pH of Buffered Solutions

Buffer Solution

• A buffered solution containing 5.0M HC2H3O2 and 3.0M NaC2H3O2 will have the same pH as one containing 0.05M HC2H3O2 and 0.03M NaC2H3O2.
• The pH of a buffered solution is dependent on the ratio of the concentrations of the acid and its conjugate base, not their individual concentrations.

Equilibrium Reaction of NH3

• NH3 + H2O ↔ NH4+ + OH-
• The equilibrium constant Kb is calculated as Kb = [(NH4+) × (OH-)]/ NH3 or Kb = (OH-)2/ NH3

Neutralization of NH3

• The volume of HCl required for complete neutralization of NH3 is calculated using the equation (M × V)NH3 = (M × V)HCl
• For example, 50 ml of HCl is required to neutralize 50 ml of 0.1M NH3

Stoichiometric Calculations

• Before the equivalence point, after adding 10 ml of HCl, the major species present are NH3, NH4+, Cl-, and H2O
• The dominant equilibrium is the dissociation of the weak acid NH4+ with Ka = Kw/Kb(for NH3)

pH at Equivalence Point

• The pH at the equivalence point is calculated using the equilibrium reaction B- + H2O ↔ HB + OH-
• The equilibrium constant Kb is calculated as [HB] [OH-] = Kw = 1.0 × 10-14
• For example, the pH at the equivalence point is 8.85 when 60.0 ml of 0.1M NaOH is added

Buffer Solutions

• Near the equivalence point, [H+] is relatively small and the addition of a small amount of OH- produces a large change in pH.
• At an intermediate pH (5), pKa = 5 because HIn is half-neutralized, [HIn] = [In-], and the color is orange.

Henderson-Hasselbalch Equation

• The equation is useful for calculating the pH of solutions when the ratio of [A-]/[HA] is known.
• The log form of the expression for Ka is called the Henderson-Hasselbalch equation.

Acid-Base Reaction

• The major species present (before any reaction occurs) are NH3, H+, Cl-, and H2O.
• NH3 reacts with H+ from the added HCl to form NH4+: NH3(aq) + H+(aq) → NH4+(aq).
• This reaction proceeds to completion as NH3 readily reacts with free protons.

pH Calculation

• The pH at the equivalence point of a weak acid with a strong base is always greater than 7.

pH Indicators

• pH indicators exhibit different colors depending on the presence or absence of a proton (H+ ion) attached to the molecule
• At low pH, the proton is attached, and the indicator appears red, with a ratio of [HIn] to [In-] of 10:1
• At high pH, the proton is absent, and the indicator appears yellow, with a ratio of [HIn] to [In-] of 1:10
• The minimum change in pH (∆pH) required to cause a color change from red to yellow is 2 units

Choice of Indicator for Titration

• When choosing an indicator for a titration, select one that changes color at approximately the pH at the equivalence point of the titration
• The indicator range is the pH range over which the indicator changes color, in this case, pH 4-6

Titration Calculations

• Molarity (M) is used to calculate the number of moles of acid or base
• The number of moles of acid or base is equal to the number of moles of H+ or OH- ions
• The equation to calculate the number of moles is: M × V × n = M × V × n

Example Titration

• A 10 mL sample of NaOH is titrated against 0.1M H2SO4, with the endpoint reached after adding 9 mL of H2SO4
• The balanced chemical equation for the reaction is: 2NaOH + H2SO4 → Na2SO4 + 2H2O
• The molarity of NaOH can be calculated using the equation: M × V × n = M × V × n

Titration of a Strong Acid with a Strong Base

• A strong acid (HCl) reacts with a strong base (NaOH) to form a salt (NaCl) and water
• The pH of the solution changes during the titration, with a sharp increase in pH at the equivalence point

Acid-Base Titration

• 60 ml of 0.1 M HCl is added to NH3, resulting in the reaction NH3 + HCl → NH4+ + H2O
• Initially, there are 5.0 mmol of NH3 and 6.0 mmol of HCl, with 0 mmol of NH4+
• After the reaction, there are 0 mmol of NH3, 0 mmol of HCl, and 5 mmol of NH4+
• The concentration of H+ is 1 mmol H+ in excess per 110 ml of solution, with a pH of -log [H+] = -log 0.009

Titration Curve of Weak Base with Strong Acid

• The titration curve is different from a weak acid-strong base titration curve in four major ways

Henderson-Hasselbalch Equation

• pH = pKa + log (salt/acid)
• Example: pH = 3.85 + log (0.25/0.75) = 3.38

Experiment 4: Acid-Base Titration

• Determination of the concentration of sodium carbonate and sodium bicarbonate in a mixture

Procedure

• Washing the burette with distilled water and then with HCl
• Washing the pipette with distilled water and then with the mixture
• Washing the flask with distilled water

Equilibrium and Dissociation

• The dominant equilibrium is the dissociation of the weak acid NH4+ with Ka = Kw/Kb (for NH3)
• pH values: 7 for strong acids, greater than 7 for weak acids