Inquiry Lab Kit for AP® Chemistry
Materials Included In Kit
Acetic acid solution, CH3CO2H, 0.1 M, 500 mL
Ammonia water solution, NH3, 0.2 M, 1000 mL
Calcium hydroxide, Ca(OH)2, reagent, 50 g*
Hydrochloric acid solution, HCl, 0.2 M, 500 mL
Methyl red indicator solution, 0.02%, 50 mL
Nitric acid solution, HNO3, 0.05 M, 500 mL
Phenolphthalein indicator solution, 1%, 30 mL
Sodium hydroxide solution, NaOH, 0.1 M, 1000 mL
Sulfuric acid solution, H2SO4, 0.1 M, 500 mL
Thymolphthalein indicator solution, 0.04%, 30 mL
pH test strips, vial of 100
*Can only be used as an analyte, not as a titrant.
Additional Materials Required
(for each lab group)
Water, distilled or deionized
Beakers, 50-, 150- and 250-mL
Graduated cylinders, 10- and 100-mL
Magnetic stirrer and stir bar or stirring rod
pH sensor or pH meter
Pipets, Beral-type, graduated
Support stand and buret clamp
Test tubes, medium, 16 mm x150 mm, 4
Test tube rack
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.
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.
- 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.
- In the Introductory Activity, each group will choose two acids and two bases to use for their titrations. Provide the concentration of one of the acids and one of the bases. The students are to determine the concentration of the other acid and other base. When students choose the volume of their unknown in the Guided-Inquiry Design and Procedure, we suggest a volume of 20 mL.
- As there are many combinations of acid–base titrations to be performed, student groups should not share their known concentrations of acids and bases with other student groups.
- 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).
- 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 and the determination of concentration of the two unknowns may be completed after lab or as homework.
- 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.
- 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 “half-neutralization” point is equal to the pKa 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. Please see Sample Data.
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.
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 pKa for a weak acid, or the pKb 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.
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
- Predict whether the pH at the equivalence point will be acidic or basic for the following classic titrations. Explain based on the properties of the conjugate acid–base pairs.
- Titration of a strong acid with a strong base.
Neutral—the conjugate base of a strong acid is neutral.
- Titration of a weak acid with a strong base.
Basic—the conjugate base of a weak acid is a weak base.
- Titration of a strong base with a strong acid.
Neutral—the conjugate acid of a strong base is neutral.
- Titration of a weak base with a strong acid.
Acidic—the conjugate acid of a weak base is a weak acid.
- Distinguish between a strong acid and a weak acid in terms of their dissociation reactions and equilibrium constants.
Strong and weak acids differ in the extent or reversibility of their dissociation reactions and the value of their equilibrium constants. Dissociation of a strong acid is irreversible and the strong acid is 100% ionized in aqueous solution. The value of Ka for a strong acid is much greater than 1. Dissociation of a weak acid is reversible and a weak acid is only partially ionized—generally, 5–10%—in aqueous solution. The value of Ka for a weak acid is less than 1. Similar distinctions apply to strong versus weak bases.
- The following values were collected for the titration of a solid weak acid with 0.100 M NaOH as the titrant. Graph the data in the chart provided and identify the pH at the equivalence point.
The pH is 8.5 at the equivalence point.
- The equilibrium constant Ka for dissociation of a weak acid HA can be determined from its titration curve with a strong base. See Equations 2 and 3.
When [A–] = [HA], Ka reduces to [H3O+], and pKa = –log Ka = 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 Question 3 to illustrate these relationships, and estimate the pKa 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 pKa of the acid is 4.8 and the value of Ka = 10–4.8 = 1.6 x 10–5.
- The following acid–base indicators are available to follow the titration shown in Question 3. Which indicator would be most appropriate for signaling the endpoint of the titration? Explain, and give the expected color to look for at the endpoint.
Thymol blue should be used since the equivalence point occurs at pH 8.5, which is within the color transition range for thymol blue. The indicator color would be green at the endpoint. If bromphenol blue were used, the indicator would be gradually changing color through the first 20 mL of base added. If bromthymol blue were used, the color change would occur too early.
*Bromthymol blue will also be an appropriate indicator for this titration. If the AP® Chemistry bundle (AP7655) was purchased, it is included in Acidity of Beverages—Inquiry Lab Kit for AP® Chemistry.
The following table shows the proper indicator to use for various titrations and their pH transition ranges. For titrations involving strong acids and strong bases only, the indicators provided are close to a neutral pH (7). Monitoring the endpoint using a pH meter as well as an indicator will help students realize that the indicators provided are close, but not exact to the equivalence point of the titration.
Titration of 25.0 mL of CH3
COOH with 0.10 M NaOH
pH Equivalence Point (calc) ______8.5______ Selected Indicator ______Phenolphthalein______
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
Answers to Questions
Answers to Guided-Inquiry Discussion Questions
Answers to Review Questions for AP® Chemistry
- With the instructor’s approval, select two acids and two bases to analyze by titration. Record the name and formula of each compound, and identify each as a strong or weak acid and base, respectively.
Weak acid = Acetic acid
Strong acids = Hydrochloric acid, nitric acid, sulfuric acid
Weak base = Ammonia
Strong base = Sodium hydroxide
Calcium hydroxide is considered a weak base because of its low solubility in water. It behaves as a strong base, however, when it is titrated with an acid, because the OH– ions are neutralized to give water. As the OH– ions are neutralized, more Ca(OH)2 solid dissolves, according to Le Chatelier’s principle, until all of the solid reacts.
- Obtain the known concentration of one each of the acids and bases. These will be used as the titrants in this investigation.
Concentrations of all reactants appear in the Materials list in the Teacher’s Notes.
- Select an appropriate indicator for titration of the unknown acid and base, respectively.
Suitable indicators are shown in the data table for the Introductory Activity.
- Set up a buret with a clamp and support stand, a pH meter or sensor, a beaker and a magnetic stirrer, if available (see Figure 2). Explain why it is desirable to clean and rinse the buret with the titrant before beginning the titration.
It is important to rinse the buret with a small amount of titrant before beginning the procedure to avoid contaminating and/or diluting the concentration of the standard solution.
- Is it necessary to know the exact volume of the “unknown” acid or base to be titrated? Explain.
The exact volume of the analyze (unknown acid or base) must be accurately or precisely measured in order to determine its concentration, which has units of moles per liter.
- It is helpful to occasionally rinse the sides of the beaker or flask with distilled water during the titration procedure. Explain why or why not it is necessary to measure the volume of rinse water used during the procedure.
It is not necessary to measure the volume of rinse water used during the titration since the stoichiometry calculations depend on the number of moles in the unknown. The number of moles does not change when it is diluted. Furthermore, only the initial volume is used in calculating the initial concentration of the unknown.
- Write a detailed, step-by-step procedure for titrating each unknown and obtaining the data for the titration curve. Include the chemicals needed, the indicators and the safety precautions required. Describe the glassware, equipment and techniques needed to ensure accuracy and precision in the quantitative analysis. Choose an amount (volume) of unknown for a convenient titration. Review the hazards of all chemicals and write appropriate safety precautions that must be followed.
Using 20 mL of analyte is optimum for titrations, ensuring that at least 10.0 mL of titrant will be needed to reach the equivalence point for maximum accuracy and also keeping the total volume of titrant convenient.
- Write the balanced chemical equation for the reaction of acetic acid with sodium hydroxide.
CH3COOH(aq) + NaOH(aq) → NaCH3COO(aq) + H2O(l)
- Calculate the average molarity of acetic acid in the vinegar using the data provided.
Trial 1: Ma = (0.50 M)(24 drops)/(15 drops) = 0.80 M
Trial 2: Ma = (0.50 M)(27 drops)/(15 drops) = 0.90 M
Trial 3: Ma = (0.50 M)(24 drops)/(15 drops) = 0.80 M
Trial 4: Ma = (0.50 M)(26 drops)/(15 drops) = 0.87 M
Trial 5: Ma = (0.50 M)(26 drops)/(15 drops) = 0.87 M
Molarity range (Trials 1–5) = 0.80–0.90 M
Estimated error = 0.85 ±0.05 M
- Calculate the average percent acetic acid in the vinegar using the following equation.
Calculate the molar mass of acetic acid:
Molar mass (C2H4O2) = (2 x 12.01) + (4 x 1.008) + (2 x 16.00) = 60.05 g/mole
Use the molar mass of acetic acid and the molarity of the vinegar solution to calculate the number of grams of acetic acid in one liter of vinegar:
Convert the number of grams of acetic acid in one liter of vinegar to percent and convert one liter to milliliters:
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.