This is an older version of this module. A newer version of the Acids and Bases I module is available
For thousands of years people have known that vinegar, lemon juice, and many other foods taste sour. However, it was not until a few hundred years ago that it was discovered why these things taste sour – because they are all acids. The term acid, in fact, comes from the Latin term acere, which means "sour". While there are many slightly different definitions of acids and bases, in this lesson we will introduce the fundamentals of acid/base chemistry.
Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases.
Bases feel slippery, change litmus blue, and become less basic when mixed with acids.
In the late 1800s, the Swedish scientist Svante Arrhenius proposed that water can dissolve many compounds by separating them into their individual ions. Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows:
Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH):
The Arrhenius definition of acids and bases explains a number of things. Arrhenius's theory explains why all acids have similar properties to each other (and, conversely, why all bases are similar): because all acids release H+ into solution (and all bases release OH-). The Arrhenius definition also explains Boyle's observation that acids and bases counteract each other. This idea, that a base can make an acid weaker, and vice versa, is called neutralization.
As you can see from the equations, acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to make the molecule H2O, or plain water:
Though Arrhenius helped explain the fundamentals of acid/base chemistry, unfortunately his theories have limits. For example, the Arrhenius definition does not explain why some substances, such as common baking soda (NaHCO3), can act like a base even though they do not contain hydroxide ions.
In 1923, the Danish scientist Johannes Brønsted and the Englishman Thomas Lowry published independent yet similar papers that refined Arrhenius' theory. In Brønsted's words, "... acids and bases are substances that are capable of splitting off or taking up hydrogen ions, respectively." The Brønsted-Lowry definition broadened the Arrhenius concept of acids and bases.
The Brønsted-Lowry definition of acids is very similar to the Arrhenius definition: Any substance that can donate a hydrogen ion is an acid. (Under the Brønsted definition, acids are often referred to as proton donors because an H+ ion, hydrogen minus its electron, is simply a proton).
The Brønsted definition of bases is, however, quite different from the Arrhenius definition. The Brønsted base is defined as any substance that can accept a hydrogen ion. In essence, a base is the opposite of an acid. NaOH and KOH, as we saw above, would still be considered bases because they can accept an H+ from an acid to form water. However, the Brønsted-Lowry definition also explains why substances that do not contain OH- can act like bases. Baking soda (NaHCO3), for example, acts like a base by accepting a hydrogen ion from an acid as illustrated below:
Under the Brønsted-Lowry definition, both acids and bases are related to the concentration of hydrogen ions present. Acids increase the concentration of hydrogen ions, while bases decrease the concentration of hydrogen ions (by accepting them). The acidity or basicity of something, therefore, can be measured by its hydrogen ion concentration.
In 1909, the Danish biochemist Sören Sörensen invented the pH scale for measuring acidity. The pH scale is described by the formula:
|pH = -log [H+]|
|Note: Concentration is commonly abbreviated by using square brackets, thus [H+] = hydrogen ion concentration. When measuring pH, [H+] is in units of moles of H+ per liter of solution.|
For example, a solution with [H+] = 1 x 10-7 moles/liter has a pH equal to 7 (a simpler way to think about pH is that it equals the exponent on the H+ concentration, ignoring the minus sign). The pH scale ranges from 0 to 14. Substances with a pH between 0 and less than 7 are acids (pH and [H+] are inversely related - lower pH means higher [H+]). Substances with a pH greater than 7 and up to 14 are bases (higher pH means lower [H+]). Right in the middle, at pH = 7, are neutral substances, for example, pure water. The relationship between [H+] and pH is shown in the table below alongside some common examples of acids and bases in everyday life.
|Acids||1 X 100||0||HCl|
|1 x 10-1||1||Stomach acid|
|1 x 10-2||2||Lemon juice|
|1 x 10-3||3||Vinegar|
|1 x 10-4||4||Soda|
|1 x 10-5||5||Rainwater|
|1 x 10-6||6||Milk|
|Neutral||1 x 10-7||7||Pure water|
|Bases||1 x 10-8||8||Egg whites|
|1 x 10-9||9||Baking soda|
|1 x 10-10||10||Tums® antacid|
|1 x 10-11||11||Ammonia|
|1 x 10-12||12||Mineral lime - Ca(OH)2|
|1 x 10-13||13||Drano®|
|1 x 10-14||14||NaOH|
Since acids and bases were first labeled and described in the 17th century, their definition has been refined over the centuries to reflect an increased understanding of their chemical properties. This module introduces the fundamentals of acid/base chemistry, including neutralization reactions. The relationship between hydrogen ion concentration [H+] and pH is shown alongside everyday examples of acids and bases.