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Chapter 19 - Ionic Equilibria in Aqueous Systems

  • The sodium ion is a spectator ion since it does not interact with water (as shown in the image attached). According to Le Châtelier's principle, adding CH3COO ion shifts the equilibrium position to the left; consequently, [H3O+] drops, decreasing the degree of acid dissociation:
    • CH3COOH(aq) + H2O(l) ⥫⥬ H3O+(aq) + CH3COO− (aq; added)
  • When we add acetic acid to a sodium acetate solution instead of water, we obtain the same effect.
  • The previously existent acetate ion prevents the acid from dissociating as much as it does in water, keeping the [H3O+] lower (and pH higher). The result is less acid dissociation in either situation.
  • The ion acetate is known as the common ion because it is present in both acetic acid and sodium acetate solutions.
  • When a particular ion is introduced to an equilibrium mixture that already includes that ion, the point of equilibrium changes away from producing it.
  • The image attached above indicates that when the concentration of acetate ion (provided by dissolving sodium acetate) increases, the percent dissociation (and the [H3O+]) of an acetic acid solution drops.
  • As a result, the common ion, A, prevents the dissociation of HA, as it reduces the acidity of the solution (higher pH).
  • As a result, the buffer components absorb virtually all of the additional H3O+ or OH. To restate, as long as the quantity of additional H3O+ or OH is modest in comparison to the quantities of the buffer components, converting one component into the other results in a minor change in the buffer-component concentration ratio and, as a result, a tiny change in [H3O+] and pH.
  • The pH variations in Sample Problem 19.1 are generally quite modest. It is worth noting that the latter two sections of the issue have a stoichiometry component.
  • When H3O+ or OH is introduced to a buffered solution, the pH changes significantly less than in an unbuffered solution.
  • A buffer is made up of a weak acid and a conjugate base (or a weak base and a conjugate acid). To be effective, the component quantities must be substantially larger than the quantity of H3O+ or OH that has been added.
  • The pH is determined by the buffer-component concentration ratio; the ratio and the pH are related.
  • The Henderson-Hasselbalch equation connects them. When H3O+ or OH is introduced to a buffer, one component interacts to produce the other, resulting in the formation of the buffer. [H3O+] (and pH) barely minimally alter.
  • A concentrated (high capacity) buffer has lower pH fluctuations than a dilute buffer. When H3O+ or OH is introduced to a buffer, one component interacts to create the other, resulting in just a small change in [H3O+] (and pH). A concentrated (high capacity) buffer has lower pH fluctuations than a dilute buffer.
  • The buffer has the greatest capacity when the pH of the buffer equals the pKa of the acid component.
  • A buffer's effective pH range is pKa 1 pH unit. To make a buffer, select a conjugate acid-base pair and determine the component ratio. Calculate the buffer concentration and adjust the final solution to the desired pH.
  • The pH of a strong acid–strong base titration begins low, steadily increases, and suddenly rockets up towards the equivalence point (pH = 7).
  • In a weak acid–strong base titration, the pH begins higher and gradually rises in the buffer zone (at the midway, pH = pKa), and then rapidly rises towards the equivalence point (pH > 7).
  • A weak base–strong acid titration curve has the inverse form of a weak acid–strong acid titration curve. The pH decreases to the equivalence point (pH 7) in the strong base scenario.
  • A weak acid with a distinct colored conjugate base form is an acid-base (pH) indicator and changes hue during the course of roughly two pH units. Polyprotic acids contain two or more acidic protons, each with its own ion.
  • Polyprotic acids are composed of two or more acidic protons, each with its own Ka value. Because the Kas varies by many orders of magnitude, each proton is titrated individually.
  • Amino acids exist in charged forms that are determined by the pH of the solution. The total charge of a protein, which can impact its function.