What Are Acids and Bases?

There are several methods of defining acids and bases. While these definitions don't contradict each other, they do vary in how inclusive they are. The most common definitions of acids and bases are Arrhenius acids and bases, Brønsted-Lowry acids and bases, and Lewis acids and bases. Antoine Lavoisier, Humphry Davy, and Justus Liebig also made observations regarding acids and bases, but didn't formalize definitions.

Svante Arrhenius Acids and Bases

The Arrhenius theory of acids and bases dates back to 1884, building on his observation that salts, such as sodium chloride, dissociate into what he termed ions when placed into water.

• acids produce H⁺ ions in aqueous solutions
• bases produce OH^- ions in aqueous solutions
• water required, so only allows for aqueous solutions
• only protic acids are allowed; required to produce hydrogen ions
• only hydroxide bases are allowed

Johannes Nicolaus Brønsted - Thomas Martin Lowry Acids and Bases

The Brønsted or Brønsted-Lowry theory describes acid-base reactions as an acid releasing a proton and a base accepting a proton. While the acid definition is pretty much the same as that proposed by Arrhenius (a hydrogen ion is a proton), the definition of what constitutes a base is much broader.

• acids are proton donors
• bases are proton acceptors
• aqueous solutions are permissible
• bases besides hydroxides are permissible
• only protic acids are allowed

Gilbert Newton Lewis Acids and Bases

The Lewis theory of acids and bases is the least restrictive model. It doesn't deal with protons at all, but deals exclusively with electron pairs.

• acids are electron pair acceptors
• bases are electron pair donors
• least restrictive of the acid-base definitions

Properties of Acids and Bases

Robert Boyle described the qualities of acids and bases in 1661. These characteristics may be used to easily distinguish between the two sets up chemicals without performing complicated tests:

Acids

• taste sour (don't taste them!)—the word 'acid' comes from the Latin acere, which means 'sour'
• acids are corrosive
• acids change litmus (a blue vegetable dye) from blue to red
• their aqueous (water) solutions conduct electric current (are electrolytes)
• react with bases to form salts and water
• evolve hydrogen gas (H₂) upon reaction with an active metal (such as alkali metals, alkaline earth metals, zinc, aluminum)

Common Acids

• citric acid (from certain fruits and veggies, notably citrus fruits)
• ascorbic acid (vitamin C, as from certain fruits)
• vinegar (5% acetic acid)
• carbonic acid (for carbonation of soft drinks)
• lactic acid (in buttermilk)

Bases

• taste bitter (don't taste them!)
• feel slippery or soapy (don't arbitrarily touch them!)
• bases don't change the color of litmus; they can turn red (acidified) litmus back to blue
• their aqueous (water) solutions conduct an electric current (are electrolytes)
• react with acids to form salts and water

Common Bases

• detergents
• soap
• lye (NaOH)
• household ammonia (aqueous)

Strong and Weak Acids and Bases

The strength of acids and bases depends on their ability to dissociate or break into their ions in water. A strong acid or strong base completely dissociates (e.g., HCl or NaOH), while a weak acid or weak base only partially dissociates (e.g., acetic acid).

The acid dissociation constant and base dissociation constant indicates the relative strength of an acid or base. The acid dissociation constant K_a is the equilibrium constant of an acid-base dissociation:

HA + H₂O ⇆ A^- + H₃O⁺

where HA is the acid and A^- is the conjugate base.

K_a = [A^-][H₃O⁺] / [HA][H₂O]

This is used to calculate pK_a, the logarithmic constant:

pk_a = - log₁₀ K_a

The larger the pK_a value, the smaller the dissociation of the acid and the weaker the acid. Strong acids have a pK_a of less than -2.