- A characteristic sour taste
- Changes the color of litmus from blue to red
- Reacts with certain metals (Figure 1) to produce gaseous H2
- Reacts with bases to form a salt and water
Aqueous acids have a pH under 7, with acidity increasing the lower the pH. Common examples of acids include acetic acid (in vinegar), sulfuric acid (used in car batteries), and tartaric acid (used in baking). As these three examples show, acids can be solutions, liquids, or solids.
There are three common definitions for acids:
- Arrhenius: acids are substances that increase the concentration of hydronium ions (H3O+) in solution.
- Brønsted-Lowry: an expansion of the Arrhenius definition, an acid is a substance that can act as a proton donor.
- Lewis: Lewis acids are electron-pair acceptors. Examples of Lewis acids can be seen here (Figure 2).
Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, and the first two definitions are most relevant. The reason why pHs of acids are less than 7 is that the concentration of hydronium ions is greater than 10−7 moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acids thus have pHs of less than 7. By the Brønsted-Lowry definition, any compound that can easily be deprotonated can also be considered an acid. Examples include alcohols and amines that contain O-H or N-H fragments.
Hydronium ions are acids according to all three definitions. Interestingly, although alcohols and amines can be Brønsted-Lowry acids as mentioned above, they can also function as Lewis bases due to the lone pairs of electrons on their oxygen and nitrogen atoms.
The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely dissociates in water; in other words, one mole of the strong acid HA dissolves in water, yielding one mole of H+ and one mole of the conjugate base, A−, and none of the protonated acid HA. In contrast, a weak acid only partially dissociates, and at equilibrium both the acid and the conjugate base are in solution. Examples of strong acids are hydrochloric acid (HCl), hydroiodic acid (HI), hydrobromic acid (HBr), perchloric acid (HClO4), nitric acid (HNO3), and sulfuric acid (H2SO4). Each of these essentially ionizes 100% in water. Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths are also often discussed in terms of the stability of the conjugate base. Stronger acids have a larger Ka and a more negative pKa than weaker acids.
A base is a substance that can accept hydrogen ions (protons) or, more generally, donate a pair of valence electrons. A soluble base is referred to as an alkali if it contains and releases hydroxide ions (OH−) quantitatively. Again, we have three common definitions:
- Arrhenius: a base is a hydroxide anion, which is strictly applicable only to alkali.
- Brønsted-Lowry: a base is a proton (hydrogen ion) acceptor.
- Lewis: a base is an electron pair donor, examples can be seen here (Figure 2).
In water, bases give solutions with a hydrogen ion activity lower than that of pure water, i.e., a pH higher than 7.0 at standard conditions. A strong base is a basic chemical compound that is able to deprotonate very weak acids in an acid-base reaction. Common examples of strong bases are the hydroxides of alkali metals and alkaline earth metals like NaOH and Ca(OH)2. Very strong bases are even able to deprotonate very weakly acidic C–H groups in the absence of water. Acids with a pKa of more than about 13 are considered very weak, and their conjugate bases are strong bases.
Acids and bases react with one another to yield two products: water and an ionic compound known as a salt. This reaction is called neutralization. Bases and acids are seen as opposites because acids increase the hydronium ion (H3O+) concentration in water, whereas bases reduce this concentration. Bases and acids are typically found in aqueous solution forms. Aqueous solutions of bases react with aqueous solutions of acids to produce water and salts in aqueous solutions in which the salts separate into their component ions. If the aqueous solution is a saturated solution with respect to a given salt solute, any additional such salt present in the solution will result in formation of a precipitate of the salt.