Unit 2
Elements of Chemical Change


The two most important classifications of compounds in inorganic chemistry are acids and bases. The following discussion forms the groundwork for understanding some of the most important chemical changes which you will encounter.

Classical Acid-Base Theory. Svante Arrhenius, in 1887, published the first satisfactory explanation of the acid-base phenomena that had been observed by chemists.

Acids. Arrhenius defined an acid as a compound that donates protons (H+) in solution. Examples would be any of the compounds you learned to name as acids earlier in this course.

NOTE: The HOH (H2O), which indicates that water is the solvent in these reactions. Both HCl and H2SO4 contribute protons in solution.

Bases. Arrhenius defined a base as any compound that donates hydroxyl (OH-) ions in solution. Again, you should be familiar with several examples from your nomenclature studies.

These classical definitions are based on the dissociation of the compounds into ions in solution. This implies that all acids and bases must contain exchangeable hydrogen and hydroxyl ions, respectively, in their formulas. This theory did explain the majority of the compounds known at the time, but there were some exceptions. Chemists knew, for example, that metal oxides (MgO, CaO, etc.) dissolved in water exhibited base-like properties. Also, ammonia (NH3) in solution exhibited the properties of a base. The attempts to explain these exceptions led to new definitions of acids and bases.

Modern Acid-Base Theory. In 1923, Bronsted and Lowry, two chemists in different countries, independently derived new definitions of acids and bases to explain the exceptions to the classical theory. The new theory they developed was named, appropriately, the Bronsted-Lowry theory. This theory differs from the classical theory in that the dissociation of water is considered as well as the dissociation of the compound.

  1. Dissociation of water. Even though we often think of water as merely being an inert solvent, it does dissociate into ions.

This is an equilibrium type reaction as indicated by the double arrow. Actually, very few ions exist at any time since they rapidly recombine to form molecular water. If we put numbers in this reaction, there are 500 million molecules of water for each hydrogen or hydroxyl ion.

  1. Bronsted-Lowry acid. By the Bronsted-Lowry theory, an acid is any compound (charged or uncharged) capable of donating a proton. This is essentially the same as the classical definition.
  2. Bronsted-Lowry base. The real value of the Bronsted-Lowry theory is in the definition of a base. A base is defined as a charged or uncharged substance capable of accepting a proton. Generally, the proton a base accepts comes from the dissociation of water.

Consider, for example, ammonia dissolved in water:

By accepting a proton from water, ammonia has effectively increased the concentration of hydroxyl ions in the solution. This would account for the properties like those of a classical base.

A second example would be magnesium oxide dissolved in water.

By accepting a proton from water, magnesium oxide has likewise increased the concentration of hydroxyl ions in the solution.

NOTE: The two theories explain all the properties of acids and bases that will be utilized in medicine. It deserves mention that there are other theories of acids and bases that explain more complex phenomena. If these are of interest to you, a college chemistry text should have a discussion of some of them.

Properties of Acids. We have defined all acids based on one common property, the ability to donate hydrogen ions in solution. Therefore, you should expect them all to exhibit a set of common properties, which they do. The properties we are concerned with are as follows:

  1. Acids change blue litmus paper to red. Litmus paper, which contains dyes sensitive to hydrogen ion concentration, turns red when there is a high concentration, blue when there is a low concentration.
  2. Acids have a sour taste. This property is familiar to you if you have ever tasted a lemon. Lemons contain citric acid, which gives them their sour taste.
  3. Acids react with metals to release hydrogen gas. For example:

Zn + 2H+ Zn++ + H2

You will notice that this reaction is an oxidation-reduction reaction. For practice, pick out the oxidizing and reducing agents.

  1. Acids react with carbonates and bicarbonates to form carbon dioxide. For example:

CaCO3 + 2HCl CaCl2 + H2O + CO2

  1. Acids react with bases to form salts and water (neutralization reaction). For example:
HCl + NaOH NaCl + H2O

Properties of Bases. In the same manner that all acids had certain properties in common, all bases have related properties. The ones that are important to the medical personnel are as follows:

Classification of Acids and Bases. Even though all acids possess certain properties in common, as do bases, not all possess them to the same degree. Some acids, for example, will completely neutralize sodium hydroxide with equal concentrations while others will only partially neutralize this base. As you might suspect, the differences in the strengths of acids results from differing abilities to donate hydrogen ions and the differences in bases from differing abilities to donate hydroxyl ions or accept hydrogen ions.

  1. Some acids and bases dissociate more readily than others when placed in solution. Those that dissociate at a rate greater than 50 percent are considered to be strong acids or bases. Weak acids and bases dissociate at a rate that is less than 50 percent. Examples:

  1. This means one mole (gram molecular weight) of HCl will produce more hydrogen ions in solution than will one mole of H2CO3 and will consequently exhibit acidic properties to a greater degree than will carbonic acid. A simpler way to say this is that HCl is a stronger acid than H2CO3.
  2. The same rationale holds for bases as well as acids. Therefore, we can divide or classify acids or bases into groups based on their dissociation--strong acids or bases (those that dissociate completely) and weak acids or bases (those that dissociate to a small degree).

Acids and Bases of Medicinal Importance. One may come in contact with a number of important acids and bases. You must be able to identify them as acids or bases and know their relative strengths. Table 2-1 shows these acids and bases. There is not an easy way to differentiate between strong and weak acids, but strong and weak bases can be differentiated based on valence. Strong bases have a positive valence of one; weak bases have a positive valence greater than one.

Safety and Antidotes. Acids and bases should be handled with care to avoid spilling on skin. They should not be taken internally unless intended for that purpose. If the skin is exposed to these compounds or is ingested, the following antidotes are recommended for first aid treatment.

  1. Acids.
  1. Bases.