Acids and bases

• Type of physical science: Chemistry
• Field of study: Chemical compounds

Acids are substances that release an H⁺ ion into a solution; bases are substances that take up an H⁺ ion from a solution. The relative amounts of acid and base determine the degree of acidity, or pH, of a solution, which is of vital importance to the progress of chemical reactions taking place in living organisms, industry, and the environment.

Overview

Acids and bases are highly reactive substances that enter into chemical interactions as reactants and also as molecules setting up conditions allowing other substances to react. For most reactions, the significant feature of acids and bases is their effect on the concentration of H⁺ ions in a solution or other locales.

An ion is any charged atom or group of atoms; an H⁺ ion is the most simple form an ion can take: a "naked" atomic nucleus consisting of a single proton carrying a positive charge, with no accompanying neutron and no surrounding cloud of electrons. One of the most inclusive definitions of acids and bases reflects the ability of these substances to alter the concentration of H⁺ ions in a medium. In these terms, an acid is defined as any substance that can release an H⁺ ion, thereby increasing the concentration of protons in the surrounding solution. A base is any substance with the opposite activity, the ability to take up an H⁺ ion, thereby reducing the concentration of protons in the solution.

The number of H⁺ ions per unit volume (or, in other words, the concentration of H⁺ ions) is a measure of the acidity of a solution. In a water solution, the concentration of H⁺ ions has a fixed relationship with another ion, the OH^- ion, which consists of an oxygen and hydrogen nucleus linked together by a cloud of surrounding electrons. The cloud contains one more electrons than the number of positive charges in the hydrogen and oxygen nuclei, giving the OH^- ion a single negative charge. In a water solution, the product of the concentrations of the H⁺ and OH^- ions is always equal to 0.00000000000001, or 1 x 10^-14 in scientific notation. These concentrations are given in moles per liter (a mole is the molecular weight of a substance in grams). In order for the H⁺ concentration to rise or fall, the concentration of OH^- ions must fall or rise proportionately, so that the product of the H⁺ and OH^- concentrations remains equal to 1 x 10^-14.

In an acid solution, the concentration of H⁺ ions is greater than the concentration of OH^- ions. Acid solutions have a sour taste and are able to attack many substances chemically, such as most metals. The greater the concentration of H⁺ ions relative to OH^- ions, the more acidic the solution. If the concentration of H⁺ ions is lower than the concentration of OH^- ions, the solution is basic. Basic solutions have a slippery feel and are also highly reactive, particularly against organic molecules (the slippery feel of basic solutions is the result of conversion of organic oils or greases on the fingers to soaps). When the concentrations of H⁺ and OH^- ions in a solution are equal (and therefore each equal to 1 x 10^-7), the solution is said to be neutral: neither acidic nor basic.

Acids and bases are characterized as strong or weak depending on the degree to which they release or take up H⁺ ions. Nitric acid (HNO₃) is an example of a strong acid. When nitric acid is placed in water, the hydrogen forming part of its molecular structure tends to separate from the rest of the molecule: HNO₃ ↔ H⁺ + NO₃^-

This separation, or dissociation, as it is called, is so complete that the number of H⁺ ions released into the solution is essentially equivalent to the number of nitric acid molecules added. The nitrate ion (NO₃^-) also produced by this dissociation is actually a base, since it can take up an H⁺ ion to form nitric acid again (the double arrow used in the reaction indicates that the reaction can theoretically go either to the right or to the left). The nitrate ion is a very weak base, however, so that its ability to take up an H⁺ ion is very limited. Strong acids typically dissociate into one or more H⁺ ions and a weak base in this manner.

A weak acid, in contrast, has only a limited ability to release H⁺ ions when placed in water solution. Acetic acid (CH₃COOH), for example, dissociates only to a limited extent to release H⁺ ions: CH₃COOH ↔ H⁺ + CH₃COO^-

Dissociation of the acid is limited because the acetate ion (CH₃COO^-) is a relatively strong base and tends to combine with most of the available H⁺ ions to form the undissociated, or CH₃COOH, form of acetic acid again.

Weak acids typically dissociate into one or more H⁺ ions and a strong base in this manner. Although the base produced when a weak acid dissociates is relatively strong and takes up many of the available H⁺ ions, enough H⁺ ions are left in the solution to overbalance the number of OH^- ions, and the solution still has an acid reaction.

The base produced when an acid dissociates to release one or more H⁺ ions is called the conjugate base of the acid. Strong acids therefore dissociate to form a weak conjugate base, and weak acids dissociate to form a strong conjugate base. A similar relationship exists between strong and weak bases: A strong base dissociates to form a weak conjugate acid, and a weak base dissociates to form a strong conjugate acid.

The relative degree of acidity produced by the addition of a strong or weak acid to water depends simply on the number of H⁺ ions released to the solution per number of acid molecules added. The H⁺ ions released by nitric acid, for example, are no "stronger," individually, than those released by acetic acid.

The limited dissociation of weak acids and bases gives these substances the ability to act as buffers, that is, substances that tend to resist changes in the acidity of a solution. For example, if additional H⁺ ions are added to a solution of acetic acid, in which the acid has partially dissociated to form acetate ions, many of the added H⁺ ions combine with the acetate ions to form acetic acid again, and are effectively removed from the solution. The ability of the acetate ions to take up the added H⁺ therefore tends to keep the solution at approximately the same degree of acidity. Conversely, if H⁺ ions are removed from an acetic acid solution, additional acetic acid molecules tend to dissociate, releasing more H⁺ ions to compensate for those removed.

The ability of weak acids or bases to act as buffers is limited, and they can be "swamped out" if the number of H⁺ ions added or released is too great to be compensated for by additional dissociation or reassociation of acid or base molecules. The ability of a weak acid or base to act as a buffer is greatly expanded if a salt of the acid or base is added to the solution. Sodium acetate, for example, is a salt of acetic acid. When added to a solution containing acetic acid, sodium acetate dissociates almost completely into sodium and acetate ions. The additional acetate ions add greatly to the ability of the system to take up or release H⁺ ions.

Applications

Acids and bases are highly reactive substances because either H⁺ or OH^- ions, when placed in excess in a solution through the presence of an acid or base, have the ability to "attack" chemical bonds in other substances. The attack, partly the result of the strongly negative or positive charge of these ions, attracts or repels electrons holding the atoms of other substances together. As a result, the chemical bonds holding the atoms together are weakened, promoting their breakage and leading to disassembly, or breakdown, of the original substances attacked, and often to interaction of the chemical groups produced by the breakdown to form new substances.

The presence of H⁺ or OH^- ions in excess also sets up conditions of acidity that can speed or reduce the rate at which other substances react. The degree of acidity and its control are particularly important in living organisms because complex organic molecules such as proteins and nucleic acids undergo drastic changes in structure and activity with changes in acidity. In general, these complex organic molecules reach and attain maximum activity only over the narrow range of acidity typical of body fluids in an organism. If the degree of acidity changes to values outside this optimal range, the activities of vital molecules are significantly altered. If the degree of acidity strays too far from optimal values, the effects can be quickly lethal. The importance of the degree of acidity or basicity to living organisms makes control of this condition absolutely necessary in biochemical experiments conducted with living systems or biological molecules.

The levels of acidity of water in the atmosphere and in streams, lakes, and the oceans is also highly critical to the quality of the environment. Atmospheric or surface water that becomes too acidic can have seriously detrimental effects on plant and animal life. Acid rain or surface water can also attack and dissolve mineral deposits such as limestone, leach and weaken concrete structures, and destroy buildings or works of art constructed from marble and many other types of natural and artificial materials.

Excessive levels of acidity in atmospheric water, rainwater, and natural bodies of water are largely the result of the effects of pollution from combustion of fossil fuels in automobiles, utilities, and factories. Sulfur present as an impurity in coal and oil produces sulfur dioxide (SO₂) when these fuels are burned. Within a short time, sulfur dioxide is converted by exposure to atmospheric oxygen into sulfur trioxide (SO₃). This substance quickly interacts with atmospheric water to form sulfuric acid. This strong acid contributes to the lung irritation caused by smog. Washing of the sulfuric acid from the atmosphere during rainstorms is one of the primary sources of acid rain.

The burning of coal and oil, and of gasoline or oil in automobile and other internal-combustion engines, also converts some of the nitrogen in Earth's atmosphere into nitrous oxide (NO). This substance slowly reacts with atmospheric oxygen to form nitrogen dioxide (NO₂). In addition to lending a brownish color to smog, nitrogen dioxide reacts with atmospheric water to form nitric acid. This strong acid is a major irritant in smog, and contributes to acid rain by two mechanisms. One is simply by being washed from the atmosphere during rainstorms; the other is its effect in greatly increasing the rate at which SO₂ is converted to SO₃, which reacts with atmospheric water to form sulfuric acid.

The importance of acidity and basicity to industry, living organisms, and the environment makes it essential to have means of evaluating and measuring these conditions. The degree of acidity or basicity of a solution is evaluated by relating it quantitatively to a number scale of 1 to 14, called the pH scale. Assignment of the degree of acidity to the pH scale is carried out by taking the negative logarithm (to the base ten) of the concentration of H⁺ ions in moles per liter in a solution. For example, in a neutral solution, H⁺ ions are present at a concentration of 0.0000001 mole (1 x 10^-7 mole). The log sub 10 of this concentration is -7, and the negative of the logarithm -7 is 7. Thus, a neutral water solution with H⁺ ions at a concentration of 1 x 10^-7 mole has a pH of 7. In general, if H⁺ ions are present in a concentration of 1 x 10^-3 mole, the pH is 3; if present in 1 x 10 to the -4 mole, the pH is 4, and so on. Solutions with pH higher than 7 have OH^- ions in excess and are basic, or alkaline; solutions with pH less than 7 have H⁺ ions in excess and are acidic.

Because pH values are logs to the base ten of the concentration, a change of one pH unit represents a concentration difference of ten times. Therefore, a solution of pH 2 has ten times as many H⁺ ions, and ten times fewer OH^- ions, than a solution at pH 3. The pH scale is thus similar to the Richter scale used to measure the strength of earthquakes, in which an earthquake falling at 4.0 on the Richter scale is ten times stronger than a 3.0 earthquake.

The positions taken on the pH scale by solutions or liquids falling within common experience give an idea of the relationship of the degree of acidity to the scale. The gastric juice in the stomach, with a pH of 1.4, is highly acidic; lemon juice and orange juice, with pHs of 2.1 and 2.8 respectively, are almost as acidic. Wine and tomato juice typically have pHs of 3.5 and 4.1, respectively; urine, with a pH of about 6, is slightly acidic. Milk, with a pH of 6.9, is only very slightly acidic. Absolutely pure water is neutral, or at pH 7; rainwater typically has a somewhat lower pH. When highly contaminated with pollution, as in acid rain, rainwater's pH may be as low as 2.5.

The fluids surrounding the cells in human bodies are neutral or very slightly basic, with pHs ranging from about 7.0 to 7.4. The pH of blood, for example, is 7.4. Seawater and baking soda, with pHs of about 8.2 and 8.5, respectively, are somewhat more basic; a soap solution has a pH of about 10.2. Limewater and milk of magnesia have a pH of about 10.5, and most household detergents, with a pH of about 11.2, are somewhat more basic. Household ammonia, with a pH of 11.9, is among the most basic solutions of common experience. Strong acids or bases that register with pHs at the ends of the scale, such as sulfuric acid, nitric acid, hydrochloric acid, and sodium hydroxide (lye), are highly caustic substances that must be handled with extreme care.

The pH of solutions can be directly measured by several means. One of the most simple and convenient is by the use of indicators such as litmus. These substances take on characteristic colors when placed in solutions of different pH. Paper strips containing litmus, for example, are colorless at pH 7, red in acid pHs, and blue in basic pHs. A wide variety of different indicators are available, each with colors that change at particular points on the pH scale. When several indicators are used in succession, the pH of a solution can be determined with a fair degree of accuracy. More precise measurements are made with electronic instruments that directly measure the H⁺ ion concentration.

Acidity or basicity is also measured chemically by titration. In this technique, a standard solution of an acid or base, containing a known number of H⁺ or OH^- ions, is added drop by drop in precisely measured quantities to a solution of unknown pH until the unknown solution becomes neutral. The amount of the standard solution required to bring the unknown to neutrality, because the number of H⁺ or OH^- ions added can be precisely determined, gives a direct measure of the degree of acidity of the original unknown solution.

Context

The reactions of solutions and substances as acids or bases have been known since ancient times. The word "acid" comes from the Latin acidus, meaning "sour." Sour substances such as vinegar were well known to the ancient Greeks and Romans. Ancient societies also knew that the sour quality of substances like vinegar could be eliminated by adding other substances, such as limewater and potash (potassium hydroxide), later known as bases. "Potash" refers to the production of this substance by roasting wood or other plant materials to an ash in iron pots. The word "alkali," an older word for "base," is derived from the Arabic word for potash, al-qalay (qalay means "to roast in a pot").

As chemistry developed, an acid came to be known as a substance that has a reactive hydrogen, and a base as a substance that has a reactive hydroxyl group, meaning an oxygen and hydrogen (OH) attached at some point. In the late 1800s, chemist Svante August Arrhenius contributed the idea that an acid splits, or dissociates, when placed in water, to release an H⁺ ion. It was also proposed that the distinguishing feature of a base is its ability to release an OH^- ion. This definition proved to be inadequate, however, because some substances capable of acting as bases, such as ammonia (NH₃), do not contain a hydroxyl group and cannot release OH^- ions.

The definition of acids and bases was made more inclusive in 1923 by Johannes Nicolaus Brønsted and Thomas M. Lowry. These chemists proposed independently that acids are substances that release an H⁺ ion, or proton, and that bases are substances that can accept an H⁺ ion, or proton. They also proposed the concept that an acid, on dissociating to release an H⁺ ion, also forms a conjugate base, and that bases form a conjugate acid when dissociating. Their proposal successfully describes the activities of most known acids and bases in water solutions.

Principal terms

ACID RAIN: rainfall made acidic by dissolved acids from atmospheric pollution

BUFFER: a substance that resists changes in acidity or pH

DISSOCIATION: the separation of a molecule into ions

INDICATOR: a substance that undergoes color changes related to acidity or pH

ION: a charged atom or group of atoms

PH: a measure, on a scale of 1 to 14, of the degree of acidity; 1 is most acid, 7 is neutral, and 14 is most basic

PROTON: an atomic particle carrying a positive charge; the nucleus of a hydrogen atom, and a hydrogen ion, consists of a single proton

TITRATION: determination of the acidity or basicity of a solution by measuring the quantity of a standard acid or base required to bring the solution to a neutral pH

Bibliography

"Acids and Bases." BBC Bitesize, www.bbc.co.uk/bitesize/guides/zyn3b9q/revision/1. Accessed 17 Oct. 2024.

Barrow, Gordon. General Chemistry. Wadsworth, 1972.

Brookshire, Bethany. "Let's Learn about Acids and Bases." Science News Explores, 23 Feb. 2021, www.snexplores.org/article/lets-learn-about-acids-and-bases. Accessed 17 Oct. 2024.

Clark, W. Mansfield. Topics in Physical Chemistry. 2nd ed., Williams & Wilkins, 1952.

Dickerson, Richard E., and Irving Geis. Chemistry, Matter, and the Universe. W. A. Benjamin, 1976.

Miller, G. Tyler, Jr. Chemistry: A Contemporary Approach. Wadsworth, 1976.

Smith, Michael B. Organic Chemistry: An Acid-Base Approach, 3rd ed., CRC, 2023.