ‘drops/mL’. Then, discard the NaOH in the graduated cylinder into the aqueous waste beaker.
   • Calibrate the pH sensor before starting the titration. Connect the pH sensor to the data acquisition system, then select 'Calibrate'. Rinse the bulb of the pH sensor with deionized water before inserting it into the pH 7 buffer. Leave the sensor submerged until the voltage stabilizes, then accept the measurement.
   • Rinse the bulb with deionized water and insert it into the pH 10 buffer. Allow the voltage to stabilize, then accept the measurement.
   • Rinse the bulb of the pH sensor again and slide it through the designated slot in the drop counter.
   • Measure 40 mL of deionized water into a clean 100-mL beaker, then, transfer 1 mL of phosphoric acid into the beaker of water.
   • Place the beaker on the stir plate under the drop counter. Carefully slide the pH sensor into the beaker.
   • Add a stir bar to the beaker, and turn the stir setting onto high. Start collecting data on the acquisition device. Then, open the bottom valve on the burette. Note: The drop rate should be about 1 drop every 2 s. After the first drop is released, check to see data are being recorded.
   • Continue the titration until the pH meter reads pH 12. Then, stop acquiring data and close the valve on the burette. Save your data on a flash drive.
   • To clean up your workspace, check the pH of all waste solutions using pH paper. Neutralize all acidic aqueous waste with baking soda and all basic waste with citric acid. Add enough of either baking soda or citric acid to the solution until it stops bubbling.
   • Flush all neutralized solutions down the sink with copious amounts of water.
   • Wash all glassware.
4. ResultsExpand
   • In the first part of this lab, you standardized a solution of NaOH using KHP to determine its actual concentration. Now, let's see how close the standardized concentration is to the 0.1 M concentration that was prepared. Determine the number of moles of KHP that was added to each flask, and by extension, the moles of NaOH. Once the solution is neutralized, the molar quantities of KHP and NaOH are equal.
   • Calculate the molarity of NaOH based on the total volume of NaOH that was added to each flask. The actual concentration is lower than the expected 0.1 molarity. This is because NaOH is hygroscopic, so it is difficult to weigh it accurately.
   • Plot the results for the phosphoric acid titration (pH vs. Volume of NaOH). Phosphoric acid is a weak triprotic acid, meaning that it has the potential to provide three protons per molecule when it dissociates in aqueous solutions. Phosphoric acid has three pKa values, one for when each proton is dissociated.
   • Look at the data. There are two sigmoidal curves, indicating two equivalence points. Each equivalence point corresponds to a dissociation constant, Ka, of phosphoric acid. Note: You stopped the experiment once the pH reached 12, so you only measured two of the three Ka values.
   • Plot the first derivative of the titration curve. The equivalence points are represented by the curve maximums.

     Table 3. Titration of phosphoric acid

     Volume1st equivalence point (mL)
     Volume1st half-equivalence point (mL)
     Volume2nd equivalence point (mL)
     Volume2nd half-equivalence point (mL)
     1st pKameasured
     1st pKatheoretical	2.16
     2nd pKameasured
     2nd pKatheoretical	7.21
     Moles of NaOH
     Moles of H3PO4
     Molarity of H3PO4

     Click Here to download Table 3
   • Find the first half-equivalence point by taking the volume of NaOH corresponding to the first equivalence point and dividing it by 2. At the half-equivalence point, the concentrations of undissociated acid and its conjugate base are the same, and the pH is equal to the pKa.
   • Look up the pH at this volume from your table of data to get a more precise value. This corresponds to the first pKa, which is reported in the literature as 2.16.
   • Repeat this to find the second pKa. The second half-equivalence point is located midway between the first and second equivalence points, which should give a pKa of about 7.2.

To begin the experiment, you will need to prepare approximately 0.1 molar sodium hydroxide using a 50 weight percent solution of sodium hydroxide. So, how much of the 50 weight percent solution do you need in order to prepare 500 milliliters of 0.1 molar sodium hydroxide? The 50 weight percent sodium hydroxide is indicative of its weight ratio.

So, if the instructor prepared 150 milliliters of solution, then 150 grams of sodium hydroxide was dissolved in 150 grams of water, and the total weight is 300 grams. Since the density of the 50 weight percent solution is 1.53 grams per milliliter, you can calculate the volume of the solution, V, in milliliters. The molar mass of sodium hydroxide is 39.998 grams per mole.

Therefore, you can solve for the number of moles, X, in the 50 weight percent solution. Then, using these two values, you can calculate the molarity M1.Now, use the following dilution formula to solve for the volume, V1, of the 50 weight percent sodium hydroxide solution with molarity M1 needed to make 500 milliliters, V2, of 0.1 molar sodium hydroxide, M2.You need to know the value of V1 before you start the experiment. To begin, put on the appropriate personal protective equipment, including gloves, chemical splash goggles, and a lab coat, which must be worn at all times.

Now, label the 500-milliliter polyethylene bottle as 0.1 molar sodium hydroxide'Then, adjust the volume on a 1-milliliter pipette to the value you calculated. Attach a tip to the pipette and use it to transfer the 50 weight percent sodium hydroxide to the polyethylene bottle. Now, use a 100-milliliter graduated cylinder to measure the amount of water.

You will need 500 milliliters minus the calculated volume of sodium hydroxide, so 497.4 milliliters. Measure this volume of water, and pour it into the bottle containing the sodium hydroxide. Once all of the water has been added, cap the bottle tightly and invert it several times to mix the solution.

Now that the sodium hydroxide has been prepared, we will determine its exact concentration, or standardize it, using the acid-base titration method. In this technique, a base like sodium hydroxide is slowly added to an acid like potassium hydrogen phthalate, or KHP. The chemical reaction that takes place in the flask is a neutralization reaction.

Here, one mole of base neutralizes one mole of acid, resulting in salt and water. This reaction is done in the presence of the indicator phenolphthalein, which is colorless at the start when the pH is acidic but turns pink as soon as you add enough sodium hydroxide to the flask to make the pH basic. To begin, label the three Erlenmeyer flasks as A'B'and C'Then, weigh 0.5 to 0.7 grams of KHP.

Record the mass and pour it into flask A.Weigh KHP for each of the other two flasks, doing your best to measure the same mass of KHP as flask A.Now, measure 50 milliliters of deionized water, and pour it into flask A.Stir the solution with a glass stirring rod until the mixture appears homogeneous. Repeat this for flasks B and C.Next, obtain the 1%phenolphthalein drop bottle from your instructor, and add two to three drops to each of the three flasks. Now, set up the titration apparatus by first attaching the burette clamp to the ring stand.

Then, securely clamp the 50-milliliter glass burette to the burette clamp. Make sure the valve of the burette is in the off position, which is positioned perpendicular to the burette. Then, label a 400-milliliter beaker as waste'and place it under the burette.

Rinse the burette by pouring about five milliliters of 0.1 molar sodium hydroxide into the burette. Open the burette valve to let all of the liquid flow into the beaker. Close the valve, and fill the burette with slightly more than 50 milliliters of sodium hydroxide.

Open the valve to release any air bubbles present in the tip of the burette. Then close the valve and record the starting volume of sodium hydroxide. Now, place flask A below the tip of the burette, and titrate the solution using 1-milliliter volumes of sodium hydroxide.

Swirl the solution after each addition. Continue adding 1-milliliter volumes to the flask until the pink color persists. This is considered the endpoint.

Record the volume of 0.1 molar sodium hydroxide added to achieve the endpoint. Now, repeat the titration for flasks B and C.You'll use this information later to calculate the actual concentration of sodium hydroxide. In this experiment, we will determine two of the three pKa values for the triprotic acid phosphoric acid using acid-base titration.

To understand these terms, review the concepts behind this lab. The reaction that's taking place in this experiment is again a neutralization reaction. Here, phosphoric acid reacts with the sodium hydroxide to form water and the salt sodium phosphate.

To set up the experiment, attach the drop counter to the ring stand with the burette clamp above it. Now, secure the plastic burette so that its tip is just above the drop counter. Connect the drop counter to the data acquisition system, and make sure the two valves of the plastic burette are both in the closed position.

Place the waste container below the burette and pour a few milliliters of 0.1 molar sodium hydroxide into the burette. Then, open both of the valves to drain the sodium hydroxide into the waste beaker. Now, close the valves again and fill the plastic burette with 25 milliliters of the 0.1 molar sodium hydroxide.

Drain about five milliliters into the waste beaker so that sodium hydroxide fills the burette tip. Make sure there are no air bubbles, then close the valves. Now, let's calibrate the drop counter.

Replace the waste beaker beneath the burette with a 10-milliliter graduated cylinder. Then, open the bottom valve on the burette while keeping the top valve closed. Turn on the data acquisition system, and make sure it is set to drop counting mode.

Then, begin the calibration. Slowly open the top valve so that the drops are released very slowly, ideally at one drop every two seconds. The drop count data should begin displaying on the screen.

Allow the drops to empty from the burette until there is 9 to 10 milliliters of 0.1 molar sodium hydroxide in the graduated cylinder. Now, close the bottom valve, and leave the top valve as is. Read the volume of sodium hydroxide in the graduated cylinder to the first decimal place, and enter this value in the data acquisition system.

Record the drops per milliliter value displayed on the screen in your lab notebook. Then, discard the sodium hydroxide in the graduated cylinder into the aqueous waste beaker. Now, let's calibrate the pH sensor before starting the titration.

Connect the pH sensor to the data acquisition system. Then, select Calibrate'Begin the calibration with the pH 7 buffer. Rinse the bulb of the pH sensor with deionized water before inserting it into the vial containing the pH 7 buffer.

Leave the sensor submerged until the voltage stabilizes. Then accept the measurement. Next, calibrate the probe with a second buffer.

Rinse the bulb with deionized water, and insert it into the vial containing the pH 10 buffer. Allow the voltage to stabilize, then accept the measurement. Now, rinse the bulb of the pH sensor again, and slide it through the designated slot in the drop counter.

Now that the equipment is set up and calibrated, let's start the titration of phosphoric acid. First, measure 40 milliliters of deionized water and pour it into a clean 100-milliliter beaker. Then, obtain 0.5 molar phosphoric acid from your instructor.

Adjust the pipette to the 1-milliliter setting, attach a new pipette tip, and transfer 1 milliliter of phosphoric acid into the beaker of water. Now, place the beaker on the stir plate under the drop counter. Lift the pH sensor slightly if needed.

Then, carefully slide the pH sensor into the beaker. Add a stir bar to the beaker, and turn the stir setting on to high. Start collecting data on the acquisition device.

Then, open the bottom valve on the burette. The drop rate should be about one drop every two seconds. After the first drop is released, check that the data are being recorded.

Continue the titration until the pH meter reads pH 12. Then, select the Stop'setting on the data acquisition system and close the valve on the burette. Save your data on a flash drive to analyze later.

Now, let's clean up our workspace. Check the pH of all of your waste solutions using pH paper. Neutralize all acidic aqueous waste with baking soda and all basic waste with citric acid.

Add enough of either baking soda or citric acid to the solution until it stops bubbling. Then, check the pH to make sure it is neutral. Flush all neutralized solutions down the sink with copious amounts of water.

Then, wash all glassware. Now, let's take a look at how to analyze our results. In the first part of this lab, you standardized a solution of sodium hydroxide using KHP to determine its actual concentration.

Now, let's see how close the standardized concentration is to the 0.1 molar concentration that was prepared. We recorded the mass of KHP that was added to each flask. So from that, we can calculate the number of moles of KHP, and by extension, the number of moles of sodium hydroxide since the molar quantities are equal once the solution is neutralized.

We also know the total volume of sodium hydroxide that was added to the flask, so we can calculate its molarity. If we do this for all three standardization trials, we see that the actual concentration is lower than the expected 0.1 molarity. This is because sodium hydroxide is hygroscopic, so it is difficult to weigh it accurately.

Now, let's examine the results for the phosphoric acid titration. Phosphoric acid is a weak triprotic acid, meaning that it has the potential to provide three protons per molecule when it dissociates in aqueous solutions. Thus, it has three pKa values, one for when each proton is dissociated.

Looking at the data, there are two sigmoidal curves. Thus, there are two equivalence points, each of which corresponds to a dissociation constant, Ka, of phosphoric acid. You stopped the experiment once the pH reached 12, so you only measured two of the three Ka values for the acid.

To calculate the equivalence points more accurately, plot the first derivative of the titration curve. The equivalence points are represented by the curve maximums. Now, take the volume of sodium hydroxide corresponding to the first equivalence point and divide it by 2 to find the first half-equivalence point.

Here, the concentrations of undissociated acid and its conjugate base are the same, and the pH is equal to the pKa. Look up the pH at this volume from your table of data to get a more precise value, which we found to be 2.7. This corresponds to the first pKa, which is reported in the literature as 2.16.

You can see that the values are close. Repeat this to find the second pKa. The second half-equivalence point is located midway between the first and second equivalence points, which should give a pKa of 7.21.