Materials Included In Kit

Additional Materials Required

Safety Precautions

Dilute acid and base solutions are severely irritating to eyes and skin and slightly toxic by ingestion and inhalation. Phenolphthalein indicator solution is an alcohol-based solution. It is a flammable liquid and is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Flush all neutralized solutions down the drain with an excess of water according to Flinn Suggested Disposal Method #26b. Neutralize and dispose of unwanted acid solutions according to Flinn Suggested Disposal Method #24b. Neutralize and dispose of unwanted base solutions according to Flinn Suggested Disposal Method #10.

Lab Hints

• Enough materials are provided in this Super Value Kit for 5 classes of 30 students each, working in pairs (75 total student groups). This laboratory activity may reasonably be completed in one 50-minute class period. The prelaboratory assignment may be completed before coming to lab and the data compilation and questions may be completed the day after the lab.

Teacher Tips

• The identities of the unknowns are:
  • Unknown HCl 1 = 0.2 M
  • Unknown HCl 2 = 0.3 M
  • Unknown HCl 3 = 0.4 M
  • Unknown NaOH 1 = 0.3 M
  • Unknown NaOH 2 = 0.4 M
  • Unknown NaOH 3 = 0.5 M
• Volumetric flasks are the preferred equipment for preparing the THAM solutions. However, because they are expensive, they may not be available for each student group. If you have just a few volumetric flasks, it is recommended that your students take turns preparing their THAM solutions in the available volumetric flasks. Once a group’s solution is prepared, have them transfer it to a labeled bottle and pass the volumetric flask on to the next group. If there are not enough volumetric flasks available to share among the class, have students use beakers that have been calibrated to hold 100 mL.
• The bromcresol green indicator solution changes from a yellow color in solutions with a pH below 3.8, to a green color, to a blue color in solutions with a pH above 5.4. The endpoint is when the green color is reached; however, when using microscale equipment, such as pipets and well plates, the green color is hard to achieve. In practice, students will observe the yellow color turn directly to blue with the addition of a single drop. Good results can still be obtained by counting the number of drops until the appearance of the blue color.
• Make sure that students always hold the pipets vertically when delivering drops to the well plates. If the pipets are held at different angles, the drop size will vary, which changes the number of drops per milliliter. Since the number of drops per milliliter was calculated holding the pipet vertically, the pipets should be held vertically during the remainder of the experimental procedure.
• Have students rinse their well plates immediately after completing the titrations. Clean each well thoroughly with a cotton swab or piece of paper towel.

Answers to Prelab Questions

1. Calculate the number of grams of tris-(hydroxymethyl) aminomethane (THAM), H₂NC(CH2OH)₃, required to prepare 100 mL of a 0.1 M THAM solution. Show all work. Describe how to prepare this solution in lab. Hint: Use steps 1–6 and Equations 1–3 in the Preparing Solutions section in the Background as a guide. Use steps 1 and 2 from the Preparing Solutions section in the Background.
   {13759_PreLabAnswers_Table_2}
   Detailed instructions for preparing solutions are provided in the Preparing Solutions section of the Background. An overview of the solution preparation is: dissolve 1.2 g of THAM in enough distilled or deionized water to prepare a total of 100 mL of solution.

Sample Data

Answers to Questions

Part A. Determining the Molarity of an Unknown Hydrochloric Acid Solution

Calculations

1. Using the average number of drops of THAM recorded in the data table and Equation 4 from the Background section, calculate the number of moles of THAM required to neutralize the HCl. Show all work. Calculate the number of moles of THAM required using the following equation
   {13759_Answers_Equation_10}
   Unknown HCl 1 (0.2 M)
   {13759_Answers_Equation_11}
   Unknown HCl 2 (0.3 M)
   {13759_Answers_Equation_12}
   Unknown HCl 3 (0.4 M)
   {13759_Answers_Equation_13}
2. The balanced equation for the reaction between HCl and THAM follows. Using the balanced equation and Equation 5 from the Background section, calculate the number of moles of HCl neutralized. Show all work.

   HCl(aq) + H₂NC(CH₂OH)₃(aq) → Cl^– H₃N+C(CH₂OH)₃(aq)

   In each case, the number of moles of HCl will be equal to the number of moles of THAM because there is a 1:1 ratio of HCl:THAM in the balanced equation. Students should calculate this using the following equation:

   {13759_Answers_Equation_14}
3. Using the number of drops of HCl recorded in the data table, the number of moles of HCl from Question 2, and Equation 6 from the Background section, calculate the molarity of the unknown HCl solution. Show all work. Write and circle the unknown number and the calculated molarity.

   Calculate the Molarity of the HCl solution using the following equation:

   {13759_Answers_Equation_15}
   Unknown HCl Solution 1 (0.2 M)
   {13759_Answers_Equation_16}
   Unknown HCl Solution 2 (0.3 M)
   {13759_Answers_Equation_17}
   Unknown HCl Solution #3 (0.4 M)
   {13759_Answers_Equation_18}

Part B. Determining the Molarity of an Unknown Sodium Hydroxide Solution

4. Using the data collected in the data table, calculate the molarity of the unknown NaOH solution. Show all work. Write and circle the unknown number and the calculated molarity.

   Unknown NaOH Solution 1 (0.3 M)
   Step 1—

   {13759_Answers_Equation_19}

   Step 2—H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l)

   {13759_Answers_Equation_20}

   Step 3—

   {13759_Answers_Equation_21}

   Unknown NaOH Solution 2 (0.4 M)
   Step 1—

   {13759_Answers_Equation_22}

   Step 2—H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l)

   {13759_Answers_Equation_23}

   Step 3—

   {13759_Answers_Equation_24}

   Unknown NaOH Solution 3 (0.5 M)
   Step 1—

   {13759_Answers_Equation_25}

   Step 2—H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l)

   {13759_Answers_Equation_26}

   Step 3—

   {13759_Answers_Equation_27}

Questions

5. Why is it important to stir the solution in the well plate with the toothpick between the addition of each drop of solution from the pipet?

   Stirring between the addition of each drop ensures that the solution is homogeneous and the OH^– ions and H⁺ ions are spread uniformly about the solution. In a homogeneous solution, the endpoint will most accurately tell when the number of OH^– ions is equal to the number of H⁺ ions. If the solution was not stirred, there might be a localized excess of OH^– ions which would make the phenolphthalein indicator solution turn pink in that area which could be misread as the endpoint.

6. Name possible sources of error in this experiment.

   Student answers will vary. Possible answers include

   1. Holding the pipet at various angles instead of always the same angle when delivering drops.
   2. Not stirring the solution in the well plate between the addition of each drop.
   3. Losing a drop if a drop lands on the side of the well plate instead of into the bulk solution.
   4. Overshooting the endpoint (too pink).
   5. Water was used when calculating the number of drops in 1 mL. However, acid and base solutions, not water, were delivered using pipets. Perhaps the acids and bases actually contain a different number of drops per mL than water.
7. Propose another method (besides titration) for determining the concentration of a solution. Explain how the concentration would be determined using this method.

   Student answers will vary. Possible answers include

   1. If the solution was colored, then the depth of color will vary with concentration. A series of stock solutions of known concentrations could be prepared and set up in a row. The color intensity will decrease as the concentration decreases. The unknown could then be placed according to its color and its concentration estimated to be between the knowns on either side of it.
   2. As concentration varies, so does density, absorbance, and conductivity. A series of stock solutions with increasing concentration could be prepared as knowns. Then the density, absorbance, or conductivity of these stock solutions could be measured and plotted on a graph. The unknown could then be tested in the same manner and its concentration determined from the graph of stock solutions.

Additional Practice Calulations

8. Calculate the molarity of a solution of potassium chloride, KCl, prepared using 54 g of solid and a total solution volume of 1.5 L. Show all work. First, the number of moles in 54 g of solid KCl must be calculated using the molecular weight (MW).
   {13759_Answers_Equation_28}
   Then, the molarity can be calculated using Equation 1 from the Background section.
   {13759_Answers_Equation_29}
9. Describe how you would prepare 1.00 L of a 0.500 M solution of cupric sulfate pentahydrate, CuSO₄•5H₂O.

   First, the number of moles of solid required to make this solution must be calculated using Equation 2 from the Background section.
   moles = Molarity x volume of solution
   moles = 0.500 M x 1.00 L moles = 0.500 moles
   
   Then the moles are converted to grams using the molecular weight (make sure to include the mass of the 5H₂O in the molecular weight).
   mass = moles x MW
   mass = 0.500 moles x 249.7 g/mole
   mass = 125 g
   
   To prepare the solution, add 125 g of solid to a 1-L volumetric flask. Dissolve the solid in a minimum amount of distilled or deionized water. Cap the flask and invert several times to mix the solution. Add water slowly and at the end drop by drop to the 1-L mark such that the bottom of the meniscus sits exactly at the 1-L mark. Cap the flask and invert several times to prepare a homogeneous solution.

10. How many grams of potassium iodide, KI, are present in 275 mL of a 0.23 M solution? Show all work.
    First, the number of moles present in 275 mL of a 0.23 M solution must be calculated using Equation 2 from the Background section.
    {13759_Answers_Equation_30}
    Then, the number of moles can be converted into grams using the molecular weight (MW).
    grams = moles x MW 
    grams = 0.063 moles x 166 g/mole
    grams = 10 grams

References

Griswold, N. E.; Neidig, H. A.; Spencer, J. N.; Stanitski, C. Laboratory Handbook for General Chemistry; Chemical Education Resources: Palmyra, PA, 1996; pp 23–24.

Introduction

How is the concentration of a solution determined? With acid and base solutions, titrations are commonly used to determine unknown concentrations. In this laboratory activity, the concentration of an unknown hydrochloric acid solution and an unknown sodium hydroxide solution will be determined using acid/base titrations.

Concepts

• Solutions

• Molarity
• Concentration
• Titrations

Background

Preparing Solutions

The amount of solute that is dissolved in a given quantity of solvent is expressed as the concentration of the solution. A dilute solution contains only a small amount of solute in a given amount of solution, while a concentrated solution contains a large amount of solute in a given amount of solution. The units most commonly used to describe the concentration of solutions are molarity units. The molarity, M, of a solution is the number of moles of solute in one liter of solution. To determine the molarity of a solution, Equation 1 can be used:

Used in conjunction with the molecular weight, MW, of a solute, Equation 1 is used to determine the number of grams of solute needed to prepare a given volume of a solution with a specific concentration. For example, consider the preparation of 500 mL of a 0.80 M solution of sodium chloride, NaCl. The following steps outline this procedure.

Step 1—Determine the number of moles necessary to prepare this solution. To do this, rearrange Equation 1 to solve for moles.

moles of solute = Molarity x volume of solution

moles of NaCl = 0.40 moles

Therefore, distilled or deionized water must be added to 0.40 moles of NaCl and the resulting solution diluted to a total volume of 500 mL to prepare a 0.80 M solution.

Step 2—Convert the number of moles to grams using the molecular weight. Therefore, 23 g of NaCl is required to prepare 500 mL of a 0.80 M sodium chloride solution.

Once the calculations have been done to determine how much solute is needed to correctly prepare the solution, precise analytical techniques must be followed when actually making the solution. The steps below outline proper solution preparation procedures.

Step 3—Obtain a piece of volumetric glassware calibrated to the volume needed. Volumetric glassware is glassware that has been calibrated (and marked) to hold a specific volume. The most common form of volumetric glassware used for preparing solutions is the volumetric flask, a flask that has a long, narrow neck with a marking on it. For a 100-mL volumetric flask, the mark on the neck indicates that when filled to the mark, the flask will contain exactly 100 mL. Because volumetric flasks are expensive, they may not be available for every student lab group. However, solutions are not commonly stored in volumetric flasks, so only a few volumetric flasks are necessary for an entire class to prepare solutions. One group can prepare a solution, then empty the solution into a labeled storage bottle and pass the volumetric flask on to another group. If no volumetric glassware is available for preparing solutions, the glassware must be calibrated before preparing the solution. To calibrate a piece of glassware, a specified volume is poured into the container and the liquid level marked.

Step 4—Precisely weigh out the required number of grams (determined in step 2) of solid on a balance in a weighing dish. Transfer the solid to a clean, dry beaker (with a larger capacity than the necessary volumetric flask). There may be a few grains of solid left on the weighing dish, so use a wash bottle filled with distilled or deionized water to rinse any remaining solid from the weighing dish into the beaker. Rinse the weighing dish several times to make sure that all of the solid was transferred. This process of transferring every bit of the solid is called quantitatively transferring. The rinse water may be enough water to dissolve all of the solid in the beaker. If it is not, add a minimum amount of distilled or deionized water to dissolve any remaining solid.

Step 5—Using a funnel, transfer the solution in the beaker to the volumetric flask. Rinse the beaker with a small amount of distilled or deionized water, transferring the rinse water through the funnel into the volumetric flask. Rinse from the beaker through the funnel into the flask several times to thoroughly rinse the beaker and the funnel.

Step 6—Fill the volumetric flask with distilled or deionized water. When the flask is about one-half to two-thirds full, cap the flask and invert it several times to make sure the solution is homogeneous. Continue filling the flask until the liquid level is almost at the mark. Fill to the mark with a pipet or wash bottle containing distilled or deionized water drop-by-drop until the bottom of the meniscus is directly on the mark. Again cap the flask and invert it several times to thoroughly mix the solution. Solutions are not generally stored in volumetric flasks, so transfer the solution to a labeled bottle and cap the bottle to prevent evaporation or contamination.

Titrations

Several techniques can be used to determine the concentration of a solution. When the solution of unknown concentration is an acid or base, one of the most widely used techniques involves carrying out a titration. In a titration involving an acid solution of unknown concentration, a given volume of the acid is placed in a container with an indicator. Then a base solution of known concentration is added to the acid solution until the endpoint is reached. The endpoint is the point at which the number of moles of hydroxide ions, OH^–, is equal to the number of moles of hydrogen ions, H⁺, in solution. The indicator is chosen so that it changes color at the endpoint. How does it do this? Refer to Figure 1 to follow the progress of a titration between hydrochloric acid, HCl, and sodium hydroxide, NaOH. Phenolphthalein is the indicator used in this titration. Before the Titration Begins. Before any NaOH is added to the acid, there are many H⁺ ions in solution. The presence of free H⁺ ions makes the solution acidic (pH below 7). At this pH, the phenolphthalein indicator is colorless.

During the Titration. As NaOH is added, the OH^– ions from the NaOH pair up with the H⁺ ions to make water molecules. As this occurs, the number of free H⁺ ions in solution decreases and the pH begins to rise. As long as there are more H⁺ ions in solution than OH^– ions, however, the pH is still below 7 and the phenolphthalein indicator is still colorless.

At the Endpoint. Recall that the endpoint of the titration is defined as the point when the number of OH^– ions added is exactly equal to the number of H⁺ ions in solution. In reality, because the NaOH is being added in drops, and each drop contains about 10¹⁹ ions, a titration cannot determine exactly when the number of ions are equal. But, it can come very close. Therefore, in practice, the endpoint is defined as the point at which the indicator changes color. In this titration, by adding just a single drop, the solution turned from colorless to pink. Therefore, at this point, it is assumed for calculation purposes, that the number of moles of OH^– ions is equal to the number of moles of H⁺ ions.

Exactly how can the data collected from the above titration be used to calculate the concentration of the HCl solution? Steps 1–3 outline the basic procedure for such titration calculations. For this titration, the following data was taken. Step 1—Determine the number of moles of base needed to neutralize the acid. Two factors must be known to calculate the number of moles of base needed to neutralize the acid—the volume of base delivered and the molarity of the base. The volume delivered is determined by counting the number of drops of NaOH added and converting drops to liters. The molarity of the NaOH is known. The number of moles of NaOH added to reach the endpoint (the pink color change) is calculated using Equation 2. Step 2—Use the balanced chemical equation to determine the number of moles of acid neutralized. The balanced chemical equation must be considered to determine how many moles of NaOH are needed to neutralize each mole of HCl.

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

From the balanced equation, it can be observed that one mole of HCl reacts with one mole of NaOH. Therefore, the ratio of NaOH to HCl is 1:1 (1 mole HCl/1 mole NaOH). Using the number of moles of base calculated in Step 1 and this ratio, the moles of acid can be calculated. It is important to make sure that the ratio is used correctly in this calculation. Set up the calculation so that moles_base cancel leaving only moles_acid.

moles_HCl = 0.00043 moles

Step 3—Calculate the molarity of the acid. To determine the molarity of the HCl solution, Equation 1 is used, substituting in the appropriate conversion factors for drops/L. A similar procedure is used to calculate the concentration of a base solution using an acid solution of known concentration. It must be pointed out that step 2 contains a critical step. In the example shown here, the moles of acid were equal to the moles of base because the ratio of the coefficients in the balanced chemical equation was 1:1. However, in a reaction involving a diprotic or triprotic acid, such as H₂SO₄ or H₃PO₄, the ratio will no longer be 1:1, but will instead be 1:2 or 1:3. This new ratio must be accounted for in step 2.

Materials

Prelab Questions

1. Calculate the number of grams of tris-(hydroxymethyl) aminomethane (THAM), H₂NC(CH₂OH)₃, required to prepare 100 mL of a 0.1 M THAM solution. Show all work on a separate sheet of paper. Describe how to prepare this solution in lab. Hint: Use steps 1–6 and Equations 1–3 in the Preparing Solutions part in the Background section as a guide.

Safety Precautions

Dilute acid and base solutions are severely irritating to eyes and skin and slightly toxic by ingestion and inhalation. Phenolphthalein indicator solution is an alcohol-based solution. It is a flammable liquid and is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part A. Determining the Number of Drops in One Milliliter

1. Fill a pipet with water. Holding it vertically, add water, drop by drop, to a clean, dry graduated cylinder. Count the number of drops required to fill the graduated cylinder to the 1-mL mark. (Add drops until the bottom of the meniscus sits exactly at the 1-mL mark.) Avoid getting any water droplets on the sides of the graduated cylinder. Record the number of drops in the data table. Repeat, filling to the 2-mL and 3-mL marks, again counting the number of drops for each mL. Record these values in the data table. Average the three values and record the average in the data table.

Part B. Determining the Molarity of an Unknown Hydrochloric Acid Solution

2. Before proceeding with step 3, check your answer from the Prelab Exercise for the number of grams of THAM required to prepare 100 mL of a 0.1 M THAM solution.
3. Dissolve the number of grams of THAM calculated in the Prelab Exercise in enough distilled or deionized water to prepare 100 mL of a 0.1 M THAM solution. To prepare this solution, use your procedure written as part of the Prelab Exercise.
4. Label two clean pipets “HCl” and “THAM.” Fill the “HCl” pipet about half-full with the hydrochloric acid solution of unknown concentration. Fill the “THAM” pipet about half-full with the 0.1 M THAM solution.
5. Holding the “HCl” pipet vertically, add 10 drops to one of the wells in the 24-well plate. Add one drop of bromcresol green indicator solution to this well. Stir with a toothpick. Note the color of the solution. Record the number of drops of HCl added in the data table.
6. Holding the “THAM” pipet vertically, add the THAM solution drop by drop, counting the number of drops. Stir with the toothpick after the addition of each drop. Continue adding the THAM solution drop by drop until the solution turns blue. Record your observations and the number of drops required to turn the solution blue in the data table.
7. Using the same labeled pipets, repeat steps 4–6 two times to obtain two more sets of data. Perform each titration in a clean well in the 24-well plate. Record the data for these trials in the data table. Calculate the averages for the three trials and record them in the data table.

Part C. Determining the Molarity of an Unknown Sodium Hydroxide Solution

8. Label two clean pipets “H₂SO₄” and “NaOH.” Fill the “H₂SO₄” pipet about half-full with the 0.1 M sulfuric acid solution. Fill the “NaOH” pipet about half-full with the unknown sodium hydroxide solution.
9. Holding the “NaOH” pipet vertically, add 10 drops to one of the clean wells in the 24-well plate. Add one drop of phenolphthalein indicator solution to this well. Stir with a clean toothpick. Note the color change. Record the number of drops of NaOH added in the data table.
10. Holding the “H₂SO₄” pipet vertically, add the sulfuric acid solution drop by drop, counting the number of drops. Stir with the toothpick after the addition of each drop. Continue adding the sulfuric acid solution drop by drop until the solution turns and remains colorless. Record your observations and the number of drops required to turn the solution colorless in the data table.
11. Using the same labeled pipets, repeat steps 8–10 two times to obtain two more sets of data. Perform each titration in a clean well in the 24-well plate. Record the data for these trials in the data table. Calculate the averages for the three trials and record them in the data table.
12. Flush the solutions in the 24-well plate down the drain with plenty of water. Rinse each well thoroughly. Dry the 24-well plate with a paper towel and a cotton swab.

Student Worksheet PDF

13759_Student1.pdf