The substances known as acids, bases, and salts are among the most important and best studied in chemistry. They are widespread in nature, in industry, and in the home.
Hydrochloric acid, present in small quantities in our stomachs, helps digest our food. Ascorbic acid, or vitamin C, is among our most important nutrients. Carbonic acid is familiar to us in the form of soda water. And the much stronger sulfuric acid is used to manufacture fertilizers, paints, synthetic fibers, and plastics.
Among the more familiar bases are lye, or caustic soda, and ammonia water. By far the best-known salt is sodium chloride, or common table salt. In addition to filling our salt shakers, sodium chloride is a basic ingredient of seawater as well as of our own blood. Other well-known salts include baking soda, or sodium bicarbonate, and Epsom salts.PROPERTIES
It is easy to describe acids in terms of what they do. Among their most familiar properties, they taste sour and corrode metals. Vinegar, for example, gets its tart taste from acetic acid.
Vinegar is a fairly dilute form of acetic acid. At full strength, this acid blisters the skin. Sulfuric acid is so corrosive that even a drop can cause deeply serious burns. In some very weak acids such as carbonic acid, the properties of sourness and corrosion are hardly noticeable to the nonchemist. Some weak acids are even used as medicines, including acetylsalicylic acid—better known as aspirin.
Bases are best known for bitter taste and the "soapy" feel they impart to water. In fact, many bases such as sodium hydroxide are essential ingredients in a number of familiar household and industrial soaps. The base ammonia is a household and industrial cleaner in its own right. Strong bases, like strong acids, can be quite corrosive and burn the skin.
When a base is mixed with an acid in the right proportions, the two substances react to form a salt and water. A salt is neutral—that is, neither an acid nor a base. Salts neither taste sour nor feel soapy in solutions, but have their own unique qualities.ACIDS
Just what is an acid? To go beyond describing an acid's properties to understanding its chemistry, we need to briefly review the basic structure of the atom and its binding forces.
The hydrogen atom, with only one proton and one electron, is the simplest of all the atoms. Suppose now that it loses its electron. All that is left is a single proton. No longer a neutral element, a hydrogen atom stripped of its electron has become an ion—an electrically charged particle. The hydrogen ion bears a single positive charge. We represent this with the symbol H+. Since the hydrogen ion consists of a single proton, chemists use the terms "hydrogen ion" and "proton" interchangeably.
Now that we know what a hydrogen ion is, we can understand one of the most widely accepted definitions of an acid: Acids are any substances that produce hydrogen ions, or protons—that is, a molecule that can lose one or more hydrogen ions from its structure. In fact, it is the hydrogen ion derived from acids that gives acidic foods their sour taste.
Let us see what happens when hydrochloric acid (hydrogen chloride, HCl) dissolves in water (H2O). Theoretically, the HCl could split into two ions—a hydrogen ion (H+) and a chlorine ion (Cl−).
In actual fact, the hydrogen ion (H+) never exists by itself in a solution of water. In such a situation, the ion immediately combines with a water molecule, H2O, as follows:
|H+ + H2O → H3O+|
This gives us the overall chemical equation for hydrochloric acid dissolved in water:
|HCl + H2O → H3O+ + Cl−|
H3O+ is called a hydronium ion—a hydrogen ion combined with a molecule of water. It is not a particularly stable molecule, however, and it easily splits back into a molecule of water and a hydrogen ion. Likewise, a hydronium ion will freely give up its "extra" hydrogen to any substance that shows an attraction for it—namely, a negative ion.
The end result of all this chemical give-and-take is a multitude of readily available hydrogen ions. So we see that HCl fits our definition of an acid.BASES
Like acids, bases ionize, or split into ions, when dissolved in water. But while acids split to produce hydrogen ions (H+), most bases produce hydroxyl ions (OH−) when dissolved in water. As you can see from its formula, the hydroxyl ion is made up of an oxygen and a hydrogen atom that together bear a single negative charge.
To illustrate, let us look at what happens when we dissolve the familiar base lye (sodium hydroxide, NaOH) in water. When NaOH ionizes, the chemical split takes place between the sodium and the oxygen atom. The resulting products are a positive sodium ion (Na+) and the negative hydroxyl ion (OH−). So we write the equation for the ionization of sodium hydroxide as follows:
|NaOH → Na+ + OH−|
Just as the hydrogen ions (H+) give acidic foods their sour taste, hydroxyl ions give basic solutions their soapy feel.
For many years, chemists defined a base as any substance that yields hydroxyl ions in water. But in modern times, chemists found that this definition was insufficient to encompass all bases. Today, the most widely accepted definition of a base can be stated as follows: a base is a chemical substance that can gain hydrogen ions, or protons, in a chemical reaction. So just as an acid is a proton producer, a base is a proton acceptor, a concept illustrated in the following reaction:
|H+ + OH− → H2O|
Not surprisingly, the negatively charged hydroxyl ion exhibits a strong attraction to a positively charged hydrogen ion. As you can see in the equation shown above, the resulting reaction forms water.
But as you will recall, in a water solution, the hydrogen ion is always carried on a water molecule in the form of a hydronium ion (H3O+). So the following balanced reaction is considerably more accurate:
|H3O+ + OH− → 2 H2O|
(two water molecules)
When such a reaction occurs, it is said that the acid and base have neutralized one another. When this "cancelling out" process occurs, a solution loses any acidic properties such as sour taste or any basic properties such as soapiness. The end product is simply water.SALT FORMATION
The last equation, concentrating on the hydronium ion and the hydroxyl ion, did not show all the ions that are present when an acid and a base react with one another in water. Suppose that we mix hydrochloric acid (HCl) with lye (NaOH), which, as we have seen, is a base. The hydrochloric acid solution has chlorine ions, Cl−, in addition to hydronium ions, H3O+. The sodium hydroxide solution has sodium ions, Na+, as well as hydroxyl ions, OH−. We can write the full equation as follows:
|H3O+ + Cl− + Na+ + OH− → 2H2O + Na+ + Cl−|
The right-hand side of the equation, giving the product of the reaction, shows water, sodium ions, and chlorine ions.
This is exactly what we get when we dissolve table salt (NaCl) in water, as follows:
|NaCl → Na+ + Cl−|
So you can see that when we react hydrochloric acid and sodium hydroxide in water, the product is simply a solution of table salt in water. If we evaporate the water, the salt will remain as a white solid.
Another example of a salt is potassium nitrate (KNO3), a colorless solid that, when mixed with charcoal and a little sulfur, makes gunpowder. The potassium nitrate molecule consists of one atom of potassium (K), one atom of nitrogen (N), and three atoms of oxygen (O3). This salt forms when nitric acid (HNO3) reacts with the base potassium hydroxide (KOH).
Specifically, nitric acid dissolves in water, forming the positive hydronium ion (H3O+) and the negative nitrate ion (NO3−) as follows:
|HNO3 + H2O → H3O+ + NO3−|
Potassium hydroxide, in water, forms hydroxyl and potassium ions (K+). So mixing nitric acid and potassium hydroxide in water produces the following balanced reaction:
|[H3O+ + NO3−] + [K+ + OH−] → K+ + NO3− + 2H2O|
When the solution is evaporated, we are left with a colorless, solid salt—potassium nitrate (KNO3).
In both of these examples, the hydrogen ions of an acid were, in effect, replaced by the metallic ions (sodium or potassium) of a base, to produce a salt. This is usually, but not always, the case in the formation of salts, which can be defined as a solid compound made up of a positive ion, or cation, and a negative ion, or anion. Some acids can lose more than one hydrogen ion, trading them in for one or more metallic ions.
For example, when sulfuric acid (H2SO4) trades in one hydrogen ion for a metallic ion, it produces an acid sulfate salt, such as NaHSO4 or KHSO4. If both hydrogens are replaced, a normal sulfate salt is produced, as is the situation in Na2SO4 or K2SO4.
In some cases, a single metallic ion carrying a double negative charge replaces two hydrogens. The magnesium ion (Mg++) reacts with sulfuric acid (H2SO4) to produce the salt magnesium sulfate (MgSO4).ACIDIC AND BASIC SOLUTIONS
Acids, bases, and salts all split into ions when dissolved in water. Water itself also ionizes, although only slightly. It forms hydrogen and hydroxyl ions, and the hydrogen ions combine with intact water molecules to form hydronium ions as follows:
Note that of the two arrows in the above equation, the smaller arrow points from left to right. This indicates that the reaction goes both ways, and that fewer hydronium and hydroxyl ions are formed than are molecules of water.
In fact, far less than 1 percent of an ordinary water solution exists in ion form. This is because the oppositely charged hydroxyl and hydronium ions are so very quick to react with one another to form water.pH. In ordinary water, there are as many hydronium (H3O+) ions as hydroxyl (OH−) ions, and the amount of each is exceedingly small. Chemists usually refer to concentrations of hydronium ions as concentrations of hydrogen ions (an H+ carried on an H2O). Accordingly, the concentration of hydrogen ions in pure water is 0.0000001 gram per liter. This number can be written as 10−7.
Solutions with a comparatively high concentration of hydrogen ions (and a slight concentration of hydroxyl ions) are called acidic. Solutions in which the concentration of hydroxyl (OH−) ions is greater than hydrogen ions are called basic, or alkaline (alkali means base). Since there are as many hydrogen as hydroxyl ions in ordinary water, it is neither acidic nor basic, but neutral.
A solution with a hydrogen-ion concentration of 0.1 gram per liter would have 1 million times the hydrogen-ion concentration of pure water. We would call such a solution strongly acidic. If the hydrogen-ion concentration were 0.0001 gram per liter, it would be 1,000 times higher than the hydrogen-ion concentration in water. This solution would still be considered acidic, but not so much so as the first.
We can express the acidity of a solution, therefore, by giving its hydrogen-ion concentration. We could write this as 10−1 and 10−4, respectively, for the two solutions above. When chemists express acidity, they simplify this notation still further by dropping the 10 and simply giving the exponent without the minus sign. So the acidity of the above two solutions would be given as 1 and 4, respectively.
This is called the solution's pH (read "P-H"). So a solution with a hydrogen-ion concentration of 10−3 grams per liter would have a pH value of 3. A neutral solution, or pure water, has a pH of 7. A pH of 6 is only slightly acidic; a pH of 1, extremely acidic.
The pH system can be applied to basic solutions, too. A pH of 8 would mean a hydrogen-ion concentration of 10−8, or 0.00000001 gram per liter. This is only one-tenth of the concentration of the hydrogen ion in pure water, and would imply that the hydroxyl (OH−) concentration would be relatively greater. Such a solution would be slightly alkaline.
So we see, the higher the pH, the more basic, or alkaline, a solution is. The pH system is the most convenient way to express the degree of acidity or alkalinity of any solution, with a range of 1 to 14.
It is a matter of great importance in many fields of science and industry to be able to determine the acidity of a solution and express it in terms of pH. Chemists, biologists, bacteriologists, agricultural experts, and many others are interested in the pH of the particular solutions with which they work.
The physiologist, for instance, is concerned with the acidity of various body fluids. The juices of the stomach are strongly acidic, having a pH of from 1.0 to 3.0. The pH of the blood ranges from 7.3 to 7.5. This means that the blood is very slightly alkaline. If the pH of the blood were to fall to 7.0, exactly neutral, or rise to 7.7, slightly more alkaline than normal, the results might be fatal. Fortunately, the body has systems of buffers to combat excessive acidity or alkalinity. These are substances that are able to remove, by chemical action, any excess of either H+ or OH−.pH Indicators. The chemist can determine the pH of a solution in various ways. One of the simplest systems involves the use of substances called indicators. When an indicator comes into contact with an acidic or basic solution, the indicator undergoes a color change that is specific to different ranges on the pH scale.
Litmus paper is among the most familiar indicators. It changes color at a pH of nearly 7, the value for a neutral solution. Other indicators change at different pH values. By testing a given solution with a number of different indicators, one can determine its pH with enough accuracy for many purposes.Titration. For a more accurate determination of pH, chemists frequently rely on a process called titration to determine the total amount of acid or base in an unknown solution. First, a specific amount of the solution is measured into a beaker or flask, and a few drops of an appropriate indicator are added.
If the indicator shows the solution to be basic, the chemist then begins to slowly add a solution with a known acid concentration. This process can be done very precisely, drop by drop, with an instrument called a burette. Eventually, the added acid will bring the basic solution to neutral (pH 7), and the indicator will change color. Then the chemist has only to note the volume of acid solution used, and he or she can easily calculate the concentration of the unknown base. Similarly, the chemist could determine the concentration of an unknown acidic solution by slowly adding an alkaline solution of known concentration.APPLICATIONS
Uses of Acids and Bases. Sulfuric acid is used in greater quantities than any other chemical. Its sales offer an excellent index of business conditions. It enters into the manufacture of explosives, dyestuffs, and drugs, and is likewise used in oil and sugar refining and in the preparation of fertilizers. Hydrochloric acid serves in the manufacture of glue, gelatin, and dextrose, and also in the cleaning of metals. Doctors prescribe very dilute hydrochloric acid for certain cases of indigestion. Nitric acid is used in large quantities in the manufacture of fertilizers and of explosives, and also in the preparation of dyes and plastics. Acetic acid is a good solvent for certain organic substances. It plays an important part in the preparation of cellulose acetate, which is used as a plastic and in the manufacture of acetate fiber. Hydrofluoric acid is used in making refrigerants and certain new plastics. It also has the useful property of being able to etch glass.
Many bases are likewise extremely useful. Sodium hydroxide is used in the rayon, film, soap, paper, and petroleum industries, and in many others. Hydrated lime (calcium hydroxide) serves to raise the pH of soils and to make insecticides and many other chemicals. Ammonium hydroxide, another important base, is employed as a raw material for making many important compounds, including the ones that are used in fertilizers.Uses of Salts. Table salt, sodium chloride, is among the most valuable of all salts. It is used to prepare various compounds of the metal sodium. It is also the principal source of chlorine, which is used for bleaching, for purifying water, and in many plastics. Sodium chloride is also an important preservative of meat, fish, hides, skins, and various other products.
The salt called sodium bicarbonate (NaHCO3) is an essential ingredient in baking powders. Epsom salts (magnesium sulfate, MgSO4 · 7H2O) and calomel (mercurous chloride, Hg2C12) have medicinal uses. Gypsum (calcium sulfate, CaSO4 · 2H2O) yields plaster of Paris and is also used in the manufacture of gypsum wallboard. Hypo (technically called sodium thiosulfate, Na2S2O3 · 5H2O) plays an important part in photography as a fixing agent.
A list of potassium salts, alone, provides a good idea of the wide-ranging uses of salts. Potassium carbonate (K2CO3) is used in making glass, dyes, and some soaps. Potassium bromide (KBr) is used in making photographic paper, plates, and developers. Potassium nitrate (KNO3) is used in the production of gunpowder, while potassium chloride (KCl) and sulfate (K2SO4) are used in making fertilizers. Pure monopotassium tartrate (KH5C4O6), called cream of tartar, is used in the manufacture of baking powder.
These are but a few examples of the ways in which acids, bases, and salts play a vital role in our daily lives.