The term acid was first used in the seventeenth century; it comes from the Latin root ac-, meaning “sharp”, as in acetum, vinegar. Some early writers suggested that acidic molecules might have sharp corners or spine-like projections that irritate the tongue or skin.
Acids have long been recognized as a distinctive class of compounds whose aqueous solutions exhibit the following properties:
- A characteristic sour taste (think of lemon juice!);
- ability to change the color of litmus* from blue to red;
- react with certain metals to produce gaseous H2;
- react with bases to form a salt and water.
*Litmus is a natural dye found in certain lichens. The name is of Scandinavian origin, e.g. lit (color) + mosi (moss) in Icelandic. "Litmus test" has acquired a meaning that transcends both Chemistry and science to denote any kind of test giving a yes/no answer.
How oxygen got mis-named
The first chemistry-based definition of an acid turned out to be wrong: in 1787, Antoine Lavoisier, as part of his masterful classification of substances, identified the known acids as a separate group of the “complex substances” (compounds). Their special nature, he postulated, derived from the presence of some common element that embodies the “acidity” principle, which he named oxygen, derived from the Greek for “acid former”.
Lavoisier had recently assigned this name to the new gaseous element that Joseph Priestly had discovered a few years earlier as the essential substance that supports combustion. Many combustion products (oxides) do give acidic solutions, and oxygen is in fact present in most acids, so Lavoisier’s mistake is understandable. In 1811 Humphrey Davy showed that muriatic (hydrochloric) acid (which Lavoisier had regarded as an element) does not contain oxygen, but this merely convinced some that chlorine was not an element but an oxygen-containing compound. Although a dozen oxygen-free acids had been discovered by 1830, it was not until about 1840 that the hydrogen theory of acids became generally accepted. By this time, the misnomer oxygen was too well established a name to be changed. The root oxy comes from the Greek word οξνς, which means "sour".
The key to understanding acids (as well as bases and salts) had to await Michael Faraday’s mid-nineteenth century discovery that solutions of salts (known as electrolytes) conduct electricity. This implies the existence of charged particles that can migrate under the influence of an electric field. Faraday named these particles ions (“wanderers”). Later studies on electrolytic solutions suggested that the properties we associate with acids are due to the presence of an excess of hydrogen ions in the solution. By 1890 the Swedish chemist Svante Arrhenius (1859-1927) was able to formulate the first useful theory of acids:
"an acidic substance is one whose molecular unit contains at least one hydrogen atom that can dissociate, or ionize, when dissolved in water, producing a hydrated hydrogen ion and an anion."
|hydrochloric acid||HCl → H+(aq) + Cl–(aq)|
|sulfuric acid||H2SO4→ H+(aq) + HSO4–(aq)|
|hydrogen sulfite ion||HSO3–(aq) → H+(aq) + SO32–(aq)|
|acetic acid||H3CCOOH → H+(aq) + H3CCOO–(aq)|
Strictly speaking, an “Arrhenius acid” must contain hydrogen. However, there are substances that do not themselves contain hydrogen, but still yield hydrogen ions when dissolved in water; the hydrogen ions come from the water itself, by reaction with the substance. A more useful operational definition of an acid is therefore the following:
There are three important points to understand about hydrogen in acids:
- Although all Arrhenius acids contain hydrogen, not all hydrogen atoms in a substance are capable of dissociating; thus the –CH3 hydrogens of acetic acid are “non-acidic”. An important part of knowing chemistry is being able to predict which hydrogen atoms in a substance will be able to dissociate into hydrogen ions; this topic is covered in a later lesson of this set.
- Those hydrogens that do dissociate can do so to different degrees. The strong acids such as HCl and HNO3 are effectively 100% dissociated in solution. Most organic acids, such as acetic acid, are weak; only a small fraction of the acid is dissociated in most solutions. HF and HCN are examples of weak inorganic acids.
- Acids that possess more than one dissociable hydrogen atom are known as polyprotic acids; H2SO4 and H3PO4 are well-known examples. Intermediate forms such as HPO42–, being capable of both accepting and losing protons, are called ampholytes. In the examples below, ampholytes are shown in green:
It turns out that hydrogen ions cannot exist as actual H+ entities in water, but don't panic! Chemists still find it convenient to pretend they are present, and to show them in equations that involve acid-base reactions. More on this in a later part of this lesson group.
The name base has long been associated with a class of compounds whose aqueous solutions are characterized by:
- a bitter taste;
- a “soapy” feeling when applied to the skin;
- ability to restore the original blue color of litmus that has been turned red by acids;
- ability to react with acids to form salts.
- react with certain metals to produce gaseous H2;
Just as an acid is a substance that liberates hydrogen ions into solution, a base yields hydroxide ions when dissolved in water:
NaOH(s) → Na+(aq) + OH–(aq)(3-1)
The abbreviations in parentheses denote the phydical state of the substance. Thus (s) refers to solid NaOH, while (aq) indicates that the two ions are dissolved in water (Latin aqua) and weakly linked to water molecules. We will make frequent use of this convention throughout this course.
Sodium hydroxide is an Arrhenius base because it contains hydroxide ions. NaOH is also a strong base; when dissolved in water, it is completely dissociatiated as indicated in (3-1) above. However, other substances which do not contain hydroxide ions can nevertheless produce them by reaction with water, and are therefore also classified as bases. Two classes of such substances are the alkali metal oxides and the hydrogen compounds of certain nonmetals:
Na2O(s) + H2O → [2 NaOH] → 2 Na+(aq) + 2 OH–(aq)(3-2)
Sodium oxide is a strong base because it reacts with water to produce sodium hydroxide which itself is a strong base.
NH3(g) + H2O → NH4+(aq) + OH–(aq)(3-3)
When ammonia gas dissolves in water only a small fraction of it reacts to produce ammonium ion and hydroxide ions; it is therefore classified as a weak base.
In order to accommodate bases such as ammonia that do not contain hydroxide ions of their own, a base is more generally defined as follows:
Acids and bases react with one another to yield two products: water, and an ionic compound known as a salt. This kind of reaction is called a neutralization reaction.
This "molecular" equation is convenient to write, but we need to re-cast it as a net ionic equation to reveal what is really going on here when the reaction takes place in water, as is almost always the case.
H+ + Cl–+ Na++ OH–→ Na+ + Cl– + H2O(4-2)
If we cancel out the ions (shown in gray) that appear on both sides (and therefore don't really participate in the reaction), we are left with the net equation
H+(aq) + OH–(aq) → H2O(4-3)
which is the fundamental process that occurs in all neutralization reactions.
Confirmation that this equation describes all neutralization reactions that take place in water is provided by experiments indicating that no matter what acid and base are combined, all liberate the same amount of heat (57.7 kJ) per mole of H+ neutralized.
The “salt” that is produced in a neutralization reaction consists simply of the anion and cation that were already present. The salt can be recovered as a solid by evaporating the water. Thus if the aqueous solution resulting from
(4-2) is evaporated to dryness, the net reaction reduces to
Na+(aq) + Cl–(aq) → NaCl(s)(4-4)
Complete neutralization of a polyprotic acid by a strong base such as NaOH consumes 1 mole of the base for each acidic hydrogen in the acid:
OH– + H2SO4 → HSO4– + H2O(4-5)
2 OH– + H2SO4 → SO4– + 2 H2O(4-6)