Materials Included In Kit

Additional Materials Required

Prelab Preparation

Standardization of Sodium Hydroxide Solution

1. Students may standardize their own sodium hydroxide solution or the instructor may standardize ahead of time.
2. Obtain a sample of potassium hydrogen phthalate (KHP) that has been previously dried in an oven at 110 °C for two hours and stored in a desiccator.
3. On an analytical balance, accurately weigh 0.4 to 0.6 g of KHP.
4. Transfer the KHP into an Erlenmeyer flask or beaker.
5. Add about 40 mL of distilled water to the flask and swirl until all the KHP is dissolved.
6. Obtain 0.10 M sodium hydroxide, NaOH, solution. This is the nominal concentration.
7. Rinse and fill the buret with the NaOH solution.
8. Add three drops of phenolphthalein solution to the KHP solution in the flask.
9. Titrate the KHP solution using the nominal 0.1 M NaOH solution. Three titrations should be conducted and averaged.
10. Calculate the concentration of NaOH. See the following sample calculation.

    Sample calculation: Assume 0.500 g of KHP was used, and the average volume of NaOH titrant required was 25.50 mL. KHP, KHC8H4O4, molar mass = 204.23 g/mole. The mole ratio for the neutralization reaction is one mole per KHP mole NaOH.

Safety Precautions

Dilute sodium hydroxide and acid solutions are irritating to skin and eyes. Avoid contact of all chemicals with eyes and skin. All 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 chemistry laboratory. Do not remove any remaining food items from the lab after they have been used in the lab. Wear chemical splash goggles, chemical-resistant gloves and a chemical resistant apron. Wash hands thoroughly with soap and water before leaving the lab. Please review 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. Excess acetic acid and hydrochloric acid solutions may be neutralized according to Flinn Suggested Methods #24a and #24b, respectively. Excess sodium hydroxide solution may be neutralized according to Flinn Suggested Disposal Method #10. The phenolphthalein solution may be saved for future use. The titrated solutions and leftover juices may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Enough materials are provided in this kit for 24 students working in pairs. Student groups may test different juices. Data may be compiled and discussed as a class. The kit includes three juice samples, however students may also choose to bring light-colored sodas to test.

• This laboratory activity can be completed in two 50-minute class periods. It is important to allow time between the Introductory Activity and the Guided-Inquiry Activity for students to discuss and design the guided-inquiry procedures. Also, all student-designed procedures must be approved for safety before students are allowed to implement them in the lab. Prelab Questions may be completed before lab begins the first day.

•  

  Three different indicators are included for the baseline activity inquiry section. Allow students to determine the best indicator to use in their titration. Their choice should be phenolphthalein.
• Burets should be rinsed with the standardized sodium hydroxide solution three times, about 5–7 mL each time, before beginning titrations. Students may standardize their own sodium hydroxide solution or the instructor may standardize ahead of time.
• Prior to beginning the titration, students may measure 5-mL of juice and conduct a rough titration with a graduated pipet and phenolphthalein indicator to determine the endpoint of the solution and color. For example, if 3-mL of sodium hydroxide are needed to reach the endpoint for 5-mL of juice, then 12-mL of sodium hydroxide would be needed to reach the endpoint in 20-mL of juice.
• Students will need a 20-mL sample of juice or soda for titration. They will dilute with 30-mL of distilled or deionized water. Analyte solution will be a total of 50 mL.
• Student groups will collect pH readings during titrations, every 1 mL added, using a pH meter. The endpoint will also be detected using an indicator. Phenolphthalein is provided in the kit.
• Students will plot the volume of sodium hydroxide added versus pH readings to construct a calibration curve of the titration.
• Students will determine the amount of acid in samples of fruit juices. They may report data in grams. See the Sample Data section.
• If pH meters are not available, pH paper may be used. A pH chart will be needed in order to plot data.
• Students may swirl the beaker or flask during the titration procedure or use a magnetic stirrer with a stir bar.

Further Extensions

Alignment to the Curriculum Framework for AP^® Chemistry

Enduring Understandings and Essential Knowledge

Atoms are conserved in physical and chemical processes. (1E)
1E2: Conservation of atoms makes it possible to compute the masses of substances involved in physical and chemical processes. Chemical processes result in the formation of new substances, and the amount of these depends on the number and the types and masses of elements in the reactants, as well as the efficiency of the transformation.

Chemical changes are represented by a balanced chemical equation that identifies the ratios with which reactants react and products form. (3A)
3A2: Quantitative information can be derived from stoichiometric calculations that utilize the mole ratios from the balanced chemical equations. The role of stoichiometry in real-world applications is important to note, so that it does not seem to be simply an exercise done only by chemists.

Learning Objectives
1.20 The student can design, and/or interpret data from, an experiment that uses titration to determine the concentration of an analyte in a solution.
3.3 The student is able to use stoichiometric calculations to predict the results of performing a reaction in the laboratory and/or to analyze deviations from the expected results.

Science Practices
1.1 The student can create representations and models of natural or man-made phenomena and systems in the domain.
2.2 The student can apply mathematical routines to quantities that describe natural phenomena.
3.1 The student can pose scientific questions.
4.2 The student can design a plan for collecting data to answer a particular scientific question.
5.1 The student can analyze data to identify patterns or relationships.
6.4 The student can make claims and predictions about natural phenomena based on scientific theories and models.
7.1 The student can connect phenomena and models across spatial and temporal scales.

Answers to Prelab Questions

1. Using the structural formula of citric acid shown in Figure 1, determine the molecular formula of citric acid and calculate its molar mass (g/mole).

   The molecular formula of citric acid is C₆H₈O₇ and its molar mass = 192.0 g/mole.

2. A 10.0-mL sample of pineapple juice was titrated with 0.100 M sodium hydroxide solution. The average volume of NaOH required to reach the endpoint was 12.8 mL.
   1. Calculate the number of moles of sodium hydroxide required to reach the endpoint.

      Moles of NaOH = Molarity (moles/L) x Volume (L)

      {13762_PreLab_Equation_1}
   2. Using the mole ratio for the neutralization reaction shown in Equation 1, determine the number of moles of citric acid in 10.0 mL of pineapple juice.
      {13762_PreLab_Equation_2}
      {13762_PreLab_Equation_3}

      = 4.27 x 10^–4 moles citric acid

   3. Multiply the number of moles of citric acid by its molar mass to calculate the mass of citric acid in 10.0 mL of the juice.

      (4.27 x 10^–4 moles citric acid)(192.0 g/mole) = 0.0819 g

   4. The concentration of acid in juices is usually expressed in grams of acid per 100 mL of juice. What is the concentration of citric acid in pineapple juice?
      {13762_PreLab_Equation_4}
3. Write a balanced chemical equation for the neutralization reaction of (a) hydrochloric acid and (b) acetic acid with sodium hydroxide.
   1. HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
   2. CH₃CO₂H(aq) + NaOH(aq) → NaCH₃CO₂(aq) + H₂O(l)
4. The titration curves for hydrochloric acid and acetic acid with sodium hydroxide are shown below. Distinguish between the strong and weak acid in terms of the initial pH, the pH at the equivalence point and the overall shape of the titration curve.

   The initial pH of a strong acid will be lower and more acidic for a strong acid, assuming they both have the same concentration (0.1 M in this example). The titration curve for a strong acid is almost flat in the region before the equivalence point. The pH then rises very suddenly or steeply, giving an almost vertical line right near the equivalence point, which occurs at pH = 7 for a strong acid. In the case of a weak acid, the rise is not quite as steep and the pH at the equivalence point is slightly basic, pH = 7.5−8. This is because the conjugate base (salt) of a weak acid is not neutral, but rather is itself a weak base.

Sample Data

Introductory Activity

Results for White Grape Juice

Standardized NaOH: 0.0971 M Amount of acid in white grape juice: Calculations shown using citric acid. Check ingredients label to know which acid is applicable to the beverage being tested.

Phenolphthalein endpoint: 11.00 mL
0.0110 L x 0.0971 M = 0.00107 mol NaO
0.00107 NaOH x 1 mol citric acid/3 mol NaOH = 3.57 x 10^–4 mol citric acid
3.57 x 10^–4 mol acid x 192 g = 0.0684 g in 20 mL juice
0.0684 g/20 mL x 88 mL = 0.301 g in one bottle of juice.
Amount of juice per bottle depends on the bottle volume. In this example, it was 3 oz or 88 mL.

Results for Pinapple Juice

Standardized NaOH: 0.0971 M Amount of acid in pineapple juice: Calculations shown using citric acid. Check ingredients label to know which acid is applicable to the beverage being tested.

Phenolphthalein endpoint: 20.60 mL
0.0206 L x 0.0971 M = 0.00200 mol NaOH
0.00200 NaOH x 1 mol citric acid/3 mol NaOH = 6.67 x 10^–4 mol citric acid
6.67 x 10^–4 mol acid x 192 g = 0.128 g in 20 mL juice
0.128 g/20 mL x 177 mL = 1.13 g in one bottle of juice.
Amount of juice per bottle depends on the bottle volume. In this example, it was 6 oz or 177 mL.

Results for Orange Juice

Standardized NaOH: 0.0971 M Amount of acid in orange juice: Calculations shown using citric acid. Check ingredients label to know which acid is applicable to the beverage being tested.

Phenolphthalein endpoint: 22.50 mL
0.0225 L x 0.0971 M = 0.00218 mol NaOH
0.002 NaOH x 1 mol citric acid/3 mol NaOH = 7.28 x 10^–4 mol citric acid
7.28 x 10^–4 mol acid x 192 g = 0.140 g in 20 mL juice
0.140 g/20 mL x 200 mL = 1.40 g in one bottle of juice.
Amount of juice per bottle depends on the bottle volume. In this example, the bottle size was 200 mL.

Answers to Questions

Guided-Inquiry Discussion

1. Choose a suitable indicator for determining the endpoint in neutralization of a weak acid with a strong base. Explain your reasoning based on the evidence obtained above as well as the titration curve data discussed in Prelab Question 4.

   Phenolphthalein is a suitable indicator for titrating a weak acid with sodium hydroxide. The endpoint color is very sudden and coincides with the addition of an equivalent amount of sodium hydroxide (that is, the equivalence point). The color change for acetic acid with bromthymol blue indicator comes before an equivalent amount of sodium hydroxide has been added. Thymol blue gives a series of color changes from yellow-orange to green and then blue when sodium hydroxide is added to acetic acid. The true endpoint is difficult to identify.

2. Would you expect any differences in the choice of an appropriate indicator for titration of a strong acid such as HCl. Why or why not?

   Phenolphthalein appears to be a suitable indicator for titration of either a strong acid or a weak acid. The pH changes very abruptly in the vicinity of the equivalence point. Addition of even one drop of sodium hydroxide after the equivalence point has been reached will increase the pH of the solution to greater than 8.0, where phenolphthalein changes from colorless to pink or red.

3. Acidic beverages generally contain weak acids, such as citric acid in citrus fruit juices, tartaric or malic acids in other fruit juices, phosphoric acid in colas, and carbonic acid in seltzers. Write balanced chemical equations and determine the mole ratio for the reaction of each acid with sodium hydroxide. Note: Use the molecular formulas of the weak acids (it is not necessary to draw out their chemical structures).

   Citric acid           C₆H₈O₇ + 3NaOH → Na₃C₆H₅O₇ + 3H₂O   1 mole citric acid/3 moles NaOH
   Phosphoric acid   H₃PO₄ + 3NaOH → Na₃PO₄ + 3H₂O            1 mole phosphoric acid/3 moles NaOH
   Tartaric acid       C₄H₆O₆ + 2NaOH → Na₂C₄H₄O₆ + 2H₂O   1 mole tartaric acid/2 moles NaOH
   Malic acid          C₄H₆O₅ + 2NaOH → Na₂C₄H₄O₅ + 2H₂O   1 mole malic acid/2 moles NaOH
   Carbonic acid     H₂CO₃ + 2NaOH → Na₂CO₃ + 2H₂O          1 mole carbonic acid/2 moles NaOH

4. The titrant used in a titration experiment is a standard solution. Explain what this means, identify the titrant and obtain the known molarity from your instructor.

   A standard solution is one whose true (accurate) concentration is precisely known, usually to 3 significant figures. The titrant in this experiment is sodium hydroxide. The concentration is 0.0971 M. Note: The concentration will vary. Please see the standardization procedure in the Prelab Preparation section.

5. Review the setup shown in Figure 2 for a titration procedure.
   • The buret should be cleaned and then rinsed with the titrant before beginning the titration. Explain why this is necessary.

     Clean the buret with distilled or deionized water and rinse the buret with two 5–10 mL portions of the titrant to avoid contaminating and/or diluting the standard solution. Its true or actual concentration must be precisely known.

   • Is it necessary to know the precise volume of beverage that will be titrated? Explain.

     The initial volume of the beverage to be titrated must be precisely measured and known because it enters into the equation for calculating the molar concentration of acid in the beverage. Molarity is moles per liter. The number of moles is determined from the volume of titrant used and the mole ratio for reaction with the acid in the beverage. The volume of the beverage must be measured.

   • Choose the type of volumetric glassware (e.g., flask, graduated cylinder or pipet) to measure the beverage(s) that will be titrated in this experiment. Explain your choice.

     The best type of volumetric glassware for measuring the initial volume of beverage is a serological pipet or a graduated cylinder. Serological pipets are more accurate but may become plugged if the juice contains pulp. A graduated cylinder is therefore more practical while still giving three significant figures in precision.

   • It’s helpful to occasionally rinse the sides of the beaker or flask with distilled water during the titration procedure. Explain why it is not necessary to know the volume of rinse water.

     The volume of rinse water does not enter into the final calculations for the concentration of acid in a beverage. Only the initial volume of the beverage must be measured.

6. Examine a buret and explain how to “read” the volume of titrant in the buret. What precision (number of significant figures) is allowed in these measurements?

   The volume is read from the top-down in a buret. The buret is marked every 0.1 mL and thus the volume may be estimated to two decimal places (see Figure 3).

   {13762_Answers_Figure_3_How to read a buret volume}
7. What data must be measured and plotted to obtain the titration curve for an acidic beverage? What is an appropriate volume interval for obtaining this data during the titration? Explain your reasoning.

   A titration curve is obtained by plotting pH on the y-axis versus volume of titrant added in mL on the x-axis. An appropriate volume interval is to measure the pH after every 1.0 mL of titrant is added until about 80% of the estimated volume needed to reach the equivalence point has been added. When approaching the equivalence point, it is best to measure the pH after every 0.2 to 0.5 mL of titrant is added. The reason is that the pH rises very steeply in the region of the equivalence point.

Answers to Review Questions for AP^® Chemistry

1. Why is phenolphthalein an appropriate indicator for titration of a strong acid with a strong base? Explain based on the pH at the equivalence point and the transition range for phenolphthalein.

   The equivalence point in the titration of a strong acid with a strong base occurs at pH 7 since all the products are neutral. Phenolphthalein is an appropriate indicator since its color transition occurs at pH 8−10. This is the region of the titration curve where the pH rise is very steep, so the endpoint is likely to differ from the equivalence point by only a drop or two of titrant.

2. A 10.00-mL sample of HCl solution was transferred to an Erlenmeyer flask and diluted by adding about 40 mL of distilled water. Phenolphthalein indicator was added, and the solution was titrated with 0.215 M NaOH until the indicator just turned pink. The exact volume of NaOH required was 22.75 mL. Calculate the concentration of HCl in the original 10.00-mL sample.

   Moles of NaOH = 0.02275 L x 0.215 moles/L = 0.00489 moles
   Mole ratio = 1 mole HCl/1 mole NaOH
   Concentration of HCl = 0.00489 moles/0.0100 L = 0.489 M

3. One student accidentally “overshot” the endpoint and added 23.90 mL of 0.215 M NaOH. Is the calculated concentration of HCl likely to be too high or too low as a result of this error?

   Overshooting the endpoint will lead to a higher calculated number of moles of NaOH required for neutralization, which will lead in turn to a higher calculated concentration of HCl.

4. The color of an indicator solution depends on pH and the relative amount of HIn and In^– at a given pH. Consider the following indicators and their acidic and basic colors, as well as the pH transition range for each.
   {13762_Answers_Table_1}

   *Alizarin has two ionizable hydrogen atoms and three color forms, H₂In, HIn^– and In₂^–.

   1. The intermediate or transition color of bromthymol blue is green. What are the relative proportions of HIn and In^– when bromthymol blue is green? Explain.

      Bromthymol blue indicator will be green when the solution contains roughly equal proportions of the yellow acid form HIn and the blue base form In⁻.

   2. A colorless solution was tested with phenolphthalein, bromthymol blue and alizarin. The solution was colorless with phenolphthalein, yellow with bromthymol blue and orange with alizarin. What is the pH of the solution? Explain.

      Solution is colorless with phenolphthalein, implying pH < 8. Yellow with bromthymol blue implies pH < 6.0. Finally, orange with alizarin occurs in the transition range for this indicator (between yellow and red), which is pH 5.5–6.8. Mixtures of indicators provide more precise measurements of pH than a single indicator.

References

AP^® Chemistry Guided-Inquiry Experiments: Applying the Science Practices; The College Board: New York, NY, 2013.

Harris, D.C. Exploring Chemical Analysis, 3rd ed.; W. H. Freeman and Company: New York, 2005.

Introduction

Common beverages may be either acidic or basic. Fruit juices, for example, get their sweet taste from sugars and their sour or tart taste from weak acids, such as citric acid. If the juice contains too much sugar, it will taste bland, but too much acid and the juice will taste sour. The concentration of acids in various consumer beverages may be determined by titration with sodium hydroxide.

Concepts

• Acids and bases

• Indicators
• Equivalence point
• Titration
• pH
• Neutralization

Background

The main acids present in fruits and fruit juices are citric acid (in citrus fruits), tartaric acid (in grapes) and malic acid (in apples). All of these are characterized as weak acids.

The amount of citric acid in citrus fruit juices can be determined by titration with a standard solution of sodium hydroxide. A standard solution is one whose concentration is accurately known, usually to three significant figures. Citric acid is a tricarboxylic acid—it has three ionizable or “active” hydrogen atoms in its structure. One mole of citric acid therefore reacts with three moles of sodium hydroxide via the acid–base neutralization reaction shown in Equation 1. Acid–base titrations are an extremely useful technique to determine the concentration of an acid or base in a sample. In titrating beverages such as orange juice, apple juice, and sodas that contain weak acids, the juice is called the analyte and a strong base is used as the titrant.

In the titration procedure, a sodium hydroxide solution of known molarity is carefully added using a buret to a measured volume of fruit juice containing an indicator. The exact volume of sodium hydroxide that must be added to reach the indicator endpoint is measured and then used to calculate the concentration of citric acid in the juice.

A sample setup for a titration is shown in Figure 2, where a buret containing the titrant is clamped to the support stand and a beaker or flask containing the analyte is set a-top a stir plate. If a pH probe is inserted into the solution, a titration curve can be constructed by plotting the pH of the solution on the y-axis versus the volume of titrant added on the x-axis. The shape of the titration curve may be used to distinguish strong and weak acids in the analyte, and also permits graphical analysis of the equivalence point. At the equivalence point, moles of added titrant are stoichiometrically related to moles of analyte in the sample. Choosing a suitable indicator for a titration is important for accurate results. Indicators signify the endpoint of a titration when a sudden change in the color of the analyte solution occurs. Indicators have different pH transition ranges and exhibit different colors in acidic versus basic solutions. The color changes arise because indicators are weak acids for which the acid form HIn and the conjugate base form In^– have different colors. An appropriate indicator for a titration is one whose color change occurs close to the theoretical pH of the equivalence point. Examples of indicators provided in this activity are shown in the following table, along with their colors and pH ranges.

Experiment Overview

The purpose of this advanced inquiry lab is to conduct acid−base titrations and determine the concentration of acid in common beverages such as orange juice or pineapple juice. The beverages contain weak acids, which will be titrated with a strong base, sodium hydroxide. The lab begins with an introductory activity to determine the proper indicator to use in the titration of acetic acid, a characteristic weak acid. The results provide a model for guided-inquiry design of a titration procedure to obtain titration curve data and calculate the molar concentration of acid in a beverage. The titration curve will be analyzed and the amount of acid in a typical serving size or bottle may also be determined. The identity of the acid in the beverage may be derived by reviewing the titration curve and reference information and by consulting the ingredients label.

Materials

Prelab Questions

1. Using the structural formula of citric acid shown in Figure 1, determine the molecular formula of citric acid and calculate its molar mass (g/mole).
2. A 10.0-mL sample of pineapple juice was titrated with 0.100 M sodium hydroxide solution. The average volume of NaOH required to reach the endpoint was 12.8 mL.
   1. Calculate the number of moles of sodium hydroxide required to reach the endpoint.
   2. Using the mole ratio for the neutralization reaction shown in Equation 1, determine the number of moles of citric acid in 10.0 mL of pineapple juice.
   3. Multiply the number of moles of citric acid by its molar mass to calculate the mass of citric acid in 10.0 mL of the juice.
   4. The concentration of acid in juices is usually expressed in grams of acid per 100 mL of juice. What is the concentration of citric acid in pineapple juice?
3. Write a balanced chemical equation for the neutralization reaction of (a) hydrochloric acid and (b) acetic acid with sodium hydroxide.
4. The titration curves for hydrochloric acid and acetic acid with sodium hydroxide are shown below. Distinguish between the strong and weak acid in terms of the initial pH, the pH at the equivalence point, and the overall shape of the titration curve.

Safety Precautions

Dilute sodium hydroxide and acid solutions are irritating to skin and eyes. Avoid contact of all chemicals with eyes and skin. All 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 chemistry laboratory. Do not remove any remaining food items from the lab after they have been used in the lab. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

Introductory Activity

Indicators for Titration of a Weak Acid

1. Label three medium test tubes B, P and T for the names of three indicators—bromthymol blue, phenolphthalein, and thymol blue—that will be studied in this activity.
2. Using a 10-mL graduated cylinder, measure and pour 2.0 mL of 0.1 M acetic acid into each test tube.
3. Add 1−2 drops of each indicator to the appropriate test tube.
4. Observe and record the initial indicator color in each test tube.
5. Rinse the graduated cylinder with distilled water and dry the cylinder.
6. Measure 3.0 mL of 0.1 M sodium hydroxide in the graduated cylinder. Using a graduated pipet, add the NaOH solution in 1 mL increments to the acetic acid solution in test tube B. Observe and record the indicator color as the base is added.
7. Note the approximate volume of NaOH that has been added when the indicator color changes.
8. Repeat steps 6−7 two times using the acetic acid−indicator solutions in test tubes P and T.
9. Rinse the test tubes with distilled water and dry them.
10. (Optional) Using a clean, 10-mL graduated cylinder, measure and pour 2.0 mL of 0.1 M hydrochloric acid into each test tube B, P and T. Repeat steps 3−8 to determine the initial and final color changes for HCl and NaOH with various indicators.

Guided-Inquiry Design and Procedure

Titration Curves and the Concentration of Acids in Fruit Juices

Form a working group with other students and discuss the following questions.

1. Choose a suitable indicator for determining the endpoint in the neutralization of a weak acid with a strong base. Explain your reasoning based on the evidence obtained above as well as the titration curve data discussed in Prelab Question 4.
2. Would you expect any differences in the choice of an appropriate indicator for the titration of a strong acid such as HCl? Why or why not?
3. Acidic beverages generally contain weak acids, such as citric acid in citrus fruit juices, tartaric or malic acids in other fruit juices, phosphoric acid in colas, and carbonic acid in seltzers. Write balanced chemical equations and determine the mole ratio for the reaction of each acid with sodium hydroxide. Note: Use the molecular formulas of the weak acids (it is not necessary to draw their chemical structures).
4. The titrant used in a titration experiment is a standard solution. Explain what this means, identify the titrant, and obtain the known molarity from your instructor.
5. Review the setup shown in Figure 2 for a titration procedure.

• The buret should be cleaned and then rinsed with the titrant before beginning the titration. Explain why this is necessary.
• Is it necessary to know the precise volume of beverage that will be titrated? Explain.
• Choose the type of volumetric glassware (e.g., flask, graduated cylinder or pipet) to measure the beverage(s) that will be titrated in this experiment. Explain the choice.
• It’s helpful to occasionally rinse the sides of the beaker or flask with distilled water during the titration procedure. Explain why it is not necessary to know the volume of rinse water.

6. Examine a buret and explain how to “read” the volume of titrant in the buret. What precision (number of significant figures) is allowed in these measurements?
7. What data must be measured and plotted to obtain the titration curve for an acidic beverage? What is an appropriate volume interval for obtaining this data during the titration? Explain your reasoning.
8. Write a detailed, step-by-step procedure for titrating a beverage to determine the concentration of weak acid, if present. Include the reagents needed, the glassware and equipment that will be used, and the appropriate measurements and observations that must be made.
9. Review the hazards of the chemicals used in the procedure and write appropriate safety precautions that must be followed during the experiment.
10. Carry out a “rough” titration to estimate the volume of beverage to be used in the experiment. Pour 5 mL of juice into a test tube, add 1−2 drops of indicator, and note the initial color. Add the titrant in 1 mL increments using a graduated pipet until the endpoint color is observed. Keep the test tube to be used as a “color standard” for the titration.
11. Choose an amount of beverage to be titrated that will require at least 10 mL but less than 20 mL of titrant. Explain why this range of titrant is optimum.
12. Carry out the titration to obtain the titration curve data. Record the results in an appropriate data table.
13. Repeat the titration as needed to check the reproducibility of the endpoint measurement. It is not necessary to use the pH meter for the additional trials. Record results.

Analyze the Results: Plot the data and explain the titration curve results, including the initial pH and the pH at the equivalence point. Determine the molar concentration of acid in the beverage sample based on its ingredient label and/or the most probable acid it contains. Calculate the mass of acid contained in a bottle or serving size of the beverage.

Student Worksheet PDF

13762_Student1.pdf