Teacher Notes

Manipulating Equilibria to Determine Sulfites in Grape Juice

Student Laboratory Kit

Materials Included In Kit

Crystal violet solution, 1% alcoholic
Hydrochloric acid solution, HCl, 1 M, 500 mL
Phosphoric acid, H3PO4, 25 mL
Sodium hydroxide solution, NaOH, 1 M, 500 mL
Strontium chloride solution, SrCl2, 10%, 10 mL

Additional Materials Required

Water, distilled or deionized
White grape juice, 100 mL
Balance, 0.01-g precision
Beaker, 25-mL
Beakers, 100-mL, 2
Beaker, 250-mL
Erlenmeyer flask, 250-mL
Erlenmeyer flask, 500-mL
Filter funnel
Filter paper, quantitative
Graduated cylinder, 10-mL
Graduated cylinder, 100-mL
Hot plate
Ice
Litmus paper
Pipets, Beral-type or Pasteur
Test tubes, 2
Watch glass

Prelab Preparation

Add 100 mL of distilled water to a 500 mL Erlenmeyer flask, followed by 15 mL of phosphoric acid. Add more distilled water until there is a total of 250 mL of solution. Add a few drops crystal violet solution to the dilute phosphoric acid solution. The best experimental results are obtained with a solution that is teal or green in color; if the solution is too blue, add more phosphoric acid, if too yellow add more water. Because crystal violet fades over time, this is best prepared fresh.

Place 50 g of strontium chloride ina 1-L Erlenmeyer flask. Add 500 mL of distilled water and mix until fully dissolved.

Safety Precautions

Hydrochloric acid solution is a corrosive liquid and is toxic by ingestion and inhalation. Sodium hydroxide solution is corrosive to body tissue, especially eyes and skin. Phosphoric acid solution is a corrosive liquid and is toxic by ingestion and inhalation. Crystal violet is a strong dye and can stain skin and clothing. Strontium chloride is an irritant and damaging to eyes. Avoid contact of all chemicals with skin and eyes. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron or laboratory coat. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information. Please follow all laboratory safety guidelines.

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. Excess hydrochloric acid may be neutralized with base and then poured down the drain with excess water according to Flinn Suggested Disposal Method #24b. Excess phosphoric acid may be neutralized with base and then poured down the drain with excess water according to Flinn Suggested Disposal Method #24b. Sodium hydroxide may be neutralized with acid and then poured down the drain with an excess water according to Flinn Suggested Disposal Method #10. Strontium chloride may be disposed of according to Flinn Suggested Disposal Method #26a.

Lab Hints

  • The amount of sulfite in white grape juice can vary between brands. When purchasing juice, check that the bottle is labeled as containing sulfites.
  • White wine can be used instead of grape juice.
  • The final color of the acidified indicator solution will vary depending on the exact composition of the initial solution.

Further Extensions

Online Educational Resources

Simulation of different strength acids and bases
https://phet.colorado.edu/en/simulation/acid-base-solutions

Reversible reaction simulation
https://phet.colorado.edu/en/simulation/legacy/reversible-reactions

Correlation to Next Generation Science Standards (NGSS)†

Science & Engineering Practices

Developing and using models
Planning and carrying out investigations
Constructing explanations and designing solutions
Obtaining, evaluation, and communicating information

Disciplinary Core Ideas

HS-PS1.A: Structure and Properties of Matter
HS-PS1.B: Chemical Reactions

Crosscutting Concepts

Patterns
Structure and function

Answers to Questions

  1. What happened to the grape juice as you added the sodium hydroxide?

    As the sodium hydroxide solution was added, the pH of the juice increased and the color darkened.

  2. What happened to the grape juice when you added the strontium chloride?

    The solution became cloudy as strontium sulfite precipitated out.

  3. What was the mass of the clean dry filter paper?

    1.15 g

  4. What color was the acidified indicator solution before being placed in the hot/cold water baths?

    The acidified solution was teal to begin with.

  5. What happened to the color of the acidified indicator solution in the ice bath?

    The solution became green.

  6. What happened to the color of the acidified indicator solution in the hot water bath?

    The solution became blue.

  7. Is the deprotonation of phosphoric acid endothermic or exothermic?

    Exothermic.

  8. What was the mass of filter paper with strontium sulfite on it?

    2.51 g

  9. What was the mass of strontium sulfite?

    1.36 g

  10. What mass of sulfite was in the grape juice?

    0.65 g

  11. What happened to the grape juice as you added the hydrochloric acid?

    The strontium sulfite dissolved back into the solution, and the juice took on its original appearance.

  12. Why did you originally need to add sodium hydroxide to the grape juice before adding strontium chloride?

    The sulfites in the juice exist in several forms, with only the deprotonated species able to form strontium sulfite. Adding sodium hydroxide shifts the equilibria within the juice so the deprotonated form is the dominant species, maximizing the amount of precipitate.

Recommended Products

Item No. Description
AP6690 Total Acidity—Titration of Fruit Juices—Student Laboratory Kit
AP3104 Filter Paper, Qualitative, 12.5 cm

Student Pages

Manipulating Equilibria to Determine Sulfites in Grape Juice

Introduction

Sulfites are commonly added to foodstuffs to prevent oxidation and microbial spoilage. However, a small percentage of the population suffer from sulfite sensitivity. Determining the concentration of sulfite ions can be complicated due to the presence of several competing equilibria within the solution. By utilizing Le Chatelier’s principle we can exploit these equilibria to permit a qualitative determination.

Watch the introductory video.

Concepts

  • Equilibrium
  • Reaction rates
  • Solubility
  • Acid–base reactions
  • Le Chatelier’s principle

Background

Most chemical reactions do not simply proceed 100% in the forward direction but rather are reversible, meaning they can go in both ways. When the forward rate and the reverse rate are equal, the system is at equilibrium. In a closed system, any reversible reaction will eventually reach a point at which the amounts of reactant and products no longer change. It is important to understand that reactions are still occurring, even though no observable changes can be measured.

We can represent this idea mathematically through the use of the equilibrium constant. Consider the following general equation for a reversible chemical reaction:

{14162_Background_Equation_1}

The equilibrium constant, Keq, for this general reaction is given by Equation 2, where the square brackets refer to the molar concentrations of the reactants and products at equilibrium. Pure liquids (including solvents) and solids have a value of 1 when used in an equilibrium equation and should be omitted as they will have no effect on the final value.

{14162_Background_Equation_2}

The equilibrium constant gets its name from the fact that for any reversible chemical reaction, the value of Keq is constant at a particular temperature, but the concentrations of reactants and products at equilibrium vary, depending on the initial amounts of materials present. The ratio of reactants and products described by Keq is always the same, as long as the system has reached equilibrium and the temperature does not change.

Any change that is made to a system at equilibrium may be considered a stress—this includes adding or removing reagents, changing the temperature, or for a reaction involving gases, changing the pressure. The rates of the forward and reverse reactions will change as a result until equilibrium is again established. Henry Le Chatelier published many studies of equilibrium systems. Le Chatelier’s principle predicts how equilibrium can be restored:

“If an equilibrium system is subjected to a stress, the system will react in such a way as to reduce the stress.”

Le Chatelier’s principle is a qualitative approach to predicting and interpreting shifts in equilibrium systems. A quantitative approach utilizes the Keq of the reaction and the reaction quotient, Q. The reaction quotient is a snapshot of the concentrations of reactants and products at a particular time. Q is calculated using the same formula as Keq (Equation 2). Depending on the instantaneous concentrations of reactants and products, Q and Keq may differ or be the same. If Q and Keq differ, the system is not at equilibrium and the rates of the forward and reverse reactions will change until Q = Keq. If Q is greater than K then the reaction will shift to the left, and if Q is less than K then the reaction will shift to the right.

Watch the video.

The sulfites in grape juice are taking part in two different equilibrium reactions:

{14162_Background_Equation_3}
{14162_Background_Equation_4}

Because fruit juices are acidic, the concentration of H+ in the solution is high and the equilibria are pushed towards the left. To analyze the amount of sulfite in the juice, we are going to precipitate it with strontium, according to the following reaction:

{14162_Background_Equation_5}

In order to maximize the amount of strontium sulfite that precipitates, we will need to shift both the reactions in Equation 3 and Equation 4 to the right. This can be achieved through the addition of sodium hydroxide. Equation 5 shows an equilibrium relating to the precipitation of a solid, and as such the equilibrium constant is instead referred to as a solubility product, is given the symbol Ksp and takes the following form:

{14162_Background_Equation_6}

You will notice that the solid strontium sulfite doesn’t appear in Equation 6. The Ksp value for strontium sulfite is 4 x 10−8, a precipitate will form if the concentration of strontium ions multiplied by the concentration of sulfite ions is greater than Ksp, so only a small excess of strontium is needed to precipitate almost all the sulfite ions. When adding base to the grape juice, the pH needs to be monitored to avoid the formation of strontium hydroxide, which has a Ksp value of 1.5 x 10−4. At a pH of 10, the concentration of hydroxide ions in solution is 1 x 10−4 M, and the concentration of strontium needed for precipitation to occur is 1.5 x 104 M. For comparison, at a pH of 14, the concentration of hydroxide ions in solution is 1 M and the concentration of strontium needed for precipitation to occur is 1.5 x 10−4 M.

Experiment Overview

This laboratory takes place over two separate lab periods. In the first laboratory session, you will precipitate and isolate strontium sulfite from white grape juice as well as examine the effect of temperature on the deprotonation of phosphoric acid. The second session involves determining the concentration of sulfite in the original sample, followed by redissolving the sulfite and restoring the juice’s original color.

Materials

Acidified crystal violet solution, 10 mL
Hydrochloric acid solution, HCl, 1 M, 25 mL
Sodium hydroxide solution, NaOH, 1 M, 25 mL
Strontium chloride solution, SrCl2, 10%, 10 mL
Water, distilled or deionized
White grape juice, 100 mL
Balance, 0.01-g precision
Beaker, 25 mL
Beakers, 100-mL, 2
Beaker, 250-mL
Erlenmeyer flask, 250-mL
Filter funnel
Filter paper, quantitative
Graduated cylinder, 10-mL
Graduated cylinder, 100-mL
Hot plate
Ice
Litmus paper
Pipet, Beral-type or Pasteur
Test tubes, 2
Watch glass

Prelab Questions

Watch the video, and complete the online activities.

Select the correct equilibrium constant for the following reactions.

{14162_PreLab_Table_1}

Based on the given Ksp predict if a precipitate will form from a 0.0095 M solution of each of the following complexes:

{14162_PreLab_Table_2}

Safety Precautions

Hydrochloric acid solution is a corrosive liquid and is toxic by ingestion and inhalation. Sodium hydroxide solution is corrosive to body tissue, especially eyes and skin. Phosphoric acid solution is a corrosive liquid and is toxic by ingestion and inhalation. Crystal violet is a strong dye and can stain skin and clothing. Strontium chloride is an irritant and damaging to eyes. Avoid contact of all chemicals with skin and eyes. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron or laboratory coat. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information. Please follow all laboratory safety guidelines.

Procedure

Part 1. Precipitation of Strontium Sulfite

  1. Pour 100 ml grape juice into a 250 mL beaker.
  2. Using a pipet, slowly add 1 M NaOH until the pH of the solution is around 9–10.
  3. Add 10 mL of the 10% strontium chloride solution to the Erlenmeyer flask.
  4. Swirl the Erlenmeyer flask and then set it aside for at least 10 minutes while the strontium sulfate precipitates from solution.
  5. Weigh a clean, dry piece of filter paper and record its mass.
  6. Gravity filter the solution. Due to the time this takes, it is a good idea complete Part 2 now before finishing Part 1.
  7. Cover and set aside the filtered juice.
  8. Wash the precipitate with 10 mL of distilled water.
  9. Set aside the filter paper to dry until the next lab session.

Part 2. Effect of Temperature on Acid Equilibria
In the Background section, we only discussed how changes in concentration affect the equilibrium; however, temperature can also have an effect. For exothermic reactions, heat can be thought of as a product, and heating an exothermic equilibrium reaction will cause it to shift to the left. Conversely, for endothermic reactions, it can be thought of as a reactant, and heating an endothermic equilibrium reaction will cause it to shift to the right.

Phosphoric acid solutions involve the following three equilibrium reactions:

{14162_Procedure_Equation_7}

The more H+ ions in solution, the lower the pH. The phosphoric acid solution you have contains crystal violet, a pH indicator that is yellow when the pH is negative and blue/violet when above 2.

  1. Obtain two 100 mL beakers.
  2. Half fill one of the beakers with approximately 50 mL of ice water, and the other with room-temperature water.
  3. Place the room-temperature beaker on a hot plate and heat it until almost boiling.
  4. Put 5 mL of the supplied acidified indicator solution into each of the two test tubes. Record the initial color of the solution.
  5. Place one of the test tubes into the ice water and the other into the hot water.
  6. Over the next 10 minutes, record any color changes in the solutions.

Part 3. Calculating the Amount of Sulfite in Juice

  1. Weigh the dry filter paper and strontium sulfite.
  2. Calculate the mass of strontium sulfite.
  3. Strontium makes up 52.3% of strontium sulfite by mass. From this piece of information, calculate the mass of sulfite ions.

Part 4. Restoring the Grape Juice

  1. Add the strontium chloride to the filtered grape juice.
  2. Slowly add hydrochloric acid to the grape juice until the pH is in the range of 3–4.
  3. Record all observations during the addition of the acid.
  4. Consult your instructor for appropriate disposal procedures.

Questions

  1. What happened to the grape juice as you added the sodium hydroxide?
  2. What happened to the grape juice when you added the strontium chloride?
  3. What was the mass of the clean dry filter paper?
  4. What color was the acidified indicator solution before being placed in the hot/cold water baths?
  5. What happened to the color of the acidified indicator solution in the ice bath?
  6. What happened to the color of the acidified indicator solution in the hot water bath?
  7. Is the deprotonation of phosphoric acid endothermic or exothermic?
  8. What was the mass of filter paper with strontium sulfite on it?
  9. What was the mass of strontium sulfite?
  10. What mass of sulfite was in the grape juice?
  11. What happened to the grape juice as you added the hydrochloric acid?
  12. Why did you originally need to add sodium hydroxide to the grape juice before adding strontium chloride?

†Next Generation Science Standards and NGSS are registered trademarks of Achieve. Neither Achieve nor the lead states and partners that developed the Next Generation Science Standards were involved in the production of this product, and do not endorse it.