Materials Included In Kit

Additional Materials Required

Prelab Preparation

Target Buffer Assignments

Assign each group a buffer solution with a specific target pH. You may give students their assigned buffer when assigning the Prelab Homework Assignment or you can give students the desired pH the day of the wet lab.

Sample pH numbers
2.90, 3.00, 3.20, 3.30 = Citric Acid/Citrate Buffer
4.55, 4.65, 4.85, 4.95 = Acetic Acid/Acetate Buffer
6.20, 6.30, 6.50, 6.60 = Carbonic Acid/Carbonate Buffer

Safety Precautions

Dilute acid and base solutions are skin and eye irritants. Acetic acid solutions may be irritating to the respiratory tract. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Avoid exposure of all chemicals to eyes and skin. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. 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 regulation that may apply, before proceeding. The acetic acid solution may be disposed of according to Flinn Suggested Disposal Method #24a. The buffer solutions, seltzer water and the sodium acetate solution may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Enough materials are provided in this kit for 24 students working in pairs or for 12 groups of students. It is important to allow time between the Prelab Homework Assignment and the lab activity. Once students turn in the homework answers and their procedure, check it for safety and accuracy before they implement it in the lab.
• A calibrated pH meter will be vital to testing your students’ solutions. Flinn has a variety of pH meters, such as the Flinn pH meter (Catalog No. AP8673), portable pH meter (Catalog No. AP7708) and benchtop models (Catalog No. AP7452 and AP7633). Whichever pH meter you choose to use, be sure to calibrate the pH meter before the lab.
• Remind students to take into account the formulas of any hydrates when calculating the molar masses of various compounds.
• Results can vary. A pH within ± 0.05 is a good reading. See sample data from the lab for more information.
• The concentration of carbonic acid can vary in seltzer water. For best results, test the seltzer water and buffer solutions before students use the seltzer water in lab. Testing and lab use of seltzer water should be done as close to the time of the lab as possible. Letting an open bottle of seltzer water sit out too long can affect the concentration.
• Depending on time, you may be able to assign groups more than one buffer pH solution to be tested. Also, depending on your class, you can assign them the pH the day of the lab or you can give students the required pH when they are completing their homework assignment.
• Another variation of the lab is to have students make a buffer from only solutions of various concentrations of the acids and salts. If you would like to challenge your students with making a buffer from various solutions, you can prepare salt solutions with distilled or deionized water and label the concentrations for students to use in lab.
• Another extension is having students make a basic buffer with ammonium hydroxide and ammonium chloride. Be aware that ammonium chloride tends to absorb water due to its hygroscopic nature. This can cause the buffer’s pH to read differently than the calculated pH. If students perform multiple trials with the ammonium chloride salt and ammonia, you can work backward from the measured pH in the lab to find out how much water was absorbed by the ammonium chloride salt.

Teacher Tips

• 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. Once students turn in the homework answers and their procedure, check it for safety and accuracy before they implement it in the lab.

Further Extensions

Alignment to the Curriculum Framework for AP^® Chemistry 

Enduring Understandings and Essential Knowledge
Atoms are conserved in physical and chemical processes. (1E)
1E1: Physical and chemical processes can be depicted symbolically; when this is done, the illustration must conserve all atoms of all types.
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.

Matter can be described by its physical properties. The physical properties of a substance generally depend on the spacing between the particles (atoms, molecules, ions) that make up the substance and the forces of attraction among them. (2A)
2A3: Solutions are homogeneous mixtures in which the physical properties are dependent on the concentration of the solute and the strengths of all interactions among the particles of the solutes and solvent.

Chemical changes are represented by a balanced chemical equation that identifies the ratios with which reactants react and products form. (3A)
3A1: A chemical change may be represented by a molecular, ionic, or net ionic equation.
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 is a dynamic, reversible state in which rates of opposing processes are equal. (6A)
6A1: In many classes of reactions, it is important to consider both the forward and reverse reaction.
6A3: When a system is at equilibrium, all macroscopic variables, such as concentrations, partial pressures, and temperature, do not change over time. Equilibrium results from an equality between the rates of the forward and reverse reactions, at which point Q = K.
6A4: The magnitude of the equilibrium constant, K, can be used to determine whether the equilibrium lies toward the reactant side or product side.

Systems at equilibrium are responsive to external perturbations, with the response leading to a change in the composition of the system. (6B)
6B1: Systems at equilibrium respond to disturbances by partially countering the effect of the disturbance (Le Chatelier’s principle).
6B2: A disturbance to a system at equilibrium causes Q to differ from K, thereby taking the system out of the original equilibrium state. The system responds by bringing Q back into agreement with K, thereby establishing a new equilibrium state.

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.
6C2: The pH is an important characteristic of aqueous solutions that can be controlled with buffers. Comparing pH to pK_a allows one to determine the protonation state of a molecule with a labile proton.

Learning Objectives
1.17 The student is able to express the law of conservation of mass quantitatively and qualitatively using symbolic representations and particulate drawings. 
1.18 The student is able to apply conservation of atoms to the rearrangement of atoms in various processes.
2.2 The student is able to explain the relative strengths of acids and bases based on molecular structure, interparticle forces, and solution equilibrium.
2.8 The student can draw and/or interpret representations of solutions that show the interactions between the solute and solvent.
3.1 Students can translate among macroscopic observations of change, chemical equations, and particle views.
3.2 The student can translate an observed chemical change into a balanced chemical equation and justify the choice of equation type (molecular, ionic or net ionic) in terms of utility for the given circumstances.
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.
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.
6.1 The student is able to, given a set of experimental observations regarding physical, chemical, biological, or environmental processes that are reversible, construct an explanation that connects the observations to the reversibility of the underlying chemical reactions or processes.
6.7 The student is able, for a reversible reaction that has a large or small K, to determine which chemical species will have very large versus very small concentrations at equilibrium.
6.8 The student is able to use Le Chatelier’s principle to predict the direction of the shift resulting from various possible stresses on a system at chemical equilibrium.
6.9 The student is able to use Le Chatelier’s principle to design a set of conditions that will optimize a desired outcome, such as product yield.
6.10 The student is able to connect Le Chatelier’s principle to comparison of Q to K be explaining the effects of the stress on Q and K.
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.
6.16 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.
6.18 The student can design a buffer solution with a target pH and buffer capacity by selecting an appropriate conjugate acidbase pair and estimating the concentrations needed to achieve the desired capacity.
6.19 The student can relate the predominant form of a chemical species involving labile proton (i.e., protonated/deprotonated form of a weak acid) to the pH of a solution and the pK_a associated with the labile proton.
6.20 The student can identify a solution as being a + solution and explain the buffer mechanism in terms of the reactions that would occur on addition of acid or base.

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.
2.3 The student can estimate numerically 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.
6.4 The student can make claims and predications about natural phenomena based on scientific theories and models.
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

1. For the above table:
   1. Write the reaction with water for each of the chemicals named.
      {12390_PreLabAnswers_Reaction_1}
   2. Label the acid and conjugate base for each of the reactions.

      See above.

   3. Calculate the pK_a for each of the reactions for all the K values in the table.
      {12390_PreLabAnswers_Table_3}
2. For a buffer containing acetic acid and sodium acetate:
   1. Calculate the pH in a solution with 0.52 moles acetic acid and 0.26 moles of sodium acetate.

      [H₃O⁺] = K_a x [HA]/[A^–]
      [H₃O⁺] =(1.8 x 10^–5) x [0.52 moles]/[0.26 moles] = 3.6 x 10^–5
      pH = –log(3.6 x 10^–5) = 4.44

   2. Draw a particulate diagram of the compounds and ions in solution that are represented in the solution in Question 2a. Use A^– to represent the acetate ion.
      {12390_PreLabAnswers_Figure_1}

      There should be twice as many acetic acid molecules as acetate ions. For example, in this answer, there are 6 acetic acid molecules for 3 acetate ions and 3 sodium ions.

3. To create a 1.00-L buffer with a pH of 9.45, how many grams of ammonium chloride must be added to 1.00 L of 0.100 M ammonia?

   [OH^–] = K_b x [B]/[BH⁺]
   pOH = 14 – 9.45 = 4.55
   [OH^–] = 10^–4.55 = 2.82 x 10^–5 M
   [OH^–] = K_b × [B]/[BH+]
   2.82 x 10^–5 M = 1.76 x 10^–5 x [0.1 M]/[BH⁺]
   [BH⁺] = 0.0624 M
   Moles of NH₄Cl needed = 0.0624 M x 1 L = 0.0624 moles NH₄Cl
   Mass of NH₄Cl = 0.0624 moles NH₄Cl x 53.49 g/mol = 3.34 g NH₄Cl
   3.34 grams of ammonium chloride would have to be added to the 1.00 L of 0.100 M ammonium hydroxide to create a buffer with a pH of 9.45.

4. When 30.0 mL of 0.1 M sodium acetate is mixed with 20.0 mL of 0.2 M acetic acid, what is the pH of the buffering solution?

   New molarities are:
   (30.0 mL) x (0.1 M sodium acetate)/(50 mL) = 0.06 M sodium acetate
   (20.0 mL) x (0.2 M acetic acid)/(50 mL) = 0.08 M acetic acid
   
   [H₃O⁺] = K_a x [HA]/[A^–]
   [H₃O⁺] = (1.8 x 10^–5) x [0.08]/[0.06] = 2.4 x 10^–5
   pH = –log(2.4 x 10^–5) = 4.62

5. The major buffer in blood is composed of the weak acid, carbonic acid (H₂CO₃), and its conjugate base, bicarbonate ion (HCO₃⁻). The normal pH of blood is 7.2−7.4, which is very far removed from the pK_a value. The pH is kept in check by the lungs, which remove CO₂ via exhalation, and by the kidneys, which excrete acid (H₃O⁺) in the urine. People with impaired lung function are not able to exchange carbon dioxide efficiently between the lungs and air. The result is an increase in the amount of CO₂ dissolved in the blood.
   1. How does this affect the buffer balance in the blood?

      Increasing the amount of dissolved CO₂ in blood will shift the equilibrium to contain more of the acid component, H₂CO₃.

      {12390_PreLabAnswers_Reaction_6}
   2. Which term, respiratory acidosis or respiratory alkalosis, would better describe the resulting condition?

      The pH of blood will decrease—this is called respiratory acidosis.

6. Why does a mixture of HCl and NaCl not produce a buffer?

   The conjugate base of the strong acid HCl is the chloride anion. Chloride ion will not react with any excess strong acid added to it in solution. In a buffer the acid component neutralizes excess base and the basic component neutralizes excess acid.

7. In the lab, you will have to create a buffer with a specific pH that will be tested by your instructor. Write a general step-by-step procedure to show how to calculate and prepare your assigned buffer. In the lab, you will be given the following chemicals:
   • Acetic acid solution, HC2H3O2, 0.1 M
   • Citric acid solution, H3C6H5O7, 0.1 M
   • Seltzer water, carbonic acid, H2CO3, assume 0.07 M
   • Sodium acetate trihydrate, solid, NaC2H3O2•3H2O
   • Sodium bicarbonate, solid, NaHCO3
   • Sodium dihydrogen citrate, solid, NaH2C6H5O7

   Each group will be allowed 50 mL of one of the above solutions, and at least 1.5 grams of each solid. Your teacher will assign a buffer with a specific pH for you to prepare for testing. Once your solution is made, your instructor will test it with their pH meter and let you know if you met your target.

   1. Think safety, first. Make sure you have the proper personal protective equipment (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. Show the necessary calculations and draw the necessary data tables in your notebook for data collection during the lab.
   5. Draw the setups you will be using in the lab. Label all equipment.

Sample Data

Assigned target pH = 4.55

1. Receive assigned buffer pH from instructor and record it in a data table.
2. Depending on the pH, measure 50 mL of the corresponding 0.1 M acid solution, and place it in a 100 mL beaker. a. If the pH is near 3.13, citric acid will be used. b. If the pH is near 4.74, acetic acid will be used. c. If the pH is near 6.37, carbonic acid (seltzer water) will be used.
3. Calculate the amount of solid salt necessary to create the buffer with the following equations.

   [H₃O⁺] = 10^–pH
   [H₃O⁺] = K_a x [HA]/[A–]

   For our assigned pH of 4.55,

   [H₃O⁺] = 10^–4.55 = 2.82 x 10^–5 M
   [H₃O⁺] = K_a x [HA]/[A^–]
   2.82 x 10^–5 M = (1.8 x 10^–5) x [0.1 M]/[A^–]
   [A^–] = 0.0639 M

   The amount of acetate ion needed in 50 mL of the buffer will be:

   0.0639 M x 0.05 L = 0.00319 moles of sodium acetate trihydrate

   0.00319 moles of sodium acetate trihydrate x (136.08 g/mol) = 0.434 g of sodium acetate trihydrate

4. Carefully mass the amount of salt necessary in a weigh boat, and add to the 50 mL of solution in your 100 mL beaker.
5. Stir the solution until all solid is dissolved.
6. Have your instructor measure your buffer with the pH meter.

Additional Data

References

College Board, The. 2014. “AP Chemistry Course and Exam Description, rev. ed.” NY: The College Board. Accessed September 19th, 2017. http://media.collegeboard.com/digitalServices/pdf/ap/ap-chemistry-course-and-exam-description.pdf

CRC Handbook of Chemistry and Physics, 95th ed. CRC Press: Boca Raton, FL, 2014

Introduction

One of the most important applications of acids and bases in chemistry and biology is that of buffers. A buffer solution resists rapid changes in pH when acids and bases are added to it. Every living cell contains natural buffer systems to maintain the constant pH needed for proper cell function. Many consumer products are also buffered to safeguard their activity. How do buffers maintain the delicate pH balance needed for life and health?

Concepts

• Acids and bases
• Conjugate acid–base pair
• Buffer
• pH

Background

The ability of buffers to resist changes in pH when acid or base is added is a result of their chemical composition. All buffers contain a mixture of a conjugate acid–base pair; either a weak acid (HA) and its conjugate base (A^–) or a weak base (B) and its conjugate acid (BH⁺). Weak acids and weak bases both dissociate slightly in water (Reactions 1 and 2).

These reactions are reversible. Both the weak acid and its conjugate base, or the weak base and its conjugate acid, are present in solution. The equilibrium constant expressions for these dissociation reactions are: Buffers control pH because the two buffering components, either HA and A^– or B and BH⁺, are able to neutralize both acids and bases added to the solution. The actual pH of a buffer solution depends on the concentration of the conjugate acid–base pair in solution. If Equation 1 is rearranged, the concentration of hydronium ions in solution is: and the pH is: If the concentrations of the acid–base pair are equal, [HA] = [A^–], then the pH of the buffer is equal to the pK_a because log(1) = 0. By varying the amounts of HA and A^– in solution, the pH of the buffer solution can be changed.

For a buffer made up of a weak base (B) and its conjugate acid (BH⁺), the solution pH calculations are similar. If Equation 2 is rearranged, the concentration of hydroxide ions (OH^–) in solution is: and the pOH is: At room temperature, if pOH is known, then pH can be calculated using Equation 7: Once the buffer is made, how does the pH remain constant when a strong acid or base is added? Let’s look at an example.

Acetic acid is a weak acid, with K_a equal to 1.8 x 10^–5. If a buffer solution is made with 0.5 moles of acetic acid and 0.5 moles of its conjugate base sodium acetate, the initial pH of the solution will be equal to the pK_a, 4.74. Now, if 0.05 moles of a strong acid is added to the buffer, the H₃O⁺ will react with 0.05 moles of the sodium acetate to form 0.05 moles of acetic acid. This produces a solution with 0.55 moles of acetic acid and 0.45 moles of its conjugate base sodium acetate. If the solution volume change is slight, then the new pH of the solution is: The pH difference is only 0.09 units.

For buffers to be effective, noticeable amounts of both parts of the conjugate acid–base pair must be present in solution. This limits the concentration ratios for HA:A^– or B:BH⁺ to between 10:1 and 1:10 and the pH range for the buffering action of any weak acid to pK_a ±1. An ideal buffer is a solution that contains equal moles of the conjugate acid–base pair.

Experiment Overview

The purpose of this activity is to complete the Prelab Homework Assignment prior to lab to promote conceptual understanding of acids, bases and buffers. You will first review and analyze buffers. After completing the homework assignment, your group will be assigned a specific pH by your instructor. Depending on the pH you are assigned, you will have to prepare the buffer that will match the provided pH. See how close your group can be! You’ll love the challenge!

Prelab Questions

Complete the following homework set, and write a lab procedure to be approved by your instructor prior to performing the lab. Along with your procedure, you will turn in any graphs, tables or figures you were asked to create in this homework set and answers to the questions. Use a separate sheet of paper if needed.

1. For Table 1:
   1. Write the reaction with water for each of the chemicals named.
   2. Label the acid and conjugate base for each of the reactions.
   3. Calculate the pK_a for each of the reactions for all the K values in the table.
2. For a buffer containing acetic acid and sodium acetate:
   1. Calculate the pH in a solution with 0.52 moles acetic acid and 0.26 moles of sodium acetate.
   2. Draw a particulate diagram of the compounds and ions in solution that are represented in the solution in Question 2a. Use A^– to represent the acetate ion.
3. To create a 1.00-L buffer with a pH of 9.45, how many grams of ammonium chloride must be added to 1.00 L of 0.100 M ammonia?
4. When 30.0 mL of 0.1 M sodium acetate is mixed with 20.0 mL of 0.2 M acetic acid, what is the pH of the buffering solution?
5. The major buffer in blood is composed of the weak acid, carbonic acid (H₂CO₃), and its conjugate base, bicarbonate ion (HCO₃⁻). The normal pH of blood is 7.2−7.4, which is very far removed from the pK_a value. The pH is kept in check by the lungs, which remove CO₂ via exhalation, and by the kidneys, which excrete acid (H₃O⁺) in the urine. People with impaired lung function are not able to exchange carbon dioxide efficiently between the lungs and air. The result is an increase in the amount of CO₂ dissolved in the blood.
   1. How does this affect the buffer balance in the blood?
   2. Which term, respiratory acidosis or respiratory alkalosis, would better describe the resulting condition?
6. Why does a mixture of HCl and NaCl not produce a buffer?
7. In the lab, you will have to create a buffer with a specific pH that will be tested by your instructor. Write a general step-by-step procedure to show how to calculate and prepare your assigned buffer. In the lab, you will be given the following chemicals:
   • Acetic acid solution, HC₂H₃O₂, 0.1 M
   • Citric acid solution, H₃C₆H₅O₇, 0.1 M
   • Seltzer water, carbonic acid, H₂CO₃, assume 0.07 M
   • Sodium acetate trihydrate, solid, NaC₂H₃O₂•3H₂O
   • Sodium bicarbonate, solid, NaHCO₃
   • Sodium dihydrogen citrate, solid, NaH₂C₆H₅O₇

   Each group will be allowed 50 mL of one of the above solutions and at least 1.5 grams of each solid. Your teacher will assign a buffer with a specific pH for you to prepare for testing. Once your solution is made, your instructor will test it with a pH meter and let you know if you met your target.

   1. Think safety first. Make sure you have the proper personal protective equipment (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. Show the necessary calculations and draw the necessary data tables in your notebook for data collection during the lab.
   5. Draw the setups you will be using in the lab. Label all equipment.

Safety Precautions

Dilute acid and base solutions are skin and eye irritants. Acetic acid solutions may be irritating to the respiratory tract. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Avoid exposure of all chemicals to eyes and skin, and notify the teacher of any spills. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.