Definitions of acids are abundant, and no single definition is used universally. Acids are typically defined either by the use of a relative scale, or by their chemical behavior.
The most familiar way for most people to recognize an acid is to see how it affects the pH of water. The pH scale is numerical, with values less than seven corresponding to an acidic pH. If a chemical is dissolved in water (which has a neutral pH of seven), and the pH is less than seven, the chemical that was dissolved is called an acid. The pH may be measured using a pH meter (which gives a numerical value) or by using a chemical indicator. Chemical indicators will change color at certain pH values. Litmus paper is a common example, and turns red when exposed to an acidic solution. Phenolphthalein is another popular indicator, remaining colorless in acidic solutions.
The use of the pH scale to define acids has a limited usefulness. It applies mostly to solutions of chemicals in water. While this encompasses most of daily experiences, it does not describe fully the nature of the chemical world, which includes solids, gases and solvents other than water.
To more completely define acids, chemists examine their chemical behavior. Several definitions have emerged over time, with three being well-known, and named after the chemist who proposed the definition.
An Arrhenius acid is a chemical that has a hydrogen atom that can dissociate (separate) from the chemical when dissolved in water, forming a hydrogen ion. This is an easily grasped concept, and applies to well-known acids like hydrochloric acid (HCl) and nitric acid (HNO3). This definition agrees well with the pH scale, but is limited only to molecules that have hydrogen that can be “lost”.
A Bronstead-Lowry acid is a chemical that can act as a “proton donor”. In chemistry parlance, “proton” in this case refers to the nucleus of a hydrogen atom. (The electron has been stripped away to form the hydrogen ion – a single, positively charged proton.) For a molecule or ion to be a proton donor (an acid), it must have an electronegative atom attached to a hydrogen. When the hydrogen leaves, the electrons remain behind on the electronegative atom. This definition is very similar to the Arrhenius definition. It also shares the same limitation. It relies on proton transfer to define acids. Outside of aqueous (water-based) environments, acid-base chemistry does not always revolve around proton exchange.
Lewis acids are known as electron acceptors. This typically refers to any electron-poor species, which includes (but is certainly not limited to) the hydrogen ion (the “proton” referred to in prior definitions), metal ions, and chemicals like boron tri-hydride which does not have enough electrons to provide boron with a full “octet”. Notice that this definition is significantly different from the previous two, even when dealing with hydrogen ions. Arrhenius, Bronstead and Lowry all regard the molecule that gives up the hydrogen ion as the acid. Lewis defines the hydrogen ion itself as the acid. This distinction is important conceptually, as it places the focus of the definition on the final acid-base reaction, and not on the ability of a molecule to dissociate. The Lewis definition is the broadest, most universally applicable definition. Its very breadth is perhaps its only limitation, as it is often more convenient to refer to a smaller set of chemicals when dealing with acids.
As with many things, context defines which definition will be used. Gardeners and swimmers will likely be satisfied with a pH reading. Young science students are typically introduced to acid-base chemistry using pH, indicators, and the Arrhenius definition. A full understanding of aqueous chemistry typically requires the Bronstead-Lowry definition, while the Lewis definition is necessary for full mastery of the subject.