Materials Included In Kit

Additional Materials Required

Prelab Preparation

Buffer Solutions, pH 2–12: Use the buffer capsules to prepare standard acid–base solutions of known pH. Buffer capsules contain pre-weighed amounts of stable, dry powders that dissolve in distilled or deionized water to give solutions of known, constant pH. Dissolve each capsule in 100 mL of distilled or deionized water.

Label the unknowns with letter codes for student use. See also the Supplementary Information in the Further Extensions section for suggestions on using household substances as unknowns in Part B.

Safety Precautions

The standard acid and base solutions used in this experiment are body tissue irritants. Avoid contact of all chemicals with eyes and skin. Food-grade items that have been brought into the lab are considered laboratory chemicals and are for lab use only. Do not taste or ingest any materials in the lab and do not remove any remaining food items after they have been used in the lab. Wear chemical splash goggles and chemical-resistant gloves and apron. Remind students to wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling, and disposal information.

Disposal

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. All of the solutions may be flushed down the drain with excess water according to Flinn Suggested Disposal Method 26b.

Lab Hints

• This experiment is designed as a fun, inquiry-based study of indicators and their uses. Students generally enjoy any kind of “natural product” chemistry and this lab is no exception. Because this experiment works well and the materials are not hazardous, the students can be given a great deal of freedom in designing their laboratory procedure. This approach usually gives students more pride and ownership of work. The experiment can reasonably be completed in one 50-minute lab period.
• The inquiry approach may not be suitable for every classroom with every group of students. The Supplementary Information in the Further Extensions section contains an alternative, step-by-step procedure, along with prepared data tables. 
• To allow more time for measuring the indicator color changes and testing unknowns, consider setting up the extraction the day before lab. The natural indicators may be extracted overnight with water at room temperature. The amount of preparation time (to measure the dried flowers and place them in a beaker with water) is minimal—15 minutes should be sufficient. Cover the beaker with Parafilm^® or plastic wrap and allow the extraction to take place overnight. On the day of lab, filter or decant the mixture and use the resulting indicator solution in Parts A and B. Alternatively, extraction of the natural indicators with hot water may be completed as a homework assignment.
• This laboratory can also be extended by using other natural products as indicators. Flowers and fruits may be extracted with hot water or 70% isopropyl (rubbing) alcohol at room temperature. There is no difference in the color charts of natural indicator solutions obtained using these different solvents. Use isopropyl alcohol only in a well-ventilated lab setting. The recommended minimum extraction time for isopropyl alcohol is 15 minutes regardless of the solvent used.
• Although the extraction step may be the most enjoyable part of this lab for students, it does tend to get a little messy, especially if the experiment is performed by multiple lab sections in the same room on the same day. Teachers may find it convenient to prepare the indicators ahead of time.
• Additional unknowns can also be used. Dilute concentrated acids (0.001 M) work great for a pH 3 unknown. Zinc nitrate (pH 6), ammonium nitrate (pH 6–7), and ammonia solutions (pH 10–12) work well. Household products also work well. Consider holding a “Measurement Fair” to have students analyze the acid–base properties of household substances using their natural indicators. The Supplementary Information in the Further Extensions section contains a list of household substances to consider as unknowns for Part B.
• A wide range of fruits and flowers contain natural acid–base indicators. The table on the following page summarizes the information obtained from a brief literature survey. The list is not meant to be exhaustive, but rather to demonstrate the variety of options suitable for classroom study.
  {13890_Hints_Table_4}

Teacher Tips

• This experiment reinforces key concepts and definitions in the chemistry of acids and bases. The behavior of natural indicators illustrates the definition of Brønsted acids (proton donors). Most natural indicators are further classified as weak acids (dissociate only partially in water and their reactions with water are reversible). The different colors observed for natural indicators thus reflect the position of equilibrium under different conditions. The color transitions are examples of Le Chatelier’s Principle in action. Use Equation 1 to predict the direction the indicator equilibrium will be shifted as a result of increasing or decreasing the H₃O⁺ concentration. The Post-Lab Questions provide an opportunity to review the relationship between pH and [H₃O⁺] in pH calculations.
• A picture is worth a thousand words! Have students draw color charts of indicator color changes using colored pencils.

Further Extensions

Supplementary Information

Alternative Procedure

1. Obtain 2–3 g of dried flowers or plant material.
2. Tear, chop, grind or crush the material and place the pieces in a 150-mL beaker.
3. Cover the sample with water. Use a minimum amount of water (approximately 50 mL) so that the resulting indicator solution will be as concentrated as possible.
4. Heat the mixture to just below the boiling point using a hot plate or Bunsen burner setup.
5. After 15 minutes, decant or filter the mixture into a clean, 100-mL beaker. The natural indicator extract should be clear, not cloudy.

Part A. Indicator Color Changes

6. Using a pipet, add 20 drops (1 mL) of each standard acid and base solution (pH 2–12) to separate wells on a 24-well reaction plate. Note the location of each solution.
7. Add 5 drops of the natural indicator solution to each well.
8. Record the color of the indicator in each “standard” well in Data Table A.
9. Do NOT discard the solutions in the standard wells until steps 10–19 have been completed.

Part B. Classifying Unknown Solutions

10. Using a pipet, add 20 drops (about 1 mL) of each unknown solution to be tested to a separate well on a 24-well reaction plate. Note the location of each solution.
11. Add 5 drops of the natural indicator to each well.
12. Record the color of the indicator in each “unknown” well in Data Table B.
13. Compare the color of each unknown to the colors of the indicator in the standard wells. Determine the pH of the standard solution that most closely matches the color of each unknown. Note: If the color of the indicator is the same in two or more standard wells, for example pH 2, 3 and 4, then find the pH range that most closely matches the color of the unknown.
14. Record the approximate pH value or pH range for each unknown in Data Table B.

Measure the pH of the Unknowns Using Synthetic Indicators

15. Based on the approximate pH value or pH range for each unknown (step 14), choose at least one indicator from the following table that will allow you either to narrow down the pH range or to confirm the pH value. Example: If the approximate pH range of an unknown is less than 4, choose thymol blue to narrow down the pH range.
    {13890_Procedure_Table_2}
16. Record the indicator selected for each unknown in Data Table B.
17. Add 20 drops (1 mL) of each unknown to separate wells on the reaction plate.
18. Add 2 drops of the appropriate indicator to the test well of each unknown.
19. Record the color of each solution in Data Table B.
20. The contents of the reaction plate may be rinsed down the drain with excess water.

Data Table A. Indicator Color Changes Data Table B. Classifying Unknown Solutions Analysis of Household Substances
The following household substances provide convenient unknowns for classifying acid–base solutions using indicators.
Alka-Seltzer^®
Ammonia, household
Antacid tablet
Aspirin tablet
Baking powder
Baking soda
Bleach
Club soda
Cola
Contact lens solution
Cream of tartar
Drain cleaning solution
Fruit Fresh^®
Ginger ale
Grapefruit juice
Hair spray
Hand lotion
Laundry detergent
Lemon-lime soda
Lemon juice
Milk
Mouthwash
Shampoo
Tea
Toothpaste
Vinegar
Vitamin C tablet
Windex^®

Answers to Prelab Questions

Phenolphthalein is a synthetic indicator that is colorless when the pH is 10. The pH range from 8–10 is the “transition range” for phenolphthalein. When phenolphthalein is added to solutions having a pH between 8 and 10, the indicator is intermediate in color between colorless and red, or various shades of pink. The color changes for phenolphthalein and two other indicators are summarized in the following color charts (see Table 1). Areas shaded with hash marks indicate pH intervals in which the color of the indicator changes from one form to another. Note: Alizarin exhibits two different color transitions, between 5–7 and 11–13, respectively. At pH values greater than 12 alizarin appears violet.

1. What will be the intermediate color of bromthymol blue in a solution of pH 7?

   The color should be intermediate between yellow and blue—green.

2. A colorless household solution was tested with the three indicators shown. The solution was colorless with phenolphthalein, yellow with bromthymol blue, and orange with alizarin. What is the pH of the solution? Be as specific as possible.

   The pH ranges suggested by the indicator colors are pH <8 (phenolphthalein), pH <6 (bromthymol blue) and pH = 5–7 (alizarin). Combining these ranges gives a pH value between 5 and 6. The pH may be reported as the middle of this interval, with a “plus-or-minus” value to indicate the range: pH = 5.5 ±0.5.

Sample Data

Sample Data Table A. Indicator Color Changes¹

¹ Colors may vary due to concentration of the indicator solution.
² Fades to greenish-brown.

Sample Data Table B. Classifying Unknown Solutions

*Do not reveal the identity of the unknowns to the students.

Sample Results Table

Answers to Questions

1. Assume that the pH 2 color of the natural indicator represents its most acidic form (HIn).
   1. What is the pH range in which the most acidic form predominates?

      For the red rose indicator as an example, the most acidic form of the indicator is pink in color. This acidic form of the indicator appears to be the predominant form of the indicator in solution at pH values less than or equal to 4.

   2. Calculate the lowest H₃O⁺ concentration at which the indicator still exists in this form.

      The lowest H₃O⁺ concentration at which the acidic form predominates is therefore 1 x 10^–4 M. Note: At pH 5–6, the solution is still pink, but it becomes very pale in color, and finally almost colorless at pH 7. The very pale pink color thus appears to be a “transition” or intermediate color.

2. Assume that the pH 12 color of the natural indicator represents its most basic form (In^–).
   1. What is the pH range in which the most basic form predominates?

      For the red rose indicator as an example, the most basic form of the indicator is yellow in color. This basic form of the indicator appears to be the predominant form of the indicator in solution at pH values greater than or equal to 10.

   2. Calculate the highest H₃O⁺ concentration at which the indicator still exists in this form.

      The highest H₃O⁺ concentration at which the basic form predominates is therefore 1 x 10^–10 M.

3. For one of the unknown acid–base solutions that you tested, explain why you chose the combination of indicators you did to determine the pH value of the solution. What is the advantage of using multiple indicators, rather than a single indicator, to determine the pH of a substance?

   In the case of Unknown C (sodium bicarbonate), the natural red rose indicator gave a pH estimate of 8–9. Bromthymol blue was selected as an alternate indicator to confirm or narrow the pH range. Bromthymol blue was green, suggesting that the pH was in the 7–8 range. The overlap of these two pH estimates suggests that the pH of the solution is very close to 8. Using multiple indicators, rather than a single indicator, often makes it possible to obtain a more precise (narrow) estimate of the pH of a substance.

4. Construct a results table to summarize the properties of the unknowns.
   1. Estimate the pH value of each unknown.
   2. Classify each solution as acidic or basic.
   3. Within each class of unknowns—acids and bases—arrange the solutions in order from least acidic to most acidic and least basic to most basic, respectively.

References

This activity is from Flinn ChemTopic^™ Labs, Volume 13, Acids and Bases; Cesa, I., Ed; Flinn Scientific: Batavia, IL, 2002.

Introduction

Roses are red, violets are blue—or are they? Red roses, as well as many other flowers and fruits, contain natural indicators that are sensitive to acids and bases. The color of a natural acid–base indicator depends on pH. One of the most well known effects of natural indicators in plants occurs in the hydrangea or snowball plant. Hydrangea flowers are blue when grown in acidic soils and pink or red in basic soils. How do the colors of natural indicators vary with pH?

Concepts

• Indicators
• Extraction
• Weak acid vs. conjugate base
• pH scale

Background

Indicators are dyes or pigments that are isolated from a variety of sources, including plants, fungi, and algae. For example, almost any flower that is red, blue, or purple in color contains a class of organic pigments called anthocyanins that change color with pH. The use of natural dyes as acid–base indicators was first reported in 1664 by Sir Robert Boyle in his collection of essays Experimental History of Colours. Indeed, Boyle made an important contribution to the early theory of acids and bases by using indicators for the classification of these substances. The idea, however, may actually have originated much earlier—medieval painters used natural dyes treated with vinegar and limewater to make different color watercolor paints.

Acid–base indicators are large organic molecules that behave as weak acids—they can donate hydrogen ions to water molecules to form their conjugate bases (Equation 1). The distinguishing characteristic of indicators is that the acid (HIn) and conjugate base (In^–) are different colors.

The abbreviation HIn represents an uncharged indicator molecule, and In– an indicator ion after it has lost a hydrogen ion. The color changes of acid–base indicators illustrate an application of reversible reactions and equilibrium. Because indicators are weak acids, the reactions summarized in Equation 1 are reversible. Reversible reactions are easily forced to go in either direction, depending on reaction conditions. The actual color of an indicator solution thus reflects the position of equilibrium for Equation 1 and depends on the concentration of H₃O⁺ ions (and hence the pH) of the solution.

There are three possible cases. (1) Most of the indicator molecules exist in the form HIn and the color of the solution is essentially the color of HIn. (2) Most of the indicator molecules exist in the form In^– and the color of the solution is essentially the color of In^–. (3) The solution contains roughly equal amounts of the two forms and the resulting color is intermediate between that of HIn and In–. The exact concentrations of H₃O⁺ at which cases 1–3 will predominate depend on the structure of the indicator and the equilibrium constant for Equation 1. Different indicators change color in different pH ranges.

Natural indicator solutions are obtained by treating flowers and fruits with a solvent to dissolve the soluble components. This process, called extraction, is similar to the procedure used to make a cup of tea using a tea bag. The solid is crushed or ground and extracted with an appropriate solvent, such as boiling water, ethyl alcohol or rubbing alcohol.

The color of an acid–base indicator depends on the concentration of H₃O⁺ ions, which is most conveniently expressed using the pH scale. The mathematical relationship between pH and [H₃O⁺] is given in Equation 2. The H₃O⁺ concentration in water ranges from 1 M in 1 M hydrochloric acid to 10^–14 M in 1 M sodium hydroxide. In pure water, which is neutral (neither acidic nor basic), the H₃O⁺ concentration is equal to 10^–7 M. The logarithm of the concentration is the “power of ten” exponent in these concentration terms. Thus, the negative logarithms (Equation 2) of typical H₃O⁺ concentrations are positive numbers from 0–14. The pH scale ranges from 0–14, with 7 being neutral. Acids have pH values less than 7, while bases have pH values greater than 7.

Within the pH range of acid solutions, either a more concentrated or a strong acid solution will have a lower pH than a less concentrated or a weak acid solution, respectively. Thus, the pH values of 0.1 and 0.01 M HCl solutions are 1 and 2, respectively, while the pH of 0.1 M acetic acid (a weak acid) is about 3. On the basic side of the pH scale, either a more concentrated or strong base solution will have a higher pH than a less concentrated or a weak base solution, respectively. Thus, the pH values of 0.1 and 0.01 M NaOH solutions are 13 and 12, respectively, while the pH of 0.1 M ammonia (a weak base) is about 11. Remember that the pH scale is logarithmic—a solution of pH 3 is ten times more acidic than a solution of pH 4, and 100 times more acidic than a solution of pH 5. Figure 1 summarizes the pH scale and the pH range of acids and bases.

Experiment Overview

The purpose of this experiment is to extract natural indicators and design a procedure to investigate their color changes as a function of pH. A set of standard acid and base solutions of known pH (pH = 2–12) will be provided. The results will be used to construct color charts of the indicators. In Part B, the natural indicators will be used, along with other known indicator solutions, to analyze the pH values of unknown solutions.

Materials

Prelab Questions

Phenolphthalein is a synthetic indicator that is colorless when the pH is 10. The pH range from 8–10 is the “transition range” for phenolphthalein. When phenolphthalein is added to solutions having a pH between 8 and 10, the indicator is intermediate in color between colorless and red, or various shades of pink. The color changes for phenolphthalein and two other indicators are summarized in the following color charts (see Table 1). Areas shaded with hash marks indicate pH intervals in which the color of the indicator changes from one form to another. Note: Alizarin exhibits two different color transitions, between 5–7 and 11–13, respectively. At pH values greater than 12 alizarin appears violet.

1. What will be the intermediate color of bromthymol blue in a solution of pH 7?
2. A colorless household solution was tested with the three indicators shown. The solution was colorless with phenolphthalein, yellow with bromthymol blue, and orange with alizarin. What is the pH of the solution? Be as specific as possible.

Safety Precautions

The standard acid and base solutions used in this experiment are body tissue irritants. Avoid contact of all chemicals with eyes and skin. Food-grade items that have been brought into the lab are considered laboratory chemicals and are for lab use only. Do not taste or ingest any materials in the lab and do not remove any remaining food items after they have been used in the lab. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Preparation: Extraction of Natural Indicators

1. For rose petals and hibiscus:
1. Obtain 2–3 g of dried flowers and place them in a 150-mL beaker with 50–60 mL of water.
2. Place the beaker on a hot plate or Bunsen burner setup and heat the flower/water mixture for 5–10 minutes.
3. Cool the mixture and decant or filter the liquid into a clean, 100-mL beaker. The natural indicator solution should be strongly colored but clear. If necessary, filter the mixture to obtain a clear solution.
2. For teas: Place one tea bag in a 150-mL beaker, add 50-mL water, and heat to just below the boiling point for 10–15 minutes. Cool and decant the liquid into a clean, 100-mL beaker.
3. For juices: Juices may be used directly “as is,” with no pretreatment, in Parts A and B. 

Part A. Indicator Color Changes

4. Design a procedure using the standard acid and base solutions of known pH to determine the color changes for the natural indicator solution and the pH intervals in which the color changes occur.
5. Construct a data table to record the results.
6. Show the data table and discuss the proposed procedure with your instructor.
7. Carry out the procedure and record the results.

Part B. Classifying Unknown Solutions

8. Design a procedure using your natural indicator solution and at least one synthetic indicator to determine the pH values of unknown solutions. Hint: Choose indicators that will give you the narrowest range possible for the pH value of each unknown. The color charts for the available indicators are shown in Table 2.
   {13890_Procedure_Table_2}
9. Construct a data table to record the results.
10. Show your data table and discuss the proposed procedure with your instructor.
11. Carry out the procedure and record the results.

Student Worksheet PDF

13890_Student1.pdf