Materials Included In Kit

Additional Materials Required

Safety Precautions

All of the acids and bases used in this lab are very corrosive to eyes, skin and other body tissues. They are toxic by ingestion. Avoid all body tissue contact. Acetic acid and hydrochloric acid are also toxic by inhalation. Avoid breathing the vapors and dispense these chemicals in a fume hood. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult material safety data sheets and acid/base safety handling and disposal procedures for additional safety and handling techniques. Keep spill control materials on hand to neutralize acids or bases in case of spills. Use sodium carbonate or sodium bicarbonate to neutralize acid solutions. Use citric acid to neutralize base spills.

Disposal

All of the solutions may be flushed down the drain with plenty of water. The used paper strips should be discarded in the solid waste disposal.

Teacher Tips

• This kit is intended to be used in an introductory chemistry curriculum to teach some of the chemical reactions of acids and bases. The laboratory instruction sheets, data tables and questions and analysis pages are designed to be reproduced for student use.
• This kit contains materials for a class of 30 students working in pairs.
• A solid waste container should be made available for students to dispose of their indicator paper strips. Caution the students that the test paper strips are not to be thrown down the drain.
• Each group will need a 24-well microplate to perform this lab. These microplates may be purchased once and then reused. (Note: If these plates are not available, all of the tests may be performed in test tubes. Alternately, the tests may be performed on clear acetate sheets using smaller volumes of solutions.)
• If a mortar and pestle is not available for grinding the antacid tablets in parts B and F, the tablet may be crushed between two sheets of paper.
• Cassette tape cases are useful for transporting and holding Beral-type pipets.
• Remind students to hold the Beral-type pipets vertically for evenly sized drops. This is especially important in the titrations.
• The use of pH meters is an excellent addition for all steps of this lab, providing availability. A Checker pH meter (Catalog No. AP8673) is a great and inexpensive option for this lab.
• For other exciting pH labs and activities, refer to A pH4 Laboratory & Activity Manual: pHysiology, pHarmacology, and other pHantastic pHenomena by Carla R. Krieger available through Flinn Scientific (Catalog No. AP4571).
• For additional introductory acid–base labs, try Acid–Base Test Kit I: Examining Properties of Acids and Bases available from Flinn Scientific (Catalog No. AP4567).
  

  Part F Teaching Tips
  
  Suggested Procedure and Grading Information

• The open-ended format of Part F is designed to allow students to use the concepts of neutralization and techniques of microtitration to evaluate the “neutralizing power” of an antacid tablet.
• Standard HCl solutions in the range of 0.1 M to 1.0 M work well. Using 0.1 M HCl more closely approximates the conditions in the human stomach but will greatly increase the volume of acid needed and will lengthen the amount of time required for the experiment. For this reason, it is suggested that students use the 1.0 M HCl provided with this kit.
• Antacid tablets for Part F are provided in the kit. One of these antacid tablets weighs approximately 1.26 g. The active ingredients in these tablets are either calcium carbonate or a mixture of calcium carbonate and magnesium hydroxide. Note: Tablets included with the kit may vary due to availability.
• Have students imagine that the beaker is the stomach and an antacid tablet was recently added. As HCl is added (stomach acid is produced), the pH will drop and the stomach will become more acidic. The antacid will neutralize some of the acid and raise the pH back up to a near-neutral value. This process will continue until the “neutralizing power” of the antacid is gone. An antacid is considered ineffective in the stomach if the pH of the stomach falls below a value of 3.0. At this point, discomfort may be felt because the antacid is no longer capable of raising the pH back up to neutral.

  For this reason, a pH of 3.0 would be a reasonable pH value to use as an endpoint in this titration experiment. An indicator, such as the methyl orange provided with this kit, which has a color change between 3.0 (red) to 4.4 (yellow), would be useful in the detection of the endpoint. A pH meter, if available, would also be useful in this experiment. If a pH meter is used, the initial antacid solution will have a pH of approximately 8.50 to 9.50. Acid should be added with stirring until the pH drops and remains at 3.0 or below. The pH meter can also be used along with an indicator.

• Other indicators with a pH range of about 3.0 may also be used beside the methyl orange provided in this kit. Some options are bromphenol blue (3.0–4.6, yellow to purple) or congo red (3.0–5.0, blue to red). You may have various groups test various indicators and share results.
• As an extension, students may time the speed of the neutralizing action of the antacid provided or of a variety of commercial brands.
• It is recommended that students perform the experiment until two trials are in agreement. Results from the trials should be averaged.
• If materials are available, you may consider having students use burets and perform a traditional titration.

  Sample student procedure (teacher’s notes are provided in parentheses)

  1. Use a mortar and pestle to grind the antacid tablet into a powder. (This tablet is enough to perform 4 trials, if necessary.)
  2. Use a weighing dish to mass the entire antacid tablet. (1.28 g)
  3. Weigh approximately ¼ of the tablet into a small beaker. (0.31 g)
  4. Add 50 mL (or so) of water. (Note: Antacids are highly insoluble in water. For this reason, heating at low temperature on a hot plate may be necessary. A higher temperature also more closely mimics the conditions in the human stomach.)
  5. Add 3 drops of methyl orange indicator solution to the antacid solution. (Solution turns yellow.)
  6. Prepare methyl orange acid and base standards (for color comparison) by placing 3 drops of methyl orange indicator into a small beaker of dilute acid solution and into a small beaker of dilute base solution. Label these and set them to the side for comparison purposes.
  7. Begin the titration by slowly adding 1.0 M HCl dropwise with constant stirring, carefully counting drops. Start with approximately 5 drops.
  8. Stir the solution and note the color changes. (When HCl is added, the yellow color of the solution changes to red. As the antacid dissolves further and neutralizes the acid, the red color fades to orange and back to yellow.)
  9. Continue adding 1.0 M HCl solution in increments of approximately 5 drops with constant stirring, carefully counting the total number of drops added.
  10. Add 1.0 M HCl solution with stirring until the pH maintains a value of approximately 3.0. (This will be detected when the solution remains pink and no longer fades back to orange or yellow.) Record the total volume of acid, in drops, needed to reach this endpoint.
  11. Repeat the titration with the other portions of the antacid. If the second trial agrees with the first, average the results. If not, continue performing trials until two trials are in agreement.
  12. Since a 1.0 M HCl solution had been used in the titration, multiply the number of drops of acid added by 10. (Note: This accounts for the fact that the stomach pH is more similar to a 0.1 M HCl solution where 10 times more acid would be needed).
  13. Determine the number of drops in 1 mL. Convert the total number of drops to number of mL of acid. This is the amount of acid that can be neutralized by the amount of antacid used in the titration. Convert this to the amount of acid that may be neutralized by one full tablet. Record this in the data table.

  Sample results: (Note: Results may vary)

• Using 0.31 g of antacid (¼ tablet), the pH reached 3 (and the solution with the methyl orange indicator remained red) after approximately 70 drops of 1.0 M HCl was added. Since the concentration of stomach acid is closer to 0.1 M HCl, we can assume that 700 drops of 0.1 M HCl would have been consumed by 0.31 g of antacid. Assuming ~25 drops/mL, 700 drops is about 28 mL for 0.31 g of antacid. Thus extrapolating our results, the entire antacid tablet (1.26 g) would be able to neutralize 114 mL of HCl. Actual results may vary.

Further Extensions

Part F. Extension—Design a Neutralization–Titration Procedure Problem
Given an antacid tablet, design a procedure to test or to determine the average volume of acid (stomach acid = HCl) that can be neutralized per antacid tablet. After designing the experiment, your team will perform the experiment and compute results.

Guidelines

1. Each team of students will only have one single antacid tablet to use in order to solve this problem. It is suggested that the tablet be ground up to a powder and portions of it used for each trial (so as to have enough for multiple trials). One quarter of a tablet in 50 mL of water is enough for each trial.
2. Imagine that the beaker is your stomach and an antacid tablet was recently added. As HCl is added (stomach acid is produced), the pH will drop and the stomach will become more acidic. The antacid will neutralize some of the acid and raise the pH back up to a near-neutral value. This process will continue until the “neutralizing power” of the antacid is gone. An antacid is considered ineffective in the stomach if the pH of the stomach falls below a value of 3. At this point, discomfort may be felt because the antacid is no longer capable of raising the pH back up to neutral.

   For this reason, a pH of 3.0 would be a reasonable pH value to use as an endpoint in this titration experiment. An indicator, such as methyl orange which has a color change between 3.0 (red) to 4.4 (yellow), would be useful in the detection of the endpoint. A pH meter, if available, would also be useful in this experiment.

3. Prelab

   Before performing your experiment, create each of the following and check with your instructor before proceeding in the laboratory

   • Materials list (including all chemicals and equipment you will need)
   • Step-by-step procedure (which a classmate could follow if needed)
   • Safety procedures
   • ata table (where your data may be recorded)
4. Lab

   Perform the experiment which you and/or your team created.
   Repeat the experiment, gathering data for at least two valid trials.
   Determine the average volume of HCl neutralized per antacid tablet (convert the per-gram amount to a per-tablet amount).
   Pool your results with the class.
   Write a conclusion based on your results.

This experiment may be further extended to include a variety of brands of commercial antacids. Students may design an experiment to compare the “neutralizing power” of each brand. Students should have access to the antacid bottles and their prices so that they can find information on active ingredients, active ingredient amounts, costs per tablet, and recommended dosages. Before the lab, students can generate hypotheses about which antacids will neutralize the most acid and which will be the most cost-effective. Since students use indicators to determine the endpoint, white antacid tablets work better than colored ones as the pigments present in some antacids may obscure indicator color changes.

Sample Data

Table 1. Formation of a Salt

Table 2. An Antacid in Action Table 3. The Rainbow Reaction Table 4. Strength and Indicators Table 5. Titration Data

Answers to Questions

1. 1. The pH of 0.1 M HCl is 1; the pH of 0.1 M NaOH is 13.
   2. The pH of an equal mixture of the HCl and NaOH is 6–7.
2. The balanced chemical equation for the neutralization reaction is HCl + NaOH → NaCl + H₂O
3. The solid residue that is left on the depression slide after evaporating the solution is sodium chloride salt with the formula NaCl.
4. The universal indicator solution (in changing from red to orange to yellow to green to blue to violet) is an indication that the solution started acidic and became increasingly basic as the antacid tablet dissolved.
5. The color change in the solution from red to violet is the result of the neutralization reaction between the antacid and the acid. As the pH rises, the indicator changes color.
6. The balanced chemical equations for the neutralization reactions between the antacid tablet (which is calcium carbonate and magnesium hydroxide) and the acid is

   CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂CO₃(aq)
   where H₂CO₃(aq) → H₂O(l) + CO₂(g)
   and Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

7. The colors that were observed in the test tube from top to bottom after performing the rainbow reaction were red, orange, yellow, green, blue and violet.
8. The fact that the violet color is on the bottom of the tube indicates that the sodium carbonate solution has a greater density than the hydrochloric acid solution. As the sodium carbonate solution settles to the bottom of the tube, a neutralization reaction takes place between the sodium carbonate and the hydrochloric acid. At various depths, the concentration of the hydrogen (H⁺) ions varies, producing layers of varying pH. The universal indicator responds accordingly, with the rainbow effect resulting. Since molecular motion is relatively slow, this effect will remain for several weeks if left undisturbed.
9. The balanced net ionic chemical equations for the neutralization reactions between the sodium carbonate and the acid are as follows

   H⁺(aq) + CO₃^2–(aq) → HCO₃^–(aq)
   H⁺(aq) + HCO₃^–(aq) → H₂CO₃(aq)
   H₂CO₃(aq) → H₂O(l) + CO₂(g)

10. The balanced chemical equations for the following neutralization reactions are:
    1. Reaction between HCl and NaOH

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

    2. Reaction between H₂SO₄ and NaOH

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

    3. Reaction between HC₂H₃O₂ and NaOH

       HC₂H₃O₂(aq) + NaOH(aq) → Na(C₂H₃O₂)(aq) + H₂O(l)

11. From the balanced equations in Question 10:
    1. 1 mole of NaOH combines with 1 mole of HCl.
    2. 2 moles of NaOH combine with 1 mole of H₂SO₄.
    3. 1 mole of NaOH combines with 1 mole of HC₂H₃O₂.
12. Looking at Table 5:
    1. The ratio of drops of NaOH to HCl is 1:1.
    2. The ratio of drops of NaOH to H₂SO₄ is 2:1.
    3. The ratio of drops of NaOH to HC₂H₃O₂ is 1:1.
    4. Answers will vary, although it is hoped that the experimental results will confirm the predicted mole ratio from the balanced chemical equations in Question 10.
13. The monoprotic acids, HCl and HC₂H₃O₂, each have one H⁺ ion per molecule of acid to react with the one OH^– ion from each molecule of NaOH. Thus the ratio is 1:1. The diprotic acid, H₂SO₄, has two H⁺ ions per molecule of acid, indicating that each H₂SO₄ needs 2 OH^– ions. Therefore, 1 molecule of H₂SO₄ reacts with 2 molecules of NaOH and the ratio is 1:2.
14. 1. The 0.1 M HCl is a stronger acid than 0.1 M HC₂H₃O₂ (even though they have the same molar concentration of 0.1 M). The pH of the HCl is 1, indicating more H⁺ ions in solution and thus a higher acid strength, while the pH of the HC₂H₃O₂ is 3, indicating less H⁺ ions in solution and thus a weaker acid strength.
    2. The HCl required slightly more NaOH to reach neutralization than the HC₂H₃O₂, even though the approximate ratios were 1:1 for both. A strong acid, such as HCl, is fully ionized in solution forming H⁺ and Cl^–. A weak acid is only partially ionized in solution. This means that while there are some H⁺ and C₂H₃O₂^– ions, there are also some molecules of HC₂H₃O₂. Because there are less H⁺ ions in solution, the HC₂H₃O₂ requires less OH^– ions in order to reach the neutralization point. This effect would be more pronounced at higher volumes.
15. The concentration of the Unknown A HCl solution can be calculated as follows using the average data as shown in Table 5:

    From the balanced equation HCl + NaOH → NaCl + H₂O, we see that at the neutralization point, moles of HCl = moles of NaOH since there is a 1:1 ratio. Thus use (Molarity 1) x (Volume 1) = (Molarity 2) x (Volume 2), or M₁ x V₁ = M₂ x V₂, where 1 is the unknown HCl and 2 is the known NaOH, so (? M HCl) x (10 drops HCl) = (0.1 M NaOH) x (22 drops NaOH)

    {12832_Answers_Equation_4}
    Note: Since the volume of the acid and the base was measured in drops, there is no need to convert from drops to mL.
16. The concentration of the Unknown B solution can be calculated as follows using the average data as shown in Table 5

    From the balanced equation H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O, we see that at the neutralization point, 1 mole H₂SO₄ = 2 moles NaOH since the mole ratio is 1:2. Thus first find the moles of NaOH needed for neutralization using the equation M x V = moles, so (0.10 M NaOH) x (30 drops) = 3.0 moles NaOH
    From the 2:1 mole ratio, we see that 3.0 moles NaOH will neutralize 1.5 moles of H₂SO₄, because

    {12832_Answers_Equation_5}
    Thus M = moles/volume = 1.5 moles H₂SO₄/10 drops H₂SO₄ = ? M H₂SO₄ = 0.15 M H₂SO₄
    Note: Since the volume of the acid and the base was measured in drops, there is no need to convert from drops to mL.)
17. A sample student procedure and sample results are outlined under Part F Teaching Tips.

References

Herron, J. D.; Sarquis, J. L.; Schrader, C. L.; Frank, D. V.; Sarquis, M.; Kukla, D. A. Chemistry; D. C. Heath: Boston, MA, 1996; Chapter 19.

Source Book, Version 1.0.; Orna, M. V.; Schreck, J. O.; Heikkinen, H., Eds.; ChemSource: New York, 1994; Vol. 1, Chapter 2.

Introduction

Acids and bases! These are familiar substances which undergo many important chemical reactions. This kit will allow examination of various reactions of acids and bases.

Background

Acids are defined as substances that release hydrogen ions (H⁺) in solution while bases release hydroxide ions (OH^–) in solution. The positive hydrogen ion, which is a proton, and the negative hydroxide ion combine together to form a neutral water molecule (H₂O) according to the following equation

To express the concentration of hydrogen ions in solution, a term called pH (the power of hydrogen ions) is used. If the con-centration of H⁺ ions is greater than the concentration of OH^– ions, then the substance is considered acidic and has a pH value of lower than 7 (i.e., 1–6). If the concentration of OH^– ions is greater than the concentration of H+ ions, then the substance is basic and has a pH value greater than 7 (i.e., 8–14). If the H+ and OH– concentrations are equal, the substance is neutral and has a pH value of 7. pH is measured with either a strip of indicator test paper, an indicator solution, or an instrument called a pH meter. The red and blue colors of litmus test paper show the presence of acidity or basicity very clearly; however, they give no indication of the strength of the acid or base. Other, more sensitive types of test papers contain acid–base indicator dyes which turn different colors at different pH values. While these “universal” test strips provide fast and convenient measurements of a range of pH, they give only approximate values. For a more exact pH measurement, a pH meter should be used. A pH meter is a specially designed voltmeter connected to a pair of ion-selective electrodes that are dipped into the solution being examined. The voltmeter measures the cell potential between the two electrodes. This potential is a function of the activity of the hydrogen ions and is converted into a hydrogen ion concentration, or pH, value. The pH value using a pH meter can be measured to 0.1 pH unit or better.

Some acids and bases produce more ions in solution than similar amounts of other acids and bases. This is related to acid or base strength. A strong acid such as hydrochloric acid, HCl, dissociates nearly 100% into its ions, H⁺ and Cl^–. A weak acid such as acetic acid, HC₂H₃O₂, only partially dissociates into its ions, H⁺ and C₂H₃O₂^–, with the majority of it remaining in the molecular form, HC₂H₃O₂. A strong acid thus donates a greater number of H⁺ ions to the solution than a weak acid and will have a lower (more acidic) pH. The large number of ions in a strong acid allows the solution to conduct electricity and is termed a strong electrolyte. A weak acid conducts an electric current to a lesser extent and is termed a weak electrolyte. Strong and weak bases can be defined in a similar manner, except that the bases produce hydroxide ions in solution. A strong base such as sodium hydroxide, NaOH, donates a greater number of OH– ions to the solution than a weak base and will have a higher (more basic) pH. Thus, a strong base is a stronger electrolyte than a weak base such as ammonium hydroxide, NH₄OH.

Acids and bases have the ability to undergo many types of reactions. When an acid is mixed with a base, the H⁺ combine with the OH^– to produce water. Reactions between acids and bases are classified as neutralization reactions because the acid and the base neutralize each other to give a solution with a neutral pH of 7. The products of a neutralization reaction are a neutral salt plus water, according to the following equations Equation 2 shows a 1:1 ratio between HCl and NaOH, indicating that one molecule of HCl is neutralized by one molecule of NaOH. Equation 3 shows a 1:2 ratio between H₂SO₄ and NaOH. This indicates that each molecule of H₂SO₄ requires two molecules of NaOH for neutralization to occur. HCl is termed a monoprotic acid because it contains only one ionizable proton to donate to a base. H₂SO₄ is called a diprotic acid as it contains two ionizable protons to donate to a base.

A common acid–base neutralization reaction is the reaction of an antacid with stomach acid, hydrochloric acid. The stomach lining produces HCl during digestion. A normal adult produces two to three liters of dilute HCl each day. While the stomach produces a small amount of acid all the time, it can be stimulated to produce more acid by the presence of food. Too much food or stress may cause the stomach to respond with an outpouring of acid, lowering the stomach pH to the point of discomfort. Antacids are chemicals that neutralize acids, thereby relieving unpleasant effects from excess HCl in the stomach.

The active ingredients used in antacids differ with manufacturers; however, most common antacids contain weak bases such as sodium bicarbonate, calcium carbonate, magnesium hydroxide, aluminum hydroxide or combinations of these. The carbonate-containing antacids react with the stomach acid to produce a neutral salt, carbon dioxide and water. The hydroxide-containing compounds react to produce a neutral salt and water.

In Parts A, B and C of this lab, various neutralization reactions will be performed. Part A consists of formation and isolation of a simple salt, Part B includes an acid neutralization using an antacid tablet, and Part C includes the use of an indicator to form a rainbow in a tube.

In Part D of this lab, various acid–base titrations will be performed. A titration is a common quantitative analytical procedure for determining the concentration of an acid or base solution. For example, a base solution of known concentration may be added to an acid solution of unknown concentration until the neutralization point is reached. This neutralization point, called the equivalence point, is the point at which the number of moles of acid (H⁺) equals the number of moles of base (OH^–) in the solution. This point is called the endpoint and is often detected with an indicator that changes color at a pH of 7. In the titrations in parts D and E of this lab, phenolphthalein will be used as the indicator. Phenolphthalein is colorless in acid and pink in base. The point at which just enough base solution has been added to the acid solution to reach neutralization will be indicated by the first detection of a color change from colorless to a faint pink color.

In Part E of this lab, the concentration of two acid solutions of unknown concentration will be determined using titration. A base solution of known concentration will be added dropwise to a measured volume of acid solution of unknown concentration. At the neutralization point, the phenolphthalein indicator will turn a faint pink color. At this point, addition of base will be stopped and the volume of base will be recorded. From the balanced equation ratio, it can be determined how many moles of base are needed to neutralize each mole of acid. Then, since concentration (expressed in molarity, M) is equal to the number of moles per volume of solution, or M = moles/volume, the equation can be rearranged to solve for moles: Moles = molarity x volume = M x V. Since the molarity and volume of the base solution added is known, the moles of base at the equivalence point can be calculated. From the stoichiometric ratio of moles of acid to moles of base, the number of moles of acid can now be determined. Since the volume of acid is known, the unknown concentration of the acid may be determined using the equation M = moles acid/volume acid.

In the Part F extension to this lab, an antacid tablet will be tested for its “neutralizing power.” The goal will be to design a procedure to determine the average volume of hydrochloric acid that may be neutralized by one antacid tablet. The concepts of neutralization, titration and use of indicators will be applied as this problem is investigated.

Materials

Safety Precautions

All of the acids and bases used in this lab are very corrosive to eyes, skin and other body tissues. They are toxic by ingestion. Avoid all body tissue contact. Acetic acid and hydrochloric acid are also toxic by inhalation. Avoid breathing the vapors and dispense these chemicals in a fume hood. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron.

Procedure

Part A. Neutralization—Formation of a Salt

1. Add about 5 drops of each of the following to two individual wells of your microplate
   • 0.1 M HCl
   • 0.1 M NaOH
2. Using a different piece of dry pH paper for each solution, dip the end of the test paper into each solution. Remove the paper immediately and record the color of the pH paper in Table 1. Use the pH indicator color chart on the pH paper container to assign a numerical pH value to each solution. Record this value in Table 1.
3. Now to a different well of your microplate, add exactly 5 drops of both 0.1 M HCl and 0.1 M NaOH to the same well. (Note: Hold the pipet vertically for evenly sized drops). Use a toothpick to mix the two solutions.
4. Using a piece of dry pH paper, dip the end of the paper into the solution. Remove the paper immediately and record the color of the pH paper in Table 1. Use the pH indicator color chart on the pH paper container to assign a numerical pH value to the solution. Record this value in Table 1.
5. Repeat step 3, this time adding exactly 5 drops of both of the solutions to a plastic microscope slide. Use a toothpick to mix the two solutions. Allow the solution on the slide to evaporate overnight. This process may be accelerated by using a heating lamp.
6. When the solution has fully evaporated, examine the residue remaining on the slide. Use a hand lens or a microscope to aid in observation. Sketch and describe the appearance of the solid in Table 1.

Part B. Neutralization of Acid Using an Antacid Tablet

7. Place about 50 mL of distilled or deionized water in a small beaker. Add about 10 drops of 0.1 M HCl and 2–3 mL of universal indicator solution (~1 pipet full) to the water in the beaker. This solution represents an “upset” stomach.
8. Use a mortar and pestle to grind a white (non-pigmented) antacid tablet into a powder. If a mortar and pestle are not available, crush the tablet between two sheets of paper. Add the powdered antacid tablet to the solution in the beaker. Stir briefly and record observations of the solution in Table 2.
9. Continue stirring for about 1–2 minutes, recording the observations in Table 2. Note: The tablet is not fully soluble in water and the powder may float on top of the solution. This is not a factor for concern in this part of the lab.
10. After the solution has finished changing color, add a few drops of 0.1 M HCl to demonstrate how an antacid reacts as the stomach continues to produce acid. Record your observations in Table 2. The contents of the beaker may be rinsed down the drain with plenty of water.

Part C. Neutralization—Forming a Rainbow Reaction

11. In a clean 16 x 125 mm test tube, place about 2–3 mL (1 pipet full) of universal indicator solution.
12. Add about 10 mL of 0.1 M HCl to the test tube.
13. Fill a Beral-type pipet about half-way with saturated sodium carbonate, Na₂CO₃, solution. Tilt the test tube slightly and slowly squeeze the saturated sodium carbonate solution down the side of the test tube. Do not attempt to layer the solutions.
14. Hold the test tube vertical and steady or set it in a test tube rack. Colored layers will begin forming immediately. Observe these layers and record your observations in Table 3. The layers will last for a week if the tube is left undisturbed.
15. To dispose of the contents of the tube, pour the solution down the drain and rinse with plenty of water.

Part D. Neutralization—Microtitration of Monoprotic and Diprotic Acids

16. Add about 5 drops of each of the following to four individual wells of your microplate
    • 0.1 M HCl
    • 0.1 M HC₂H₃O₂
    • 0.1 M H₂SO₄
    • 0.1 M NaOH
17. Using a different piece of dry pH paper for each solution, dip the end of the paper into each solution. Remove the paper immediately and record the color of the pH paper in Table 4. Use the pH indicator color chart on the pH paper container to assign a numerical pH value to each solution. Record this value in Table 4.
18. To each solution, add 2 drops of phenolphthalein indicator solution. Record the color of each solution in Table 4. Note the color of the phenolphthalein in each solution.
19. Use your results from step 17 to label each solution as one of the following: strong acid, weak acid, strong base, weak base. Record your classification in Table 4.

Titration 1—HCl with NaOH

20. Now place exactly 10 drops of 0.1 M HCl in an individual well of your microplate. (Reminder: Hold the pipet vertically for evenly sized drops.) Add 3 drops of phenolphthalein indicator solution. Place your microplate on a sheet of white paper for better viewing of the color change.
21. Slowly, and with careful counting of the drops, add 0.1 M NaOH dropwise until the solution just turns and remains a faint pink color. Use a toothpick to mix the solutions. The neutralization point or endpoint is the point at which the faintest pink color remains without fading back to colorless. Record the exact number of drops of NaOH used to neutralize the HCl in Table 5.
22. Repeat steps 20–21. Again use 10 drops of 0.1 M HCl and titrate with 0.1 M NaOH. Record your results in Table 5. Compute the average number of drops of NaOH needed for neutralization. Record this value in Table 5.

Titration 2—H₂SO₄ with NaOH

23. Repeat steps 20–21, this time using 10 drops of 0.1 M H₂SO₄ and again titrating with 0.1 M NaOH. Conduct two trials and record your results in Table 5. Compute the average number of drops of NaOH needed to arrive at the neutralization point. Record this value in Table 5.

Titration 3—HC₂H₃O₂ with NaOH

24. Repeat steps 20–21, this time using 10 drops of 0.1 M HC₂H₃O₂ and again titrating with 0.1 M NaOH. Conduct two trials and record your results in Table 5. Compute the average number of drops of NaOH needed to arrive at the neutralization point. Record this value in Table 5.
25. Clean the well plate by pouring the solutions down the drain and rinsing the well plate with plenty of water. Caution: Take care when rinsing the plates so that the solution does not splash out.

Part E. Determine the Concentration of an Acid via Titration with NaOH

Titration 4—Unknown A (? M HCl) with NaOH

26. Place exactly 10 drops of Unknown A (? M HCl) in an individual well of your microplate. Add 3 drops of phenolphthalein indicator solution. Place your microplate on a sheet of white paper for better viewing of the color change.
27. Slowly, and with careful counting of the drops, add 0.1 M NaOH dropwise until the solution just turns and remains a faint pink color. The titration should be performed as outlined in step 21. Record the exact number of drops of NaOH used to neutralize the unknown concentration of acid in Table 5.
28. Repeat steps 26–27, again using Unknown A. Record your results in Table 5. Compute the average drops of NaOH. Record this value in Table 5.

Titration 5—Unknown B (? M H2SO4) with NaOH

29. Repeat steps 26–27, this time using Unknown B (? M H₂SO₄). Gather data for two or three trials and record your results in Table 5. Compute the average drops of NaOH. Record this value in Table 5. 

Student Worksheet PDF

12832_Student1.pdf