conjugate acid-base pair

Examples of conjugate acid-base pair in the following topics:

• Buffer Range and Capacity

  • A buffer solution usually contains a weak acid and its conjugate base.
  • When H+ is added to a buffer, the weak acid's conjugate base will accept a proton (H+), thereby "absorbing" the H+ before the pH of the solution lowers significantly.
  • Similarly, when OH- is added, the weak acid will donate a proton (H+) to its conjugate base, thereby resisting any increase in pH before shifting to a new equilibrium point.
  • Each conjugate acid-base pair has a characteristic pH range where it works as an effective buffer.
  • In other words, the pH of the equimolar solution of acid (e.g., when the ratio of the concentration of acid and conjugate base is 1:1) is equal to the pKa.

• Absolute Concentrations of the Acid and Conjugate Base

  • A buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.
  • Therefore, it is very important to be able to identify acid and conjugate base pairs.
  • The conjugate acid is created by accepting (adding) a proton (H+) donated by the conjugate base.
  • A concentrated buffer can neutralize more added acid or base than a dilute buffer, because it contains more acid/conjugate base.
  • 8.1.3 Deduce the formula of the conjugate acid/base of any Brønsted-Lowry base/acid IB Chemistry SL - YouTube

• Relative Amounts of Acid and Base

  • A buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.
  • You can use one of these acid/conjugate base pairs:
  • Which pair should be used?
  • Extrapolating further from this, a buffer is most effective when the concentrations of acid and conjugate base (or base and conjugate acid) are approximately equal—in other words, when the log [base]/[acid] equals 0 and the pH equals the pKa.
  • The change is minimized if the concentrations of acid and conjugate base are equal.

• Acid-Base Reactions

  • Structurally related acid-base pairs, such as {H-A and A:(–)} or {B:(–) and B-H} are called conjugate pairs.
  • When the pH of an aqueous solution or mixture is equal to the pKa of an acidic component, the concentrations of the acid and base conjugate forms must be equal (the log of 1 is 0).
  • However, a more common procedure is to report the acidities of the conjugate acids of the bases (these conjugate acids are often "onium" cations).
  • According to the Lewis theory, an acid is an electron pair acceptor, and a base is an electron pair donor.
  • The resulting mixture of non-bonded Lewis acid/base pairs has been termed "frustrated", and exhibits unusual chemical behavior.

• The Brønsted-Lowry Definition of Acids and Bases

  • Keep in mind that acids and bases must always react in pairs.
  • Here, a conjugate base is the species that is left over after the Brønsted acid donates its proton.
  • The conjugate acid is the species that is formed when the Brønsted base accepts a proton from the Brønsted acid.
  • The products include the acetate ion, which is the conjugate base formed in the reaction, as well as hydronium ion, which is the conjugate acid formed.
  • The conjugate acid formed in the reaction is the ammonium ion, and the conjugate base formed is hydroxide.

• Acidity of Amines

  • We normally think of amines as bases, but it must be remembered that 1º and 2º-amines are also very weak acids (ammonia has a pKa = 34).
  • pKa is being used as a measure of the acidity of the amine itself rather than its conjugate acid, as in the previous section.
  • The first compound is a typical 2º-amine, and the three next to it are characterized by varying degrees of nitrogen electron pair delocalization.
  • The acids shown here may be converted to their conjugate bases by reaction with bases derived from weaker acids (stronger bases).
  • For complete conversion to the conjugate base, as shown, a reagent base roughly a million times stronger is required.

• Nature of Acids and Bases

  • Lewis acid: any substance that can accept a pair of electrons.
  • By definition, a strong acid is one that completely dissociates in water; in other words, one mole of the generic strong acid, HA, will yield one mole of H+, one mole of the conjugate base, A−, with none of the unprotonated acid HA remaining in solution.
  • At equilibrium, both the acid and the conjugate base will be present, along with a significant amount of the undissociated species, HA.
  • Acid strengths are also often discussed in terms of the stability of the conjugate base.
  • Acids + Bases Made Easy!

• Basicity of Amines

  • A review of basic acid-base concepts should be helpful to the following discussion.
  • Although 4-dimethylaminopyridine (DMAP) might appear to be a base similar in strength to pyridine or N,N-dimethylaniline, it is actually more than ten thousand times stronger, thanks to charge delocalization in its conjugate acid.
  • Here, as shown below, resonance stabilization of the base is small, due to charge separation, while the conjugate acid is stabilized strongly by charge delocalization.
  • The relationship of amine basicity to the acidity of the corresponding conjugate acids may be summarized in a fashion analogous to that noted earlier for acids.
  • Strong bases have weak conjugate acids, and weak bases have strong conjugate acids.

• Brønsted Acids and Bases

  • The conjugate base is the ion or molecule that remains after the acid has donated its proton, and the conjugate acid is the species created after the base accepts the proton.
  • Water is amphoteric, which means it can act as either an acid or a base.
  • The acetate ion CH3CO2- is the conjugate base of acetic acid, and the hydronium ion H3O+ is the conjugate acid of the base, water:
  • The hydroxide ion is the conjugate base of water, which acts as an acid, and the ammonium ion is the conjugate acid of the base, ammonia.
  • Identify the Brønsted acid, Brønsted base, conjugate acid, and conjugate base in an acid-base reaction.

• Strong Bases

  • As discussed in the previous concepts on bases, a base is a substance that can: donate hydroxide ions in solution (Arrhenius definition); accept H+ ions (protons) (Bronsted-Lowry definition); or donate a pair of valence electrons (Lewis definition).
  • Strong bases are capable of deprotonating weak acids; very strong bases can deprotonate very weakly acidic C–H groups in the absence of water.
  • When writing out the dissociation equation of a strong base, assume that the reverse reaction does not occur, because the conjugate acid of a strong base is very weak.
  • Usually, these bases are created by adding pure alkali metals in their neutral state, such as sodium, to the conjugate acid.
  • Unlike weak bases, which exist in equilibrium with their conjugate acids, the strong base reacts completely with water, and none of the original anion remains after the base is added to solution.