# VI: Acid–Base Equilibria

Acids and bases have been defined differently by three sets of theories. One is the Arrhenius definition, which revolves around the idea that acids are substances that ionize (break off) in an aqueous solution to produce hydrogen (H+) ions while bases produce hydroxide (OH-) ions in solution. On the other hand, the Bronsted-Lowry definition defines acids as substances that donate protons (H+) whereas bases are substances that accept protons. Also, the Lewis theory of acids and bases states that acids are electron pair acceptors while bases are electron pair donors. Acids and bases can be defined by their physical and chemical observations.

• 6.1: Classifications of Acids and Bases
In chemistry, acids and bases have been defined differently by three sets of theories: One is the Arrhenius definition defined above, which revolves around the idea that acids are substances that ionize (break off) in an aqueous solution to produce hydrogen (H+) ions while bases produce hydroxide (OH-) ions in solution. The other two definitions are discussed in detail alter in the chapter and include the Brønsted-Lowry definition and the Lewis theory.
• 6.2: The Brønsted-Lowry Scheme
We now turn our attention to acid–base reactions to see how the concepts of chemical equilibrium and equilibrium constants can deepen our understanding of this kind of chemical behavior. We begin with a qualitative description of acid–base equilibria in terms of the Brønsted–Lowry model and then proceed to a quantitative description the following sections.
• 6.3: Acid and Base Strength
The strengths of Brønsted-Lowry acids and bases in aqueous solutions can be determined by their acid or base ionization constants. Stronger acids form weaker conjugate bases, and weaker acids form stronger conjugate bases. Thus strong acids are completely ionized in aqueous solution because their conjugate bases are weaker bases than water. Weak acids are only partially ionized because their conjugate bases are strong enough to compete successfully with water for possession of protons.
• 6.4: Equilibria Involving Weak Acids and Bases
The strengths of Brønsted-Lowry acids and bases in aqueous solutions can be determined by their acid or base ionization constants. Stronger acids form weaker conjugate bases, and weaker acids form stronger conjugate bases. Thus strong acids are completely ionized in aqueous solution because their conjugate bases are weaker bases than water. Weak acids are only partially ionized because their conjugate bases are strong enough to compete successfully with water for possession of protons.
• 6.5: Buffer Solutions
Buffers are solutions that resist a change in pH after adding an acid or a base. Buffers contain a weak acid ($$HA$$) and its conjugate weak base ($$A^−$$). Adding a strong electrolyte that contains one ion in common with a reaction system that is at equilibrium shifts the equilibrium in such a way as to reduce the concentration of the common ion. Buffers are characterized by their pH range and buffer capacity.
• 6.6: Acid-Base Titration Curves
The shape of a titration curve, a plot of pH versus the amount of acid or base added, provides important information about what is occurring in solution during a titration. The shapes of titration curves for weak acids and bases depend dramatically on the identity of the compound. The equivalence point of an acid–base titration is the point at which exactly enough acid or base has been added to react completely with the other component.
• 6.7: Polyprotic Acids
An acid that contains more than one ionizable proton is a polyprotic acid. The protons of these acids ionize in steps. The differences in the acid ionization constants for the successive ionizations of the protons in a polyprotic acid usually vary by roughly five orders of magnitude. It is often (but not always) appropriate to break down the calculations of the concentrations of the ions in solution into a series of steps.
• 6.8: A Deeper Look - Exact Treatment of Acid-Base Equilibria
The methods for dealing with acid-base equilibria that we developed in the earlier units of this series are widely used in ordinary practice. Although many of these involve approximations of various kinds, the results are usually good enough for most purposes. In this unit, we look at exact, or "comprehensive" treatment of some of the more common kinds of acid-base equilibria problems.
• 6.9: Organic Acids and Bases - Structure and Reactivity
Inductive effects and charge delocalization significantly influence the acidity or basicity of a compound. The acid–base strength of a molecule depends strongly on its structure. The weaker the A–H or B–H+ bond, the more likely it is to dissociate to form an H+ ion. In addition, any factor that stabilizes the lone pair on the conjugate base favors the dissociation of H+, making the conjugate acid a stronger acid.
• 6.10: Lewis Bases as Nucleophilic Reactants in Organic Reactions
Lewis proposed that the electron pair is the dominant actor in acid-base chemistry. An Lewis acid is a substance that accepts a pair of electrons, and in doing so, forms a covalent bond with the entity that supplies the electrons. A Lewis base is a substance that donates an unshared pair of electrons to a recipient species with which the electrons can be shared. Lewis acis/base theory is a powerful tool for describing many chemical reactions used in organic and inorganic chemistry.