Concept Summary

Definitions

• Arrhenius acids dissociate to produce an excess of hydrogen ions in solution. Arrhenius bases dissociate to produce an excess of hydroxide ions in solution.
• Brønsted–Lowry acids are species that can donate hydrogen ions. Brønsted–Lowry bases are species that can accept hydrogen ions.
• Lewis acids are electron-pair acceptors. Lewis bases are electron-pair donors.
• All Arrhenius acids and bases are Brønsted–Lowry acids and bases, and all Brønsted–Lowry acids and bases are Lewis acids and bases; however, the converse of these statements is not necessarily true (that is, not all Lewis acids and bases are Brønsted–Lowry acids and bases, and not all Brønsted–Lowry acids and bases are Arrhenius acids and bases).

• Amphoteric species are those that can behave as an acid or base. Amphiprotic species are amphoteric species that specifically can behave as a Brønsted–Lowry acid or Brønsted–Lowry base.

  • Water is a classic example of an amphoteric, amphiprotic species—it can accept a hydrogen ion to become a hydronium ion, or it can donate a hydrogen ion to become a hydroxide ion.
  • Conjugate species of polyvalent acids and bases can also behave as amphoteric and amphiprotic species.

Properties

• The water dissociation constant, K_w, is 10⁻¹⁴ at 298 K. Like other equilibrium constants, K_w is only affected by changes in temperature.
• pH and pOH can be calculated given the concentrations of H₃O⁺ and OH⁻ ions, respectively. In aqueous solutions, pH + pOH = 14 at 298 K.
• Strong acids and bases completely dissociate in solution.
• Weak acids and bases do not completely dissociate in solution and have corresponding dissociation constants (K_a and K_b, respectively).

• In the Brønsted–Lowry definition, acids have conjugate bases that are formed when the acid is deprotonated. Bases have conjugate acids that are formed when the base is protonated.

  • Strong acids and bases have very weak (inert) conjugates.
  • Weak acids and bases have weak conjugates.
• Neutralization reactions form salts and (sometimes) water.

Polyvalence and Normality

• An equivalent is defined as one mole of the species of interest.
• In acid–base chemistry, normality is the concentration of acid or base equivalents in solution.
• Polyvalent acids and bases are those that can donate or accept multiple electrons. The normality of a solution containing a polyvalent species is the molarity of the acid or base times the number of protons it can donate or accept.

Titration and Buffers

• Titrations are used to determine the concentration of a known reactant in a solution.

  • The titrant has a known concentration and is added slowly to the titrand to reach the equivalence point.
  • The titrand has an unknown concentration but a known volume.
• The half-equivalence point is the midpoint of the buffering region, in which half of the titrant has been protonated (or deprotonated); thus, [HA] = [A⁻] and a buffer is formed.

• The equivalence point is indicated by the steepest slope in a titration curve; it is reached when the number of acid equivalents in the original solution equals the number of base equivalents added, or vice-versa.

  • Strong acid and strong base titrations have equivalence points at pH = 7.
  • Weak acid and strong base titrations have equivalence points at pH > 7.
  • Weak base and strong acid titrations have equivalence points at pH < 7.
  • Weak acid and weak base titrations can have equivalence points above or below 7, depending on the relative strength of the acid and base.

• Indicators are weak acids or bases that display different colors in their protonated and deprotonated forms.

  • The indicator chosen for a titration should have a pK_a close to the pH of the expected equivalence point.
  • The endpoint of a titration is when the indicator reaches its final color.
• Multiple buffering regions and equivalence points are observed in polyvalent acid and base titrations.
• Buffer solutions consist of a mixture of a weak acid and its conjugate salt or a weak base and its conjugate salt; they resist large fluctuations in pH.
• Buffering capacity refers to the ability of a buffer to resist changes in pH; maximal buffering capacity is seen within 1 pH point of the pK_a of the acid in the buffer solution.
• The Henderson–Hasselbalch equation quantifies the relationship between pH and pK_a for weak acids and between pOH and pK_b for weak bases; when a solution is optimally buffered, pH = pK_a and pOH = pK_b.