CHEMISTRY. -
I. DEFINITIONS AND DIVISIONS
The word chemistry may derive from the Egyptian or Coptic terms «km.t» or «Kèmi», meaning «black earth» or «Egypt»: chemistry would be the science of Egypt, the region where it is said to have been first cultivated. The Arabs added the Arabic article «al», whence the name alchemy by which this science was designated until the early modern period, when it came to be called simply chemistry again.Chemistry is the science that studies the specific properties and substantial transformations of all material bodies insofar as they are such: it makes exception for nuclear transformations, which belong to physics because they are studied using its methods.
It is divided into two branches. The specific properties and transformations of substances, without reference to particular substances or reactions, are studied in a general way in physical chemistry or theoretical chemistry; in addition to the means proper to chemistry, it employs
The study of mathematics and physics. The properties and transformation processes concerning individual substances or each group of substances are examined in systematic descriptive general chemistry; for didactic and practical reasons, this is divided into inorganic or mineral chemistry and organic or carbon compound chemistry. When various substances or particular groups of them are studied with more restricted aims or in specific natural or artificial associations, chemistry gives rise to the numerous branches of its applications: astrochemistry, geochemistry, biochemistry; industrial, agricultural, analytical, pharmaceutical, bromatological, metallurgical, photographic, dye, textile chemistry, etc.
II. LAWS OF CHEMICAL COMBINATIONS; ATOMIC-MOLECULAR HYPOTHESIS
Nature has provided man with a great variety of substances. Some of these can even be found in a pure state, such as diamond, quartz, and other crystalline substances; almost all other bodies, however, are more or less homogeneous mixtures or more or less organized associations of various chemical species. First, prehistoric man, then more cultured peoples, with greater effort in the Middle Ages by the alchemists and in the modern age by chemists, made numerous attempts to modify bodies, especially through the action of heat. Thus, changes of state (evaporation, condensation, fusion) and other predominantly physical transformations (mixing, dissolution, crystallization, drying) were obtained; sometimes chemical species were separated in a state of relative purity (sodium chloride by evaporation of seawater, sulfur by sublimation of crude ore, pure water by distillation of natural water, alcohol by distillation of wine). At other times, technical operations based on chemical reactions were performed, such as the preparation of cast iron by heating iron ores with coal, copper by a similar operation on copper ores, lime by prolonged heating of limestone (prehistoric times), and glass by fusing siliceous sand with plant ashes and lime (ancient Chinese and Egyptian civilizations).Chemical transformations received a proper interpretation when their quantitative aspect began to be considered, that is, when the weight ratios according to which bodies combine were taken into account, and, when gaseous substances participate or are formed, also the volume variations and the ratios between gas volumes.
It was thus observed that some substances, when subjected to physical and chemical means, no longer transformed, or gave products always of greater weight (combinations, syntheses), but never products of lesser weight than the starting material (decompositions, analyses): certain substances, that is, reveal the possible ability to enter into combination but not to yield simpler products. Thus, chemical elements were discovered—those simple substances that cannot be further decomposed or simplified, from whose combinations all other pure substances are formed, and from these all bodies.
Combinations between elements do not occur in any manner but obey laws, some of which, the stoichiometric ones, were discovered at the dawn of chemistry:
Principle of the conservation of mass or weight (Lavoisier, 1743–94). The sum of the masses (or weights) of the substances resulting from a chemical reaction is equal to the sum of the masses (or weights) of the substances that reacted. This principle is also called the conservation of matter; but improperly, since the scientist observes only the physical and chemical properties of matter, not its essence or complete definition, which belongs to philosophy. It is also expressed in the following materialistic formulations: “nothing is created and nothing is destroyed, but everything is transformed”; or “the quantity of matter contained in the universe is constant,” though with less precision; indeed, the concepts of creation and destruction are not the object of positive experience; nor are there sufficient elements to establish a law referring to the entire universe, since only an extremely small fraction of it is known.
Principle of definite proportions (L. G. Proust, French chemist, 1755–1826). Pure bodies exhibit constant composition.
Law of multiple proportions (G. Dalton, English chemist, 1766–1844). In a series of pure substances that share two elements in common, the quantities of one element combined with a fixed quantity of the other are in simple ratios to one another.
Law on combinations of gaseous substances (G. Gay-Lussac, French chemist, 1778–1850). Gases or vapors combine with one another in simple volume ratios when measured under equal conditions of temperature and pressure.
It was later found that these laws are not exactly observed in every case, for different reasons peculiar to each law. The principle of the conservation of mass is limited by possible transformations of a fraction of mass into energy (nuclear reactions, radioactive transformations); the laws of definite and multiple proportions are limited by the possibility of the existence of chemical combinations (e.g., natural iron sulfides of the reported type and pyrite) in which, within certain limits, two elements are combined in ratios that are not simple; the law on gaseous combinations is limited by the properties of real gases, which deviate from those of perfect gases.
The knowledge and application of these laws paved the way for the atomic hypothesis formulated by G. Dalton in 1804: 1) there exist as many different kinds of atoms as there are elements constituting matter; 2) all atoms of the same element have the same weight. This hypothesis was not yet sufficient to calculate the relative weights of atoms from the weight ratios according to which elements combine with one another. To achieve this, Avogadro’s intuition (1776–1856) was needed; in various writings from 1811 to 1838, he proposed the famous hypothesis that bears his name: “equal volumes of gas under the same conditions of temperature and pressure contain an equal number of molecules.” His writings, however, remained almost unknown and thus unproductive for the progress of science until they were studied and valued by Cannizzaro (1826–1910), who succeeded in drawing the attention of scientists to them at the Karlsruhe Congress in 1860. His work can be summarized in the following “Cannizzaro’s rule,” which allows the determination of the atomic weight of an element when the molecular weight of its compounds is known: “by the atomic weight of an element is meant the smallest quantity of it contained in the molecular weight of its compounds.”
Once the elementary percentage composition and the molecular weight were determined, and the atomic weights were known, the number of atoms of each kind constituting the molecule became known—that is, the chemical formulas that summarize all this knowledge in an extremely simple and complete manner. A deeper study of the properties of substances and, in particular, the observation that certain groupings of atoms passed unchanged from one substance to another in reactions (the theory of radicals) led to the establishment of structural formulas, which summarize in an equally simple, though less certain, manner not only the composition but also some chemical properties of the substances represented.
III. PERIODIC CLASSIFICATION OF THE ELEMENTS
Just as in the 18th century for the descriptive natural sciences, so in the first half of the last century there were early attempts at classifying the chemical elements; but it was only the Russian chemist Mendeleev who, in 1869, succeeded in proposing a classification thatFew modifications were made before it was definitively incorporated into science. Although constructed on empirical criteria, it was later seen, at the beginning of our century, that this classification had a rational basis, because the properties of the elements, and thus their position in the classification, were a direct consequence of the electronic structure of the atoms of which the elements are composed.
Assuming familiarity with the physical concepts contained in the entry MATTER (q.v.), a concise illustration of Mendeleev’s table will be provided here.
It should be noted that the classification is based first on knowledge of the atomic weights of the elements and then on their physical and chemical properties and those of the compounds derived from them; among these properties, the characteristic valences of each element are particularly important.
By valence is meant the number of hydrogen atoms with which an atom of the element combines or which it replaces in its various compounds. The same element can exhibit different valences depending on the compounds in which it participates. As an example, consider three of the simplest compounds of nitrogen. In ammonia, where one nitrogen atom combines with three hydrogen atoms, nitrogen acts as trivalent; in nitric oxide, where it combines with one oxygen atom and thus replaces two hydrogen atoms, it acts as divalent; in nitric acid, where it replaces five hydrogen atoms (from three water molecules), nitrogen acts as pentavalent:
\[\begin{array}{ccc}
\text{H} & \text{H} & \text{H} \\
\end{array}\]
To highlight the periodic variations with increasing atomic weight, the following should be considered: the maximum valence that the element exhibits in its combinations with oxygen, and, for nonmetals alone, the maximum valence with respect to hydrogen (metals either do not combine with hydrogen or form unstable or difficult-to-prepare compounds).
With this in mind, arrange the symbols of the chemical elements in a horizontal line in order of increasing atomic weight, starting from He, H, Li, etc. (q.v., below, table 1). Up to fluorine, the elements all have distinctly different properties from one another; but upon reaching neon, it is found that it is always inert: it does not combine with hydrogen, oxygen, or any other element, precisely like helium; it is therefore written in the same column beneath II. Then come sodium, magnesium, etc., whose properties are undoubtedly similar to lithium, beryllium, etc., respectively; and so the process continues until the third row. From titanium onward, the similarities are no longer so marked (Mn, for example, is a metal that does not combine with hydrogen, whereas chlorine is a gas that forms a stable compound with hydrogen, hydrochloric acid), but formal similarities remain, particularly in the maximum valence with respect to oxygen; this is taken into account by placing the element in the relevant column but shifted to the left. After manganese, there is no inert element, but the metals Fe, Ni, and Co, which cannot be written under Ar or K; they are therefore placed consecutively on the same row.
Due to clear similarities between Zn and Mg, As and P, Se and S, Br and Cl—more pronounced from left to right—the element copper is also placed in the same column as Li, Na, and K, even though the similarities between it and these are barely discernible; this is accounted for by writing it shifted to the right. With these precautions, the process continues up to lanthanum. From cerium to lutetium, periodicity undergoes a long pause: in fact, for the rare-earth metals, which have very similar properties, there is no place in the table, and they are written separately. Periodicity resumes with hafnium and then continues without other notable peculiarities up to uranium.
Due to its valence similarity with the alkali metals Li, Na, etc., hydrogen can be removed from the first row and placed above lithium or above fluorine. The table is thus constructed: it should only be noted that, to preserve evident similarities, Ni and Co, K and Ar, I and Te have been inverted, making an exception to the progression of atomic weights in favor of the properties.
All the elements are thus classified:
- into 6 periods, of which the first two, from He to F and from Ne to Cl, are small, with 8 elements each; the next three are large, with 18 elements each, plus the 14 rare-earth elements for the fifth; the sixth is interrupted with uranium and remains with only 6 terms;
- into 9 groups, from 0 to 8, plus the group of inert elements; group 8 contains 3 elements in each large period; the remaining groups, from 1 to 7, are subdivided, only in the large periods, into two subgroups each.
At present, the periodic system of elements is preferably presented as in the table in cols. 1541–42, in which:
a) the numbering of the periods changes, with H and He belonging to the new 1st period, while the old 1st period, which begins with Li and no longer with Ne, is designated as the 2nd period; a similar shift in numbering applies to the subsequent periods. This is justified by the convenience of aligning the periodic classification with the electronic structure of atoms;
b) the group of noble gases is placed on the right side of the table, at the end of each period rather than at the beginning of the next;
c) the maximum valence of individual elements with respect to oxygen is highlighted by indicating the formula of the typical oxides;
d) the subdivision into subgroups is given greater emphasis;
e) the rare earths are listed separately and in the table are simply indicated by the expression “rare earths”;
f) for individual elements, the following are highlighted: the atomic number, the symbol, the atomic weight, and the number—though not always certain—of isotopes (q.V. MATTER) found in nature.
Element 43 is poorly known: it has successively borne the names of Masurium (Ma) and Technetium (Tc). Element 61 has successively borne the names of Illinium (Il), Florencium (Fi), and Promethium (Pm). Lutetium (71) is also called Cassiopeium (Cp). Radon (86) is also called Emanation (Em) or Niton (Nt). Element 87, poorly known,
Internal and therefore very stable: their decomposition into elementary components requires a great expenditure of energy.
The elements that combine to form very stable substances possess a great reciprocal affinity. This is a term that comes to us from alchemy, already used by St Albert the Great to designate the union between similar substances in conformity with the axiom of the ancients: *similia similibus*. The term was retained because even at the dawn of chemistry it was believed that similar or akin elements unite to form combinations in preference to dissimilar ones. It was later seen that this could be true for some categories of substances, while for others the opposite was true, stable combinations being formed between elements with opposite characteristics, such as sodium and chlorine, calcium and oxygen. But the word affinity has remained to indicate the tendency of elements to enter into combination and, in general, of chemical substances to react with one another.
It was M. Berthelot, with his celebrated law of maximum heat evolution (1867), who provided a criterion for measuring affinity: “Chemical transformations that occur spontaneously, i.e. without the contribution of energy supplied from outside, tend to produce the system of substances whose formation is accompanied by the greatest evolution of heat.”
The rule indicated by Berthelot worked well in many cases, but it was not of general application and could therefore not constitute a natural law. The difficulty was resolved by J. van’t Hoff, who replaced heat as a measure of affinity with the maximum work that can be obtained from the reaction when it is conducted under ideal conditions of reversibility. One of the most practical ways—though possible in only a few cases—to achieve reversible processes, and thus obtain maximum work and hence a measure of affinity, consists in transforming the chemical energy of the reaction into electrical energy, which occurs in electric cells.
It is important to note that the affinity of a chemical process depends not only on the nature of the reaction but also on the concentrations of the reacting and resulting substances and on the temperature, so that even these circumstances determine the direction in which a reaction proceeds. The study of the relationships between the concentration of substances, their specific heat, temperature, energy exchanges, and affinity is perhaps the finest chapter of physical chemistry, entirely founded on a fruitful application of the second law of thermodynamics. Further progress in this field was made by Nernst in 1906 with the formulation of his theorem, or third law of thermodynamics, which specifies the relationships between thermal level and maximum work at low temperatures, in the vicinity of absolute zero.
V. CHEMICAL REACTIONS IN LIVING ORGANISMS AND THE SECOND LAW OF THERMODYNAMICS
The question of vital force agitated chemists and biologists of the 18th and 19th centuries: they wondered whether the production of organic substances required a force different from those at play in ordinary physical and chemical processes. This was resolved experimentally with Wöhler’s synthesis of urea in 1826, starting from an inorganic substance, ammonium cyanate, followed by the synthesis of numerous other organic substances. Today it is held that even the most complex and essential chemical reactions for life obey the same laws as all other chemical reactions; yet science still cannot explain that aspect of biochemistry (q.v.) by which reactions are initiated, regulated, and utilized for the growth and preservation of the individual or the species.While the question has thus been resolved in qualitative terms, as stated, with the synthesis of countless organic substances (about 1 million up to 1940–49), other aspects of the question remain undefined. A topic that has been raised even in recent times is the following: does life obey the second law of thermodynamics (q.v.)?
Consider a higher plant that, using two low-energy products, water and carbon dioxide, and small quantities of other substances, constructs an extremely differentiated organism composed of high-energy substances such as cellulose, starch, sugars, fats, proteins, lignin, etc. At the end of the growth process, the radiant energy consumed—energy theoretically fully convertible into work—is found in three forms: some is stored in the plant as chemical energy of organic substances, part is mobile energy convertible into work, and part is degraded energy no longer convertible into work but only into heat. The remainder has been dissipated into the environment as heat during the growth process: overall, there has been a degradation of energy. A plant, while obeying the first law of thermodynamics (or conservation of energy), would contravene the second law if it were capable of constructing its tissues from carbon dioxide and water by utilizing, at least in part, the heat at uniform temperature of the environment in which it lives—which, as seen from the example given, does not occur in higher plants and has not been demonstrated even in lower forms of life.
Science cannot explain life, a complex phenomenon one aspect of which lies beyond its competence; it can only affirm that the individual elementary physical and chemical processes, from which arises the aspect of life accessible to it, obey the laws of physical chemistry.
VI. CONSTITUTION OF MOLECULES
Modern theoretical chemistry, having resolved stoichiometric and energetic questions, now concerns itself with the structure of molecules and the relationships between molecular constitution and the physical and chemical properties of substances.Numerous quantitative relationships are known, especially those derived from the application of the second law of thermodynamics. They link together some of the following quantities: 1) the temperature at which the chemical process occurs, 2) the maximum work it can yield, 3) the equilibrium constant of the reaction, 4) the variation of this constant with changing temperature, 5) the latent heat of fusion and 6) the cryoscopic constant, 7) the latent heat of vaporization and 8) the ebullioscopic constant, 9) the specific heat of the substances participating in the physical or chemical process, etc.
Such relations, essentially physical in nature, do not provide sufficient material for constructing a deductive, rational science, but they cover a modest sector of chemistry, which still retains some characteristics of a descriptive and empirical science.
It must be acknowledged, however, that the results obtained from experimentation with physical and chemical methods on substances and chemical processes represent today an imposing body of knowledge, both from a speculative and a practical standpoint.
By irradiating crystals with X-rays of known wavelength and examining the diffraction patterns formed by these rays upon exiting the crystal, it has been possible to measure the exact distances between atoms in the molecule and the minimum distances to which atoms of contiguous molecules can approach (q.V. figures).
By using infrared, visible, and ultraviolet light
Chemistry — Model of the benzylpenicillin molecule.
By means of violet (see optics) light, absorption spectra, fluorescence spectra, and Raman spectra have been obtained, which allow the measurement of the various molecular electronic levels, the oscillation frequency of atoms and atomic groups, the rotational constants of molecules, etc. And since in certain cases it is possible to calculate these values when the following are known: the arrangement of the atoms in the molecule, the atomic weights, and the magnitude of the interatomic forces that oppose the displacement of the atoms from their equilibrium positions, useful information has been obtained on the actual structure of molecules by assimilating them to models.
From the measurement of the dielectric and magnetic constants (see electrology) of substances and from the measurement of the variations they undergo under the action of electric and magnetic fields respectively, the electric and magnetic moments of the molecules have been calculated, and useful deductions have been drawn regarding their shape, the distribution of electric charges, and the shape of magnetic fields.
The results of such and other numerous investigations have in many cases provided satisfactory answers to certain questions concerning the structure of molecules and the properties of substances.
When molecules (or ions) combine to form crystals, they solve a "minimum problem": in accordance with the shape of the molecules and the intermolecular forces, they assume the arrangement that corresponds to the minimum internal energy and often, though not always, to the minimum volume.
When atoms combine to form molecules, they must satisfy a greater number of conditions, not all of which are well defined; among these are valence, the direction of valences, affinity, and stability. These are nothing more than different aspects under which the properties of atoms manifest themselves.
It has already been said of valence in connection with the periodic classification of the elements. The three hydrogen atoms in the ammonia molecule, cited as an example at that time, do not assume any arbitrary position around the nitrogen atom, nor do they assume a position of maximum symmetry, i.e., the three vertices of an equilateral triangle with the nitrogen atom at the center; instead, they occupy the base vertices of a triangular pyramid, the fourth vertex of which is occupied by the nitrogen atom, so that the hydrogen atoms do not surround the nitrogen atom but are all on one side. In water (see), the two hydrogen atoms do not lie on a straight line with the oxygen atom; rather, the first two form an angle of 104°–31′ with the latter. In methane, on the other hand, the four hydrogen atoms are arranged around the carbon atom in positions corresponding to maximum symmetry, i.e., the vertices of a regular tetrahedron with the carbon atom at the center.
Affinity, already mentioned, considered as the greater or lesser tendency that equal or different atoms have to combine to form a true chemical bond, varies first of all with the nature of the atoms but also depends on other circumstances: in particular, on the type of bond, as will be discussed below, and on the nature and arrangement of the other atoms in the same molecule. It is measured by the energy required to separate the two atoms to infinity, i.e., to break the bond, while leaving the other conditions unchanged. Some typical metalloids (H, N, O) in the elementary state form very stable diatomic molecules and thus constitute examples of very energetic bonds between two identical atoms. Carbon is the only element whose atoms have a marked tendency to combine in even very long chains or in rings or systems of rings, even very complex ones; a less pronounced tendency in this direction is also possessed by silicon and sulfur. Oxygen has the capacity to combine with all elements, naturally excluding the zero-valent elements of the helium series, which do not form stable compounds with any element.
The stability of chemical combinations depends in particular on temperature and the heat of formation, as well as on other circumstances. Every substance has a limiting temperature of stability above which it decomposes into other, more stable and often simpler substances. The influence of the heat of formation on the stability of substances has already been mentioned: substances that form from the elements with a large evolution of heat are stable at ordinary temperatures, whereas those that form with an absorption of heat are unstable. The thermal effect varies with the temperature at which the reaction takes place.
The forces that hold atoms together in molecules are of various types; the most important are those corresponding to electrovalency and covalency. An electric valence or typical heteropolar bond occurs when one of the two atoms transfers an electron to the other:
Chemistry — Interatomic distances (in Angstroms) in benzylpenicillin molecules.
Between the two, the first positive and the second negative, an electrostatic attraction is established: this is the case of sodium chloride, NaCl. A covalence or homopolar bond occurs when the two atoms share one electron each, so that the two atoms remain bound together by a pair of electrons that in a certain way belongs simultaneously to both atoms. The resulting attraction cannot be expressed by means of classical physics schemes but can only be explained with the mathematical tools proper to quantum wave mechanics.
The forces that establish an attraction between molecules in crystals or in liquids, and also in the gaseous state when the molecules approach beyond a certain limit, are not essentially different from those that are established between atoms of the same molecule; but since the electrical charges are largely already neutralized within the molecule and the electrons are already engaged in intramolecular bonds, the attractive forces that can develop are less strong than those between atoms of the same molecule, and equilibrium with the repulsive forces is established at a greater distance. When intermolecular attractions reach the same order of magnitude as intramolecular ones or exceed them, condensations, polymerizations, or other chemical reactions may occur.
VII. FUTURE OF CHEMISTRY — Chemistry is distinguished from physics by having its own formal object and its own research methods; but from a strictly philosophical and systematic point of view, it can also be considered as a branch of physics, inasmuch as the former does not possess problems of its own distinct from those of the latter, as, for example, biology and astronomy do; all the more so since the boundaries between physics and chemistry become increasingly blurred with the progress of research. There will come a time when the present chemistry may give rise to a rational science largely constituted by advanced mathematical procedures; and experience will then have no other purpose than to seek the values to be introduced into these procedures and to verify the predictions they will provide. But there are no signs that this is imminent. Experimental and mathematical means are still lacking to resolve questions that, as they are currently framed, present extraordinary complexity. For many phenomena an explanation can be given, but not an exact quantitative formulation: even though the molecular structure of many substances is known with certainty, the exact values of some of their physical and chemical properties cannot be predicted. As a practical science, however, it can be considered that chemistry has reached a high degree of development: it has provided technology with countless procedures that have contributed profoundly to transforming human social life.