Weak Acids, Weak Bases, and Salts

Weak Acids, Weak Bases, and Salts Introduction: Weak acids and weak bases only dissociate to a small degree (typically less than 5%). Weak acids vary in their tendency to ionize in water. A weak acid that ionizes to a greater extent will form more H3O+ ions and thus will have a larger Ka (acid ionization constant) value than a weak acid that ionizes to a smaller extent. By comparing the Ka values for different weak acids, the relative strength of weak acids can be determined. Weak bases vary in their relative ability to accept protons. Thus, a weak base with a larger Kb (base ionization constant) value is stronger than a weak base with a smaller Kb. A salt is the ionic compound formed during an acid-base neutralization reaction. Almost all salts are strong electrolytes (soluble in water) and thus exist as ions in aqueous solutions. Salt solutions can be neutral, acidic, or basic depending on the behavior of the component ions. Cations from the strong bases (group IA and group IIA) and anions from the monoprotic strong acids (Cl-, Br-, I-, NO3- and ClO4-) will not react with water and are neutral ions. Salt solutions that contain only these ions will be neutral. For example, NaCl and SrBr2 are neutral salts. Cations that are derived from weak bases can react with water (hydrolyze) to form acidic solutions. For example, NH4+ is the conjugate acid of the weak base NH3, so NH4+ acts as a weak acid and donates a proton to water to produce H3O+ ions: NH4+ (aq) + H2O(l) ⇌ H3O+(aq) + NH3(aq) On the other hand, anions that are derived from weak acids tend to form basic solutions. For example, CNis the conjugate base of the weak acid HCN. Thus, CN- acts as a weak base and accepts a proton from water to produce OH- ions: CN- (aq) + H2O(l) ⇌ HCN(aq) + OH-(aq) The pH values for various acidic, basic, and salt solutions will be measured in this experiment. For each acidic solution, you will use the pH to calculate the equilibrium concentration of the hydronium ions, [H3O+]eq. The initial concentration of the weak acid, [HA]i, and the [H3O+]eq will then be used in an ICE table to calculate a Ka value for each acidic solution. For each basic solution, [OH-]eq will need to be calculated and used in an ICE table to determine a Kb value. Percent dissociation values will also be calculated for all six solutions. Acid ionization constants will be used to compare the relative strength of the acidic solutions; base ionization constants will be used to compare the relative strength of the basic solutions. Refer to Sections 14.1-14.4 of Openstax Chemistry for information on defining Bronsted-Lowry Acids and Bases, calculating [H3O+] from pH, relative strengths of acids and bases, and hydrolysis of salt solutions. Equations to use for calculations: + - Acid Dissociation: HA(aq) + H2O(l) ⇌ H3O (aq) + A (aq) + pH = - log [H3O ]eq + [H3O ] = 10 -pH Ka = [H 3 O + ]eq [A − ] eq % dissociation = [HA] eq [H 3 O + ]eq [HA]i  100 page 1 of 9 Base Dissociation: B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq) pOH = - log [OH-]eq [OH-] = 10-pOH Kb = [BH + ]eq [OH − ]eq [B] eq % dissociation = [OH − ]eq [B]i  100 pH + pOH = 14.00 Materials: 6 – 30 mL beakers 1 – 250 mL beaker Chromebook GoLink pH probe DI water KimWipes CH3COOH(aq) NH4Cl(aq) HC2H3O3(aq) NaCH3COO(aq) NH3(aq) Na2CO3(aq) Note: Handle the pH probe carefully to avoid breakage fees! Procedure: 1. Obtain about 15 mL of each of the following solutions in labelled 30 mL beakers: acetic acid, ammonium chloride, glycolic acid, sodium acetate, ammonia, and sodium carbonate. Refer to the “Transferring liquids” technique. 2. Assemble the Chromebook, pH probe, and GoLink system as directed in the “Using and calibrating a pH probe” technique. In the Vernier Graphical Analysis window, click the 3-square icon in the upper right corner to select “Meter”. 3. Use the 250 mL beaker to collect the pH probe rinsings, and rinse the pH probe with deionized water every time it is transferred from one solution to another. Measure the pH of all six solutions being careful not to contaminate solutions. 4. Using pH paper, test the pH of each of the six solutions. Record your results. 5. Have your instructor double-check your pH readings before you put away the computer equipment. When you are done with the pH sensor, rinse it with deionized water and place in the sensor storage container. Clean and return your glassware. Note: The solutions we test will generally be more acidic than we would predict them to be. A major reason for this increased acidity is the presence of dissolved CO2 in aqueous solutions. CO2 reacts with water to generate H+: CO2(g) + H2O(l) ⇌ H+(aq) + HCO3-(aq) The solubility of CO2 is greatest in basic solutions and least in acidic ones so basic solutions will be affected more than acidic solutions. Waste Disposal: Discard your solutions in the waste container in the fume hood. Clean-Up: Rinse all glassware with lots of tap water and then a final rinse with DI water. Wipe your entire bench with a damp paper towel. Put all equipment back where you found it. page 2 of 9 Weak Acids, Weak Bases, and Salts Pre-Lab Questions and Calculations You will complete this quiz in Canvas 1 hour before your lab period. This page will not be turned in or graded. You may use this page to set up your calculations before you take the Canvas quiz. 1. Which of the solutions used in lab this week are predicted to be acidic? 2. Which of the solutions used in lab this week are predicted to be basic? 3. Acidic ions tend to have a pH value that is _________ 4. If the pH of a 0.45 M weak acid is measured to be 4.32, what is its Ka (acid ionization constant) value? 5. Which one of the following salts is predicted to be basic? a. b. c. d. KCl CaBr2 LiNO2 NH4I 6. Why is KF a basic salt? 7. A student accidentally takes too much solution for this week's lab. What should the student do with the excess solution? 8. Why is it important to rinse the pH probe with DI water after taking a pH measurement? page 3 of 9 Name: _____________________________ Partners: ________________________ Weak Acids, Weak Bases, and Salts Lab Report Turn in Pages 4-6 as your Graded Lab Report. Data and Calculations Table 1: Concentrations and pH values of solutions Original Solutions Concentration pH probe value pH paper color/value Instructor Initials: _________ Calculations: Acidic solution: Show a sample calculation for: a) all equilibrium concentrations ([HA]eq, [H3O+]eq, and [A-]eq) ; b) Ka ; and c) percent dissociation for one acidic solution. Include a sample ICE table showing how you calculated your equilibrium concentrations. Report your calculated values for the three acidic solutions in Results: Table 2. page 4 of 9 Basic solution: Show a sample calculation for: a) all equilibrium concentrations ([B]eq, [OH-]eq, and [BH+]eq); b) Kb ; and c) percent dissociation for one basic solution. Include a sample ICE table showing how you calculated your equilibrium concentrations. Report your calculated values for the three basic solutions in Results: Table 3. Results Table 2: Results of acidic solution calculations Acid pH [HA]i, M [H3O+]eq, M [A–]eq, M [HA]eq, M Ka % dissociation [BH+]eq, M [B]eq, M Kb % dissociation CH3COOH NH4Cl HC2H3O3 Table 3: Results of basic solution calculations Base pH [B]i, M [OH-]eq , M NaCH3COO NH3 Na2CO3 page 5 of 9 Conclusion: Summarize the trends for the pH readings, K values, and % ionization values relative to the strength of the acidic solutions. Do the same for the basic solutions. What was the strongest and weakest acidic solution and how did you determine this? What was the strongest and weakest basic solution and how did you determine this? page 6 of 9 Post-Lab Questions – These questions will not be graded as part of your lab report grade. You will be responsible for the information in these questions and able to answer these or similar questions on the post-lab quiz at the start of next week’s lab. 1. List your solutions (use chemical formulas) in order of increasing pH. Lowest pH ________________________________________________________________ Highest pH 2. Rank your acidic solutions (use chemical formulas) by increasing pH, Ka, and % dissociation: pH __________________________________________________ Ka __________________________________________________ % dissociation __________________________________________________ 3. Rank your basic solutions (use chemical formulas) by increasing pH, Kb, and % dissociation: pH __________________________________________________ Kb __________________________________________________ % dissociation __________________________________________________ 4. Identify what type of cation and anion (acidic, basic, or neutral) typically make up each type of salt below. Circle your answer for each. Acidic salt: cation: acidic basic neutral anion: acidic basic neutral Neutral salt: cation: acidic basic neutral anion: acidic basic neutral Basic salt: cation: acidic basic neutral anion: acidic basic neutral 5. For each of the 3 salt solutions used in lab: 1) write the dissociation equation; 2) label the cation as acidic, basic, or neutral; 3) label the anion as acidic, basic or neutral. NaF has been done as an example. e. g. Salt = NaF , dissociation equation: NaF(aq) → Na+(aq) + F-(aq) Salt 1 = _______ Salt 2 = _______ formula for cation: _Na+_ (circle one) acidic basic neutral formula for anion: _ F-_ (circle one) basic neutral acidic 1) dissociation equation: ______________________________________ 2) formula for cation: _________ (circle one) acidic basic neutral 3) formula for anion: _________ (circle one) acidic basic neutral 1) dissociation equation: ______________________________________ 2) formula for cation: _________ (circle one) acidic basic neutral 3) formula for anion: _________ (circle one) acidic basic neutral page 7 of 9 Salt 3 = _______ 1) dissociation equation: ______________________________________ 2) formula for cation: _________ (circle one) acidic basic neutral 3) formula for anion: _________ (circle one) acidic basic neutral 6. From what acid and what base were the salts below made? Ca(NO3)2: ____________________________ NH4Br: ______________________________ BaCl2: _______________________________ 7. Calculate the Ka value for a 0.010 M solution of acetic acid that is measured to have a pH of 3.38. Show all your work below! Calculation: a) How does this calculated value above compare to the Ka value of 1.8 x 10-5 reported in the textbook? b) State the Ka value you calculated for the 0.50 M CH3COOH solution in this experiment and compare it to the calculated value above. c) Should these Ka values be about the same or different? Explain why. 8. Discuss at least 2 sources of error in this experiment. How did they affect your results and how would you correct them if you were to repeat the experiment? 9. Complete the table below. Classify each of the following solutions as a weak acid, weak base, acidic salt, basic salt or neutral salt. substance type of compound substance K2CO3(aq) H3PO4(aq) NH4ClO3(aq) CaCl2(aq) CsF(aq) KNO3(aq) NH(CH3)2(aq) (an amine) Sr(ClO2)2(aq) type of compound CH3COOH(aq) page 8 of 9 10. Complete the table below for each solution used in this lab. NaF has been done as an example. Example: Sodium Fluoride (NaF) 0.50 M Solution F- Species Expected to Hydrolyze Na+ Spectator Ion (if none, write N/A) F-(aq) + H2O(l)  HF(aq) + OH-(aq) Net-Ionic Equation for Hydrolysis Kb = [HF][OH-] / [F-] Equilibrium Expression (specify Ka, Kb) Acetic Acid (CH3COOH) Ammonium chloride (NH4Cl) Glycolic Acid (HC2H3O3) Sodium acetate (NaCH3COO) Ammonia (NH3) Sodium Carbonate (Na2CO3) page 9 of 9