Intro Chem II - 1.1 Buffer Solutions BLANK PDF

1.1 Buffer Solutions Dr. Jamie Young Learning Objectives Calculate the pH of a buffer solution Calculate the pH change when acid or base is added to a buffer solution Specify the composition of a buffer solution with a given pH Understand buffer “effectiveness”...

1.1 Buffer Solutions Dr. Jamie Young Learning Objectives Calculate the pH of a buffer solution Calculate the pH change when acid or base is added to a buffer solution Specify the composition of a buffer solution with a given pH Understand buffer “effectiveness” 2 Buffers Buffers are solutions that resist (small) changes in pH when an acid or base is added. They act by neutralizing excess acid or base that is added to the buffered solution. Blood has a mixture of H2CO3 and HCO3− Pharmaceuticals can be modified using buffers Shampoo is a buffer of citric acid and NaOH Introductory Chemistry II 3 Making a Buffer Introductory Chemistry II 4 How Acid Buffers Work: Adding Base Buffers work by applying Le Châtelier’s Principle to weak acid equilibrium. Buffers contain significant amounts of the weak acid molecules, HA. These molecules react with added base to neutralize it: you can also think of the H3O+ combining with the OH− to make H2O; the H3O+ is then replaced by the shifting equilibrium. Introductory Chemistry II 5 How Acid Buffers Work: Adding Acid HA(aq) + H2O(l) ⇌ A−(aq) + H3O+(aq) The buffer solution also contains significant amounts of the conjugate base anion, A−. These ions react with the added acid to make HA: After the equilibrium shifts, the concentration of H3O+ (H+) is kept constant. Introductory Chemistry II 6 Common Ion Effect HA(aq) + H2O(l) ⇌ A−(aq) + H3O+(aq) Adding a salt containing the anion NaA, which is the conjugate base of the acid (the common ion), shifts the position of equilibrium to the left. This causes the pH to be higher than the pH of the acid solution: Introductory Chemistry II 7 Practice: What is the pH of a buffer that is 0.100 M CH3COOH and 0.100 M CH3COONa? 1. Write the reaction for the acid with water. [HA] [A−] [H3O+] 2. Construct an ICE table for Initial the reaction. Change 3. Enter the initial Equilibrium concentrations – assuming the [H3O+] from water is ≈ 0. Introductory Chemistry II 8 Practice: What is the pH of a buffer that is 0.100 M CH3COOH and 0.100 M CH3COONa? CH3COOH + H2O ⇌ CH3COO− + H3O+ 4. Represent the change in the concentrations in terms [HA] [A−] [H3O+] of x. Initial 0.100 0.100 ≈0 5. Sum the columns to find Change the equilibrium concentrations in terms of x. Equilibrium 6. Substitute into the equilibrium constant expression. Introductory Chemistry II 9 Practice: What is the pH of a buffer that is 0.100 M CH3COOH and 0.100 M CH3COONa? Ka for CH3COOH = 7. Determine the value of Ka. [HA] [A−] [H3O+] Initial 0.100 0.100 ≈0 8. Due to Ka being small, Change -x +x +x [HA]eq = [HA]i and [A-]eq = [A-]i, then solve for x. Equilibrium 0.100-x 0.100+x x Introductory Chemistry II 10 Practice: What is the pH of a buffer that is 0.100 M CH3COOH and 0.100 M CH3COONa? Ka for CH3COOH = 1.8 x 10−5 9. Check if the approximation is valid seeing [HA] [A−] [H3O+] if x < 5% of [CH3COOH]i. Initial 0.100 0.100 ≈0 Change -x +x +x Equilibrium 0.100-x 0.100+x x Introductory Chemistry II 11 Practice: What is the pH of a buffer that is 0.100 M CH3COOH and 0.100 M CH3COONa? 10. Substitute x into the [HA] [A−] [H3O+] equilibrium concentration Initial 0.100 0.100 ≈0 definitions and solve. Change -x +x +x Equilibrium 0.100-x 0.100+x x Introductory Chemistry II 12 Practice: What is the pH of a buffer that is 0.100 M CH3COOH and 0.100 M CH3COONa? 11. Substitute [H3O+] into the [HA] [A−] [H3O+] formula for pH and solve. Initial 0.100 0.100 ≈0 Change -x +x +x Equilibrium Introductory Chemistry II 13 Common Ion Effect Introductory Chemistry II 14 Practice: What is the pH of a buffer that is 0.100 M CH3COOH and 0.100 M CH3COONa? 12. Check by substituting [HA] [A−] [H3O+] the equilibrium Initial 0.100 0.100 ≈0 concentrations back into the equilibrium constant Change -x +x +x expression and comparing Equilibrium 0.100 0.100 1.8x10-5 the calculated Ka to the given Ka. Introductory Chemistry II 15 Henderson-Hasselbalch Equation Calculating the pH of a buffer solution can be simplified by using an equation derived from the Ka expression called the Henderson-Hasselbalch equation. The equation calculates the pH of a buffer from the pKa and initial concentrations of the weak acid and salt of the conjugate base as long as the “x is small” approximation is valid Introductory Chemistry II 16 Deriving the HH Equation Introductory Chemistry II 17 Do I Use the Full Analysis or HH? The Henderson-Hasselbalch equation is generally good enough when the “x is small” approximation is applicable Generally, this approximation will work when both of the following are true: a) b) For most problems, this means that the initial acid and salt concentrations should be over 100x to 1000x larger than the value of Ka Introductory Chemistry II 18 Example: What is the pH of a buffer that is 0.050 M C6H5COOH and 0.150 M C6H5COONa? 1. Assume the [HA] and [A−] equilibrium concentrations are the same as the initial. 2. Substitute into the Henderson-Hasselbalch Equation 3. check the “x is small” approximation Introductory Chemistry II 19 Practice: What is the pH of a buffer that is 0.14 M HF (Ka = 7.08 x 10-4) and 0.071 M KF? 1. Find the pKa from the given Ka. 2. Assume the [HA] and [A−] equilibrium concentrations are the same as the initial. 3. Substitute into the Henderson-Hasselbalch Equation. 4. Check the “x is small” approximation. Introductory Chemistry II 20 Basic Buffers B:(aq) + H2O(l) ⇌ H:B+(aq) + OH−(aq) Weak base Conjugate acid Buffers can also be made by mixing a weak base, (B:), with a soluble salt of its conjugate acid, H:B+Cl− Buffer solution Introductory Chemistry II 21 HH Equation for Basic Buffers The Henderson-Hasselbalch equation is written for a chemical reaction with a weak acid reactant and its conjugate base as a product The chemical equation of a basic buffer is written with a weak base as a reactant and its conjugate acid as a product: To apply the Henderson-Hasselbalch equation, the chemical equation of the basic buffer must be looked at like an acid reaction: ✓ this does not affect the concentrations, just the way we are looking at the reaction Introductory Chemistry II 22 Relationship between pKa and pKb Just as there is a relationship between the Ka of a weak acid and Kb of its conjugate base, there is also a relationship between the pKa of a weak acid and the pKb of its conjugate base: Introductory Chemistry II 23 Example: What is the pH of a buffer that is 0.50 M NH3 (pKb = 4.75) and 0.20 M NH4Cl? 1. Find the pKa of the conjugate acid (NH4+) from the given pKb 2. Assume the [B] and [HB+] equilibrium concentrations are the same as the initial 3. Substitute into the Henderson-Hasselbalch equation 4. Check the “x is small” approximation Introductory Chemistry II 24 HH Equation for Basic Buffers The Henderson-Hasselbalch equation is written for a chemical reaction with a weak acid reactant and its conjugate base as a product The chemical equation of a basic buffer is written with a weak base as a reactant and its conjugate acid as a product: We can rewrite the Henderson-Hasselbalch equation for the chemical equation of the basic buffer in terms of pOH: Introductory Chemistry II 25 Does the pH of a Buffer Change? Though buffers do resist change in pH when acid or base is added to them, their pH does change. Calculating the new pH after adding acid or base requires breaking the problem into two parts: 1. a stoichiometry calculation for the reaction of the added chemical with one of the ingredients of the buffer to reduce its initial concentration and increase the concentration of the other ✓ added acid reacts with the A− to make more HA ✓ added base reacts with the HA to make more A− 2. an equilibrium calculation of [H3O+] using the new initial values of [HA] and [A−] Introductory Chemistry II 26 Buffering Effectiveness A good buffer should be able to neutralize moderate amounts of added acid or base. Of course, there is a limit to how much can be added before the pH changes significantly. The buffer capacity is The buffer range is Introductory Chemistry II 27 Buffer Capacity Buffer capacity is the amount of acid or base that can be added to a buffer without causing a large change in pH. Buffers intended to work with added acid generally have [base] > [acid]. Buffers intended to work with added base generally have [acid] > [base]. Introductory Chemistry II 28 Buffer Range A buffer will be effective when: Substituting into the Henderson-Hasselbalch equation we can calculate the maximum and minimum pH at which the buffer will be effective Lowest pH Highest pH When choosing an acid to make a buffer, choose one whose pKa is closest to the pH of the buffer. Introductory Chemistry II 29 iClicker Which of the following acids would be the best choice to combine with its sodium salt to make a buffer with pH 4.25? a. Chlorous acid, HClO2 pKa = 1.95 b. Nitrous acid, HNO2 pKa = 3.34 c. Formic acid, HCHO2 pKa = 3.74 d. Hypochlorous acid, HClO pKa = 7.54 Introductory Chemistry II 30 What ratio of NaCHO2 : HCHO2 would be required to make a buffer with pH 4.25? Formic acid, HCHO2 pKa = 3.74 Introductory Chemistry II 31 Effectiveness of Buffers A buffer will be most effective when A buffer will be effective when A buffer will be more effective when A buffer is generally most effective when Introductory Chemistry II 32