The number of hydrogen ions present in a solution is a measure of the acidity of the solution. All acids do not ionize completely when dissolved in water, i.e., all the molecules of acid do not ionize and exist in the solution as electrically-charged particles. The hydrogen ion concentration is a measure, therefore, of the amount of dissociated acid rather than of the an amount of acid present. Strong acids dissociate more freely than weak acids; hydrochloric acid, for example, dissociates freely into H+ and Cl− whereas carbonic acid, a weak acid, dissociates much less freely into H+ and CO3–. The number of free hydrogen ions is a measure of its acidity rather than an indication of the type of molecule from which the hydrogen ions originated.
The alkalinity of a solution is dependent upon the number of hydroxyl ions present. Water is a neutral solution because each molecule contains one H+ and one OH−. For each molecule of water dissociated, there is one H+ and one OH−, each one neutralizing the other.
A neutral solution, such as water, where the number of hydrogen ions is balanced by the same number of hydroxyl ions, has a pH of 7.0. The range of the pH scale is from 0 to 14.
The pH scale runs from 1 to 14; neutrality being at pH 7.0.
Solutions having a pH lesser than 7 are acidic, i.e., the concentration of H+ is greater than that of OH−.
Conversely, solutions having a pH more than 7 are basic or alkaline, i.e., denote an excess of OH− over H+
Brönsted−Lowry concept of Acids and Bases
In aqueous systems, the addition or removal of hydrogen ions is best understood in terms of the Brönsted−Lowry concept of acids and bases, propounded in 1923.
A Brönsted−Lowry acid is defined as a substance that can donate a proton (H+); conversely, a Brönsted−Lowry base is a substance that can accept a proton.
A proton donor (i.e., an acid) and its corresponding proton acceptor (i.e., a base) make up a conjugate (coniungereL = to join together) acid−base pair. This broad definition of acids and bases includes many substances that are not usually
considered acidic or basic.
Strong and Weak Acids
There are two general classes of acids — strong and weak.
A strong acid is defined as a substance that has a greater tendency to lose its proton and therefore completely dissociates (or ionizes) in water, such as HCl and H2SO4.
A weak acid, on the other hand, is a molecule that has a lesser tendency to lose its proton (or, in other words, displays a high affinity for its proton) and, therefore, does not readily dissociate in water, such as CH3COOH.
A buffer solution is one that resists a change in pH on the addition of acid (H+) or base (OH−), more effectively than an equal volume of water. Most commonly, the buffer solution consists of a mixture of a weak Brönsted acid and its conjugate base; for example, mixtures of acetic acid and sodium acetate or of ammonium hydroxide and ammonium chloride are buffer solutions. A buffer system consists of a weak acid (the proton donor) and its conjugate base (the proton acceptor).
Cells and organisms maintain a specific and constant cytosolic pH, keeping biomolecules in their optimal ionic state, usually near pH 7.
In multicelled organisms, the pH of the extracellular fluids (blood, for example) is also tightly regulated. Constancy of pH is achieved primarily by biological buffers : mixtures of weak acids and their conjugate bases. A certain amount of many of these is usually present in the body and cellular fluids, and so the maintenance of a constant pH depends on a complex system.