Materials Included In Kit

Additional Materials Required

Safety Precautions

Phenolphthalein and thymolphthalein solutions contain alcohol and are flammable liquids; they are toxic by ingestion. Do not use near flames or other sources of ignition. Dilute sodium hydroxide solution is slightly toxic by ingestion and skin absorption and is irritating to skin and eyes. The ammonia water solution is mildly toxic by ingestion and inhalation, irritating to body tissues and a lachrymator. Solid calcium hydroxide is toxic by inhalation, irritating to body tissue and its solution is caustic. Acetic acid may cause respiratory tract irritation. Hydrochloric acid is slightly toxic by inhalation and ingestion, a severe body tissue irritant and corrosive to eyes. Sulfuric acid is slightly toxic by ingestion and severely irritating to body tissues, especially eyes. Nitric acid is irritating to body tissues. Avoid contact of all chemicals with eyes and skin. 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. The acids may be neutralized according to Flinn Suggested Disposal Method #24a. The bases may be neutralized according to Flinn Suggested Disposal Method #10. The titrated solutions are considered neutral and may be rinsed down the drain with plenty of water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• This laboratory activity was specifically written, per teacher request, to be completed in one 50-minute class period. It is important to allow time between the Prelab Homework Assignment and the Lab Activity. Prior to beginning the homework, show the students the titration setup and chemicals—this will get the procedure thought process rolling. Once students turn in the homework answers, graphs and figures and their procedure, check it for safety and accuracy before implementation in the lab.
• Guide the students to use 10 mL of analyte for chemical economy.|Make sure to check the chemicals for accuracy before students start.
• Bromythymol blue indicator solution is included in Acidity of Beverages—Inquiry Lab Kit for AP^® Chemistry, available from Flinn Scientific, Catalog No. AP7645.
• If the appropriate chemicals are available, students may standardize the titrants. Standardization is incorporated in AP^® Chemistry Investigation 4, Acidity of Beverages—Inquiry Lab Kit for AP^® Chemistry, Flinn Scientific, Catalog No. AP7645.
• Students should rinse the buret with the titrant. Explain to students that rinsing the buret with water may change the initial concentration of the titrant.
• Students should monitor the pH changes during the titration after every 0.50-mL to 1.00-mL addition of titrant. Economical “personal” pH meters, such as the Flinn pH meter (Catalog No. AP8673) may be used.
• If magnetic stirrers are not available, have the students swirl the flask after each addition of NaOH solution, then measure the pH value.
• Remind students to read the volume in a buret from the top-down. A buret is marked every 0.1 mL and thus the volume may be estimated to two decimal places (see Figure 4).
  {14106_PreLabAnswers_Figure_4}

Teacher Tips

• Quantitative analysis represents a nearly invisible application of chemistry in our daily lives. To illustrate the importance of quantitative analysis, ask students how they would feel if they could not trust that the water they drink or the medicines they take had been tested to assure quality and safety.
• You may include more indicators than are available in this kit as options during the lab.
• This is an excellent inquiry opportunity because everything you need to know about a weak acid can be determined from the titration curve! The initial pH of the solution provides information about the strength of the acid, or its concentration. The pH at the equivalence point reveals that the conjugate base of the weak acid is also a weak base, while the volume of base needed to reach the equivalence point is used to calculate the concentration of the acid. Finally, the pH at the “halfneutralization” point is equal to the pK_a of the weak acid, and the shape of the curve in this region demonstrates the buffering capacity of a weak acid and its conjugate base. Help students integrate all of these key learning objectives by having them present their findings in a poster or presentation.
• Due to the low solubility of calcium hydroxide in water, an accurate solution with known molarity cannot be made.

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.

Chemical reactions can be classified by considering what the reactants are, what the products are, or how they change from one into the other. Classes of chemical reactions include synthesis, decomposition, acid–base, and oxidation–reduction reactions. (3B)
3B2: In a neutralization reaction, protons are transferred from an acid to a base.

Chemical equilibrium plays an important role in acid–base chemistry and in solubility. (6C)
6C1: Chemical equilibrium reasoning can be used to describe the proton-transfer reactions of acid–base chemistry.

Learning Objectives
1.18 The student is able to apply conservation of atoms to the rearrangement of atoms in various processes.
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.
3.4 The student is able to relate quantities (measured mass of substances, volumes of solutions, or volumes and pressures of gases) to identify stoichiometric relationships for a reaction, including situations involving limiting reactants and situations in which the reaction has not gone to completion.
3.7 The student is able to identify compounds as Brønsted-Lowry acids, bases, and/or conjugate acid–base pairs, using proton-transfer reactions to justify the identification.
6.11 The student can generate or use a particulate representation of an acid (strong or weak or polyprotic) and a strong base to explain the species that will have large versus small concentrations at equilibrium.
6.12 The student can reason about the distinction between strong and weak acid solutions with similar values of pH, including the percent ionization of the acids, the concentrations needed to achieve the same pH, and the amount of base needed to reach the equivalence point in a titration.
6.13 The student can interpret titration data for monoprotic or polyprotic acids involving titration of a weak or strong acid by a strong base (or a weak or strong base by a strong acid) to determine the concentration of the titrant and the pK_a for a weak acid, or the pK_b for a weak base.
6.15 The student can identify a given solution as containing a mixture of strong acids and/or bases and calculate or estimate the pH (and concentrations of all chemical species) in the resulting solution.
6.16 The student can identify a given solution as being the solution of a monoprotic weak acid or base (including salts in which one ion is a weak acid or base), calculate the pH and concentration of all species in the solution, and/or infer the relative strengths of the weak acids or bases from given equilibrium concentrations.
6.17 The student can, given an arbitrary mixture of weak and strong acids and bases (including polyprotic systems), determine which species will react strongly with one another (i.e. with K > 1) and what species will be present in large concentrations at equilibrium.

Science Practices
1.2 The student can describe representations and models of natural or man-made phenomena and systems in the domain.
1.4 The student can use representations and models to analyze situations or solve problems qualitatively and quantitatively.
2.1 The student can justify the selection of a mathematical routine to solve problems. (Appropriateness of selected mathematical routine.)
2.2 The student can apply mathematical routines to quantities that describe natural phenomena.
3.1 The student can pose scientific questions.
3.3 The student can evaluate scientific questions.
4.1 The student can justify the selection of the kind of data needed to answer a particular scientific question.
4.2 The student can design a plan for collecting data to answer a particular scientific question.
4.3 The student can collect data to answer a particular scientific question.
4.4 The student can evaluate sources of data to answer a particular scientific question.
5.1 The student can analyze data to identify patterns or relationships.
5.3 The student can evaluate the evidence provided by data sets in relation to a particular scientific question.
6.1 The student can justify claims with evidence.
6.2 The student can construct explanations of phenomena based on evidence produced through scientific practices.
7.2 The student can connect concepts in and across domain(s) to generalize or extrapolate in and/or across enduring understandings and/or big ideas.

Answers to Prelab Questions

Introduction to Acids and Bases

1. Define and provide an example list of common strong acids and bases and weak acids and weak bases.

   Strong acids and bases dissociate completely while weak acids and bases partially dissociate and form an equilibrium in an aqueous solution. Examples of each: strong acids: HCl, H₂SO₄, HNO₃, strong bases: group I and II hydroxides (select group II hydroxides are slightly soluble in water, where there is dissolved base, it is completely ionized), weak acids: acetic acid, oxalic acid, weak bases: ammonia, amines, pyridines.

2. See the particulate models in Figure 2 to answer Questions a–d.
   {14106_PreLabAnswers_Figure_2}
   1. Write the chemical reactions taking place.

      See Figure 2.

   2. Identify conjugate pairs by labeling the reactants and products as acids or bases.

      HF is the conjugate acid; F^– is the conjugate base. H₂O is the conjugate base; H₃0⁺ is the conjugate acid. HCl is the conjugate acid; Cl^– is the conjugate base.

   3. Which chemical reaction illuminates the light bulb the brightest? Which light bulb is dimmer? Explain.

      The beaker containing the HCl solution illuminates the bulb brighter than the beaker containing the HF solution. Since HCl is a strong acid, it is a better conductor because it completely ionizes in solution.

   4. Write the equilibrium constant expression, K_a, for each and predict the relative magnitudes of K_a. Describe what structural factors influence dissociation and, thus, the K_a value.

      HCl: HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl^–(aq)
      K_a = [H₃O⁺][Cl^–]/[HCl]

      {14106_PreLabAnswers_Reaction_1}

      K_a = [H₃O⁺][F^–]/[HF]

      In the beaker containing aqueous HCl: since the strong acid molecules completely ionize in solution and produce hydronium and chloride ions, [HA] (HCl) is very small, thus K_a is very large. In the case of aqueous HF, which does not completely ionize, K_a is small because [HA] (HF) is large.
      
      The structural feature of HF: the strong hydrogen bonding is attributed to making it a weak acid. Since the fluoride ion is very electronegative (in fact, fluorine is the most electronegative element on the periodic table) it “traps” the hydrogen atoms and does not allow complete dissociation. The chloride ion in HCl is less electronegative, it does not feature the same structural characteristics as HF. In addition, HCl has a bigger bond length than HF, and is thus weaker.

3. Draw a similar particulate model from Question 2 for an aqueous solution of 0.01 M ammonia.
   {14106_PreLabAnswers_Figure_5}
4. Answer a–e regarding the relationships between [H⁺], [OH^–] and pH.
   1. Water may behave as an acid or a base:
      {14106_PreLabAnswers_Reaction_3}
      where the equilibrium constant, K_w, equals 1 x 10^–14 at 25 °C. Write the equilibrium expression.

      K_w = [H⁺][OH^–] = 1 x 10^–14 at 25 °C

   2. What are the concentrations of [H⁺] and [OH^–] from 4a?

      Each has a concentration of 1 x 10^–7 M.

   3. Calculate the pH of 0.10 M HCl and 0.10 M CH₃COOH. Show your work. Hint: pH = –log[H⁺]

      HCl: pH = –log[0.10] = 1
      CH₃COOH: K_a = 1.75 x 10^–5
      K_a = [H₃O⁺][A^–]/[HA] = (x)(x)/I – x
      x²/0.10 = 1.75 x 10^–5 = x = 0.00132 M
      pH = –log(0.00132) = 2.87

   4. Seek out education resources and observe the pH values of common acids and bases, both strong and weak. Can a weak acid and a strong acid have the same pH?

      Yes, there will be differences in percent ionization, but both may exhibit the same pH values due to the same [H₃O⁺] in solution.

   5. Go back to Question 3. Write the equilibrium constant expression, calculate it, and calculate the pH. Show your work. Hint: K_a for ammonium ion is 5.69 x 10^–10.
      {14106_PreLabAnswers_Reaction_2}

      K_a x K_b = K_w
      K_b = K_w/K_a = 1.0 x 10^–14/5.69 x 10^–10 = 1.76 x 10^–5
      K_b = [NH₄⁺][ OH^–]/[NH₃] = x²/0.01 = 1.76 x 10^–5 = x = 4.20 x 10^–4
      [H⁺] = 1 x 10^–14/4.20 x 10^–4 = 2.38 x 10^–11
      pH = –log(2.38 x 10^–11) = 10.6

Neutralization Reactions

5. Acids and bases go to completion via neutralization reactions, thus titrations are applicable. Refer to educational resources and provide an example of the chemical reactions for the solutions in a–c.

   Student answers will vary.

   1. A mixture between a strong acid and a strong base.

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

   2. A strong base mixed with weak acid.

      NaOH(aq) + CH₃COOH(aq) → NaCOOCH₃(aq) + H₂O(l)

   3. A strong acid mixed with weak base.

      HCl(aq) + NH₃(aq) → NH₄Cl(aq)

Design a Titration Experiment

6. A student has to design a titration experiment to determine the concentration of an unknown acid. Answer Questions a–f.
   1. See Figure 3. Label the equipment the student needs to perform a titration. Write answers next to the blanks.
      {14106_PreLabAnswers_Figure_3}
   2. The student performs a “rough” titration using beakers, beral-type pipets and a pH meter. The purpose is to determine the proper acid–base indicator in signaling the endpoint in the titration of various combinations of acids and bases. See the following student data table. Fill in the table for the best indicator to use in each. Note: More than one indicator may apply.
      {14106_PreLabAnswers_Table_1}
      {14106_PreLabAnswers_Table_2}
   3. After the “rough titration,” the student gathered materials and performed a titration three times. Graph the students’ titration data below of a weak acid (acetic acid) titrated with a strong base (0.10 M sodium hydroxide) utilizing a spreadsheet program. The student selected thymol blue as the indicator. Give a reason why this is a good indicator to use for this titration.
      {14106_PreLabAnswers_Figure_6}

      Thymol blue is a good indicator because its pH transition interval is 8.0–9.2. The pH at the equivalence point is 8.5.

   4. On the generated titration plot from part c, using arrows, label the pH at half equivalence, the one half-equivalence point volume, the equivalence point volume, the pH at the equivalence point and the end titration.

      See titration curve from 6c.

   5. Calculate K_a.

      When [A^–] = [HA], K_a reduces to [H₃O⁺] and pK_a = –log K_a = pH.
      This condition is met at the “half-equivalence point” in the titration. When one-half the volume of NaOH needed to reach the equivalence point has been added, the pH of the solution is equal to the pKa value for the acid. Draw dashed lines on the titration curve in part C to illustrate these relationships, and estimate the pK_a value for the weak acid. The volume to reach the equivalence point is 25 mL. At half-neutralization, 12.5 mL, the pH = 4.8. Therefore, the pK_a of the acid is 4.8 and the value of K_a = 10^–4.8 = 1.6 x 10^–5.

   6. Calculate the concentration of the acid (acetic acid).
      {14106_PreLabAnswers_Equation_2}
   7. Write a general step-by-step procedure to determine the concentration of an unknown acid or base. On lab day, your instructor will provide you with a titration curve to reproduce. The procedure should be written in the lab notebook to be used on the day of lab.

      Helpful tips:

      1. Think safety, first. Make sure you have the proper PPE available to perform this lab (i.e., goggles, apron and gloves).
      2. Make a list of the equipment and glassware needed for this lab.
      3. Number the steps in your procedure; remember to be as detailed as possible, from set-up to clean-up.
      4. Draw necessary data tables in your notebook for data collection during the lab.
      5. Be prepared to pick your acid and base on lab day to reproduce the given titration curve. Watch the amount of chemical volumes for the three trials—you don’t want to run out!
      6. To record the volume in a buret, read it from the top-down. A buret is marked every 0.1 mL and thus the volume may be estimated to two decimal places (see Figure 4).
         {14106_PreLabAnswers_Figure_4}

Sample Data

Titration of 25.0 mL of CH₃COOH with 0.10 M NaOH (Note: the sample data provided were performed with various analyte sample volumes, ranging from 10 mL to 25 mL.

pH Equivalence Point (calc) ______8.5______ Selected Indicator ______Phenolphthalein______

Titration Curves 

The initial volume of analyte was 10.0 mL in each case. Note that the concentrations of acid and base were not equimolar in each case. The pH at the equivilence point and the shape of the curve are the essential features. Titration Curves with Calculations

Teacher Handouts

14106_Teacher1.pdf

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

Experience and learn the concepts you need to help you succeed on the AP^® Chemistry exam with this guided-inquiry activity! We can study acid–base chemistry by applying general principles of chemical equilibrium. An exciting application is utilizing titrations to determine the concentration of acid or base in a solution. This is an important quantitative lab technique because acids and bases are present in our everyday lives—they exist in commercial products from beverages, to cleaners, to OTC medications and, essentially, in the natural world around us. The homework set will guide you through understanding Brønsted-Lowry acid–base chemistry of strong acids and bases and weak acids and bases and how to properly interpret their titration data. Then, design your very own experiment for lab day! You will be challenged to reproduce a given titration curve, graphical analysis will confirm your results. Did you choose the correct chemicals to reproduce the curve you were given?

Concepts

• Strong and weak acids
• Equilibrium constant
• K_a
• Titration
• Strong and weak bases
• Equivalence point
• Indicators

Background

Titration is a method of volumetric analysis—the use of volume measurements to analyze an unknown. In acid–base chemistry, titration is most often used to analyze the amount of acid or base in a sample or solution. Consider a solution containing an unknown amount of hydrochloric acid. In a titration experiment, a known volume of the hydrochloric acid solution would be “titrated” by slowly adding dropwise a standard solution of a strong base such as sodium hydroxide. (A standard solution is one whose concentration is accurately known.) The titrant, sodium hydroxide in this case, reacts with and consumes the acid via a neutralization reaction (Equation 1). The exact volume of base needed to react completely with the acid is measured. This is called the equivalence point of the titration—the point at which stoichiometric amounts of the acid and base have combined.

Knowing the exact concentration and volume of the added titrant gives the number of moles of sodium hydroxide, which is, in turn, related by the mole ratio to the number of moles of hydrochloric acid initially present in the unknown.

Either acids or bases may be titrated to determine their concentration by choosing an appropriate standard solution as the titrant. Indicators are usually added to acid–base titrations to detect the equivalence point. The endpoint of the titration is the point at which the indicator changes color and signals that the equivalence point has indeed been reached. For example, in the case of the neutralization reaction shown in Equation 1, the pH of the solution would be acidic (< 7) before the equivalence point and basic (> 7) after the equivalence point if excess sodium hydroxide is added. The pH at the equivalence point should be exactly 7, corresponding to the neutral products—sodium chloride and water. An indicator that changes color around pH 7 is therefore a suitable indicator for the titration of a strong acid with a strong base.

The progress of an acid–base titration can also be followed by measuring the pH of the solution being analyzed as a function of the volume of titrant added. A plot of the resulting data is called a pH curve or titration curve. Titration curves allow a precise determination of the equivalence point of the titration without the use of an indicator.

The graph of pH versus volume of NaOH added for the titration of HCl is shown in Figure 1. Note the significant change in pH in the vicinity of the equivalence point. When a weak acid is titrated with a strong base, the equivalence point is not at pH 7, but rather is on the basic side. The value of the equilibrium constant for the dissociation of a weak acid can be obtained from its titration curve with a strong base. The shape of the titration curve for a weak acid with a strong base is explored in the Prelab Homework Assignment, along with the equilibrium constant determination.

Experiment Overview

The purpose of this activity is to complete the homework assignment prior to lab to promote conceptual understanding of acids and bases and the experimental design of a classic titration. You will be presented with one of four titration curve plots to successfully reproduce on lab day. The homework will build your knowledge to interpret particulate model diagrams, extrapolate from a titration curve plot, perform pertinent calculations, and ultimately write your very own procedure. Take ownership of selecting the correct chemicals—an acid, a base and an indicator—to correctly reproduce the titration curve. Learn to setup the titration equipment to ensure success and accurate results. You’ll love the challenge.

Prelab Questions

See Student PDF.

Safety Precautions

Phenolphthalein and thymolphthalein solutions contain alcohol and are flammable liquids; they are toxic by ingestion. Do not use near flames or other sources of ignition. Dilute sodium hydroxide solution is slightly toxic by ingestion and skin absorption and is irritating to skin and eyes. The ammonia water solution is mildly toxic by ingestion and inhalation, irritating to body tissues, and a lachrymator. Solid calcium hydroxide is toxic by inhalation, irritating to body tissue, and its solution is caustic. Acetic acid may cause respiratory tract irritation. Hydrochloric acid is slightly toxic by inhalation and ingestion, a severe body tissue irritant, and corrosive to eyes. Sulfuric acid is slightly toxic by ingestion and severely irritating to body tissues, especially eyes. Nitric acid is irritating to body tissues. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Student Worksheet PDF

14106_Student1.pdf