VOLUME VI slice I Châtelet to Chicago (continued 3)

CHEKE, SIR JOHN (1514-1557), English classical scholar, was the son of Peter Cheke, esquire-bedell of Cambridge University. He was educated at St John’s College, Cambridge, where he became a fellow in 1529. While there he adopted the principles of the Reformation. His learning gained him an exhibition from the king, and in 1540, on Henry VIII.’s foundation of the regius professorships, he was elected to the chair of Greek. Amongst his pupils at St John’s were Lord Burghley, who married Cheke’s sister Mary, and Roger Ascham, who in The Schoolmaster gives Cheke the highest praise for scholarship and character. Together with Sir Thomas Smith, he introduced a new method of Greek pronunciation very similar to that commonly used in England in the 19th century. It was strenuously opposed in the University, where the continental method prevailed, and Bishop Gardiner, as chancellor, issued a decree against it (June 1542); but Cheke ultimately triumphed. On the 10th of July 1554, he was chosen as tutor to Prince Edward, and after his pupil’s accession to the throne he continued his instructions. Cheke took a fairly active share in public life; he sat, as member for Bletchingley, for the parliaments of 1547 and 1552-1553; he was made provost of King’s College, Cambridge (April 1, 1548), was one of the commissioners for visiting that university as well as Oxford and Eton, and was appointed with seven divines to draw up a body of laws for the governance of the church. On the 11th of October 1551 he was knighted; in 1553 he was made one of the secretaries of state, and sworn of the privy council. His zeal for Protestantism induced him to follow the duke of Northumberland, and he filled the office of secretary of state for Lady Jane Grey during her nine days’ reign. In consequence Mary threw him into the Tower (July 27, 1553), and confiscated his wealth. He was, however, released on the 13th of September 1554, and granted permission to travel abroad. He went first to Basel, then visited Italy, giving lectures in Greek at Padua, and finally settled at Strassburg, teaching Greek for his living. In the spring of 1556 he visited Brussels to see his wife; on his way back, between Brussels and Antwerp, he and Sir Peter Carew were treacherously seized (May 15) by order of Philip of Spain, hurried over to England, and imprisoned in the Tower. Cheke was visited by two priests and by Dr John Feckenham, dean of St Paul’s, whom he had formerly tried to convert to Protestantism, and, terrified by a threat of the stake, he gave way and was received into the Church of Rome by Cardinal Pole, being cruelly forced to make two public recantations. Overcome with shame, he did not long survive, but died in London on the 13th of September 1557, carrying, as T. Fuller says (Church History), “God’s pardon and all good men’s pity along with him.” About 1547 Cheke married Mary, daughter of Richard Hill, sergeant of the wine-cellar to Henry VIII., and by her he had three sons. The descendants of one of these, Henry, known only for his translation of an Italian morality play Freewyl (Tragedio del Libero Arbitrio) by Nigri de Bassano, settled at Pyrgo in Essex.

Thomas Wilson, in the epistle prefixed to his translation of the Olynthiacs of Demosthenes (1570), has a long and most interesting eulogy of Cheke; and Thomas Nash, in To the Gentlemen Students, prefixed to Robert Greene’s Menaphon (1589), calls him “the Exchequer of eloquence, Sir Ihon Cheke, a man of men, supernaturally traded in all tongues.” Many of Cheke’s works are still in MS., some have been altogether lost. One of the most interesting from a historical point of view is the Hurt of Sedition how greueous it is to a Communewelth (1549), written on the occasion of Ket’s rebellion, republished in 1569, 1576 and 1641, on the last occasion with a life of the author by Gerard Langbaine. Others are D. Joannis Chrysostomi homiliae duae (1543), D. Joannis Chrysostomi de providentia Dei (1545), The Gospel according to St Matthew ... translated (c. 1550; ed. James Goodwin, 1843), De obitu Martini Buceri (1551), (Leo VI.’s) de Apparatu bellico (Basel, 1554; but dedicated to Henry VIII., 1544), Carmen Heroicum, aut epitaphium in Antonium Deneium (1551), De pronuntiatione Graecae ... linguae (Basel, 1555). He also translated several Greek works, and lectured admirably upon Demosthenes.

His Life was written by John Strype (1821); additions by J. Gough Nichols in Archaeologia (1860), xxxviii. 98, 127.

CHELLIAN, the name given by the French anthropologist G. de Mortillet to the first epoch of the Quaternary period when the earliest human remains are discoverable. The word is derived from the French town Chelles in the department of Seine-et-Marne. The climate of the Chellian epoch was warm and humid as evidenced by the wild growth of fig-trees and laurels. The animals characteristic of the epoch are the Elephas antiquus, the rhinoceros, the cave-bear, the hippopotamus and the striped hyaena. Man existed and belonged to the Neanderthal type. The implements characteristic of the period are flints chipped into leaf-shaped forms and held in the hand when used. The drift-beds of St Acheul (Amiens), of Menchecourt (Abbeville), of Hoxne (Suffolk), and the detrital laterite of Madras are considered by de Mortillet to be synchronous with the Chellian beds.

See Gabriel de Mortillet, Le Préhistorique (1900); Lord Avebury, Prehistoric Times (1900).

CHELMSFORD, FREDERIC THESIGER, 1st Baron (1794-1878), lord chancellor of England, was the third son of Charles Thesiger, and was born in London on the 15th of April 1794. His father, collector of customs at St Vincent’s, was the son of a Saxon gentleman who had migrated to England and become secretary to Lord Rockingham, and was the brother of Sir Frederic Thesiger, naval A.D.C. to Nelson at Copenhagen. Young Frederic Thesiger was originally destined for a naval career, and he served as a midshipman on board the “Cambrian” frigate in 1807 at the second bombardment of Copenhagen. His only surviving brother, however, died about this time, and he became entitled to succeed to a valuable estate in the West Indies, so it was decided that he should leave the navy and study law, with a view to practising in the West Indies and eventually managing his property in person. Another change of fortune, however, awaited him, for a volcano destroyed the family estate, and he was thrown back upon his prospect of a legal practice in the West Indies. He proceeded to enter at Gray’s Inn in 1813, and was called on the 18th of November 1818, another change in his prospects being brought about by the strong advice of Godfrey Sykes, a special pleader in whose chambers he had been a pupil, that he should remain to try his fortune in England. He accordingly joined the home circuit, and soon got into good practice at the Surrey sessions, while he also made a fortunate purchase in buying the right to appear in the old palace court (see Lord Steward). In 1824 he distinguished himself by his defence of Joseph Hunt when on his trial at Hertford with John Thurtell for the murder of Wm. Weare; and eight years later at Chelmsford assizes he won a hard-fought action in an ejectment case after three trials, to which he attributed so much of his subsequent success that when he was raised to the peerage he assumed the title Lord Chelmsford. In 1834 he was made king’s counsel, and in 1835 was briefed in the Dublin election inquiry which unseated Daniel O’Connell. In 1840 he was elected M.P. for Woodstock. In 1844 he became solicitor-general, but having ceased to enjoy the favour of the duke of Marlborough, lost his seat for Woodstock and had to find another at Abingdon. In 1845 he became attorney-general, holding the post until the fall of the Peel administration on the 3rd of July 1846. Thus by three days Thesiger missed being chief justice of the common pleas, for on the 6th of July Sir Nicholas Tindal died, and the seat on the bench, which would have been Thesiger’s as of right, fell to the Liberal attorney-general, Sir Thomas Wilde. Sir Frederic Thesiger remained in parliament, changing his seat, however, again in 1852, and becoming member for Stamford. During this period he enjoyed a very large practice at the bar, being employed in many causes célèbres. On Lord Derby coming into office for the second time in 1858, Sir Frederic Thesiger was raised straight from the bar to the lord chancellorship (as were Lord Brougham, Lord Selborne and Lord Halsbury). In the following year Lord Derby resigned and his cabinet was broken up. Again in 1866, on Lord Derby coming into office for the third time, Lord Chelmsford became lord chancellor for a short period. In 1868 Lord Derby retired, and Disraeli, who took his place as prime minister, wished for Lord Cairns as lord chancellor. Lord Chelmsford was very sore at his supersession and the manner of it, but, according to Lord Malmesbury he retired under a compact made before he took office. Ten years later Lord Chelmsford died in London on the 5th of October 1878. Lord Chelmsford had married in 1822 Anna Maria Tinling. He left four sons and three daughters, of whom the eldest, Frederick Augustus, 2nd Baron Chelmsford (1827-1905), earned distinction as a soldier, while the third, Alfred Henry Thesiger (1838-1880) was made a lord justice of appeal and a privy councillor in 1877, at the early age of thirty-nine, but died only three years later.

See Lives of the Chancellors (1908), by J.B. Atlay, who has had the advantage of access to an unpublished autobiography of Lord Chelmsford’s.

CHELMSFORD, a market town and municipal borough, and the county town of Essex, England, in the Chelmsford parliamentary division, 30 m. E.N.E. from London by the Great Eastern railway. Pop. (1901) 12,580. It is situated in the valley of the Chelmer, at the confluence of the Cann, and has communication by the river with Maldon and the Blackwater estuary 11 m. east. Besides the parish church of St Mary, a graceful Perpendicular edifice, largely rebuilt, the town has a grammar school founded by Edward VI., an endowed charity school and a museum. It is the seat of the county assizes and quarter sessions, and has a handsome shire hall; the county gaol is near the town. Its corn and cattle markets are among the largest in the county; for the first a fine exchange is provided. In the centre of the square in which the corn exchange is situated stands a bronze statue of Lord Chief-Justice Tindal (1776-1846), a native of the parish. There are agricultural implement and iron foundries, large electric light and engineering works, breweries, tanneries, maltings and extensive corn mills. There is a race-course 2 m. south of the town. The borough is under a mayor, 6 aldermen and 18 councillors. Area 2308 acres.

A place of settlement since Palaeolithic times, Chelmsford (Chilmersford, Chelmeresford, Chelmesford) owed its importance to its position on the road from London to Colchester. It consisted of two manors: that of Moulsham, which remained in the possession of Westminster Abbey from Saxon times till the reign of Henry VIII., when it was granted to Thomas Mildmay; and that of Bishop’s Hall, which was held by the bishops of London from the reign of Edward the Confessor to 1545, when it passed to the crown and was granted to Thomas Mildmay in 1563. The medieval history of Chelmsford centred round the manor of Bishop’s Hall. Early in the 12th century Bishop Maurice built the bridge over the Chelmer which brought the road from London directly through the town, thus making it an important stopping-place. The town was not incorporated until 1888. In 1225 Chelmsford was made the centre for the collection of fifteenths from the county of Essex, and in 1227 it became the regular seat of assizes and quarter-sessions. Edward I. confirmed Bishop Richard de Gravesend in his rights of frank pledge in Chelmsford in 1290, and in 1395 Richard II. granted the return of writs to Bishop Robert de Braybroke. In 1377 writs were issued for the return of representatives from Chelmsford to parliament, but no return of members has been found. In 1199 the bishop obtained the grant of a weekly market at the yearly rent of one palfrey, and in 1201 that of an annual fair, now discontinued, for four days from the feast of St Philip and St James.

CHELSEA, a western metropolitan borough of London, England, bounded E. by the city of Westminster, N.W. by Kensington, S.W. by Fulham, and S. by the river Thames. Pop. (1901) 73,842. Its chief thoroughfare is Sloane Street, containing handsome houses and good shops, running south from Knightsbridge to Sloane Square. Hence King’s Road leads west, a wholly commercial highway, named in honour of Charles II., and recalling the king’s private road from St James’s Palace to Fulham, which was maintained until the reign of George IV. The main roads south communicate with the Victoria or Chelsea, Albert and Battersea bridges over the Thames. The beautiful Chelsea embankment, planted with trees and lined with fine houses and, in part, with public gardens, stretches between Victoria and Battersea bridges. The better residential portion of Chelsea is the eastern, near Sloane Street and along the river; the western, extending north to Fulham Road, is mainly a poor quarter.

Chelsea, especially the riverside district, abounds in historical associations. At Cealchythe a synod was held in 785. A similar name occurs in a Saxon charter of the 11th century and in Domesday; in the 16th century it is Chelcith. The later termination ey or ea was associated with the insular character of the land, and the prefix with a gravel bank (ceosol; cf. Chesil Bank, Dorsetshire) thrown up by the river; but the early suffix hythe is common in the meaning of a haven. The manor was originally in the possession of Westminster Abbey, but its history is fragmentary until Tudor times. It then came into the hands of Henry VIII., passed from him to his wife Catharine Parr, and thereafter had a succession of owners, among whom were the Howards, to whom it was granted by Queen Elizabeth, and the Cheynes, from whom it was purchased in 1712 by Sir Hans Sloane, after which it passed to the Cadogans. The memorials which crowd the picturesque church and churchyard of St Luke near the river, commonly known as the Old Church, to a great extent epitomize the history of Chelsea. Such are those of Sir Thomas More (d. 1535); Lord Bray, lord of the manor (1539), his father and son; Lady Jane Guyldeford, duchess of Northumberland, who died “at her maner of Chelse” in 1555; Lord and Lady Dacre (1594-1595); Sir John Lawrence (1638); Lady Jane Cheyne (1698); Francis Thomas, “director of the china porcelain manufactory, Lawrence Street, Chelsea” (1770); Sir Hans Sloane (1753); Thomas Shadwell, poet laureate (1602); Woodfall the printer of Junius (1844), and many others. More’s tomb is dated 1532, as he set it up himself, though it is doubtful whether he lies beneath it. His house was near the present Beaufort Street. In the 18th and 19th centuries Chelsea, especially the parts about the embankment and Cheyne Walk, was the home of many eminent men, particularly of writers and artists, with whom this pleasant quarter has long been in favour. Thus in the earlier part of the period named, Atterbury and Swift lived in Church Lane, Steele and Smollett in Monmouth House. Later, the names of Turner, Rossetti, Whistler, Leigh Hunt, Carlyle (whose house in Cheyne Row is preserved as a public memorial), Count D’Orsay, and Isambard Brunel, are intimately connected with Chelsea. At Lindsey House Count Zinzendorf established a Moravian Society (c. 1750). Sir Robert Walpole’s residence was extant till 1810; and till 1824 the bishops of Winchester had a palace in Cheyne Walk. Queen’s House, the home of D.G. Rossetti (when it was called Tudor House), is believed to take name from Catharine of Braganza.

Chelsea was noted at different periods for two famous places of entertainment, Ranelagh (q.v.) in the second half of the 18th century, and Cremorne Gardens (q.v.) in the middle of the 19th. Don Saltero’s museum, which formed the attraction of a popular coffee-house, was formed of curiosities from Sir Hans Sloane’s famous collections. It was Sloane who gave to the Apothecaries’ Company the ground which they had leased in 1673 for the Physick Garden, which is still extant, but ceased in 1902 to be maintained by the Company. At Chelsea Sir John Danvers (d. 1655) introduced the Italian style of gardening which was so greatly admired by Bacon and soon after became prevalent in England. Chelsea was formerly famous for a manufacture of buns; the original Chelsea bun-house, claiming royal patronage, stood until 1839, and one of its successors until 1888. The porcelain works existed for some 25 years before 1769, when they were sold and removed to Derby. Examples of the original Chelsea ware (see Ceramics) are of great value.

Of buildings and institutions the most notable is Chelsea Royal Hospital for invalid soldiers, initiated by Charles II. (according to tradition on the suggestion of Nell Gwynne), and opened in 1694. The hospital itself accommodates upwards of 500 men, but a system of out-pensioning was found necessary from the outset, and now relieves large numbers throughout the empire. The picturesque building by Wren stands in extensive grounds, which include the former Ranelagh Gardens. A theological college (King James’s) formerly occupied the site; it was founded in 1610 and was intended to be of great size, but the scheme was unsuccessful, and only a small part of the buildings was erected. In the vicinity are the Chelsea Barracks (not actually in the borough). The Royal Military Asylum for boys, commonly called the Duke of York’s school, founded in 1801 by Frederick, duke of York, for the education of children connected with the army, was removed in 1909 to new quarters at Dover. Other institutions are the Whitelands training college for school-mistresses, in which Ruskin took deep interest; the St Mark’s college for school-masters; the Victoria and the Cheyne hospitals for children, a cancer hospital, the South-western polytechnic, and a public library containing an excellent collection relative to local history.

The parliamentary borough of Chelsea returns one member, and includes, as a detached portion, Kensal Town, north of Kensington. The borough council consists of a mayor, 6 aldermen and 36 councillors. Area, 659.6 acres.

CHELSEA, a city of Suffolk county, Massachusetts, U.S.A., a suburb of Boston. Pop. (1890) 27,909; (1900) 34,072, of whom 11,203 were foreign-born; (1910) 32,452. It is situated on a peninsula between the Mystic and Chelsea rivers, and Charlestown and East Boston, and is connected with East Boston and Charlestown by bridges. It is served by the Boston & Maine and (for freight) by the Boston & Albany railways. The United States maintains here naval and marine hospitals, and the state a soldiers’ home. Chelsea’s interests are primarily industrial. The value of the city’s factory products in 1905 was $13,879,159, the principal items being rubber and elastic goods ($3,635,211) and boots and shoes ($2,044,250.) The manufacture of stoves, and of mucilage and paste are important industries. Flexible tubing for electric wires (first made at Chelsea 1889) and art tiles are important products. The first settlement was established in 1624 by Samuel Maverick (c. 1602-c. 1670), the first settler (about 1629) of Noddle’s Island (or East Boston), and one of the first slave-holders in Massachusetts; a loyalist and Churchman, in 1664 he was appointed with three others by Charles II. on an important commission sent to Massachusetts and the other New England colonies (see Nicolls, Richard), and spent the last years of his life in New York. Until 1739, under the name of Winnisimmet, Chelsea formed a part of Boston, but in that year it was made a township; it became a city in 1857. In May 1775 a British schooner in the Mystic defended by a force of marines was taken by colonial militia under General John Stark and Israel Putnam,—one of the first conflicts of the War of Independence. A terrible fire swept the central part of the city on the 12th of April 1908.

See Mellen Chamberlain (and others), History of Chelsea (2 vols., Boston, 1908), published by the Massachusetts Historical Society.

CHELTENHAM, a municipal and parliamentary borough of Gloucestershire, England, 109 m. W. by N. of London by the Great Western railway; served also by the west and north line of the Midland railway. Pop. (1901) 49,439. The town is well situated in the valley of the Chelt, a small tributary of the Severn, under the high line of the Cotteswold Hills to the east, and is in high repute as a health resort. Mineral springs were accidentally discovered in 1716. The Montpellier and Pittville Springs supply handsome pump rooms standing in public gardens, and are the property of the corporation. The Montpellier waters are sulphated, and are valuable for their diuretic effect, and as a stimulant to the liver and alimentary canal. The alkaline-saline waters of Pittville are efficacious against diseases resulting from excess of uric acid. The parish church of St Mary dates from the 14th century, but is almost completely modernized. The town, moreover, is wholly modern in appearance. Assembly rooms opened in 1815 by the duke of Wellington were removed in 1901. A new town hall, including a central spa and assembly rooms, was opened in 1903. There are numerous other handsome buildings, especially in High Street, and the Promenade forms a beautiful broad thoroughfare, lined with trees. The town is famous as an educational centre. Cheltenham College (1842) provides education for boys in three departments, classical, military and commercial; and includes a preparatory school. The Ladies’ College (1854), long conducted by Miss Beale (q.v.), is one of the most successful in England. The Normal Training College was founded in 1846 for the training of teachers, male and female, in national and parochial schools. A free grammar school was founded in 1568 by Richard Pate, recorder of Gloucester. The art gallery and museum may be mentioned also. The parliamentary borough returns one member. The municipal borough is under a mayor, 6 aldermen and 18 councillors. Area, 4726 acres. The urban district of Charlton Kings (pop. 3806) forms a south-eastern suburb of Cheltenham.

The site of a British village and burying-ground, Cheltenham (Celtanhomme, Chiltham, Chelteham) was a village with a church in 803. The manor belonged to the crown; it was granted to Henry de Bohun, earl of Hereford, late in the 12th century, but in 1199 was exchanged for other lands with the king. It was granted to William de Longespée, earl of Salisbury, in 1219, but resumed on his death and granted in dower to Eleanor of Provence in 1243. In 1252 the abbey of Fécamp purchased the manor, and it afterwards belonged to the priory of Cormeille, but was confiscated in 1415 as the possession of an alien priory, and was granted in 1461 to the abbey of Lyon, by which it was held until, once more returning to the crown at the Dissolution, it was granted to the family of Dutton. The town is first mentioned in 1223, when William de Longespée leased the benefit of the markets, fairs and hundred of Cheltenham to the men of the town for three years; the lease was renewed by Henry III. in 1226, and again in 1230 for ten years. A market town in the time of Camden, it was governed by commissioners from the 18th century in 1876, when it was incorporated; it became a parliamentary borough in 1832. Henry III. in 1230 had granted to the men of Cheltenham a market on each Thursday, and a fair on the vigil, feast and morrow of St James. Although Camden mentions a considerable trade in malt, the spinning of woollen yarn was the only industry in 1779. After the discovery of springs in 1716, and the erection of a pump-room in 1738, Cheltenham rapidly became fashionable, the visit of George III. and the royal princesses in 1788 ensuring its popularity.

See S. Moreau, A Tour to Cheltenham Spa (Bath, 1738).

CHELYABINSK, a town of Russia, in the Orenburg government, at the east foot of the Urals, is the head of the Siberian railway, 624 m. by rail E.N.E. of Samara and 154 m. by rail S.S.E. of Ekaterinburg. Pop. (1900) 25,505. It has tanneries and distilleries, and is the centre of the trade in corn and produce of cattle for the Ural iron-works. The town was founded in 1658.

CHELYS (Gr. χέλυς, tortoise; Lat. testudo), the common lyre of the ancient Greeks, which had a convex back of tortoiseshell or of wood shaped like the shell. The word chelys was used in allusion to the oldest lyre of the Greeks which was said to have been invented by Hermes. According to tradition he was attracted by sounds of music while walking on the banks of the Nile, and found they proceeded from the shell of a tortoise across which were stretched tendons which the wind had set in vibration (Homeric Hymn to Hermes, 47-51). The word has been applied arbitrarily since classic times to various stringed instruments, some bowed and some twanged, probably owing to the back being much vaulted. Kircher (Musurgia, i. 486) applied the name of chelys to a kind of viol with eight strings. Numerous representations of the chelys lyre or testudo occur on the Greek vases, in which the actual tortoiseshell is depicted; a good illustration is given in Le Antichità, di Ercolano (vol. i. pl. 43). Propertius (iv. 6) calls the instrument the lyra testudinea. Scaliger (on Manilius, Astronomicon, Proleg. 420) was probably the first writer to draw attention to the difference, between chelys and cithara (q.v.).

(K. S.)

CHEMICAL ACTION, the term given to any process in which change in chemical composition occurs. Such processes may be set up by the application of some form of energy (heat, light, electricity, &c.) to a substance, or by the mixing of two or more substances together. If two or more substances be mixed one of three things may occur. First, the particles may be mechanically intermingled, the degree of association being dependent upon the fineness of the particles, &c. Secondly, the substances may intermolecularly penetrate, as in the case of gas-mixtures and solutions. Or thirdly they may react chemically. The question whether, in any given case, we have to deal with a physical mixture or a chemical compound is often decided by the occurrence of very striking phenomena. To take a simple example:—oxygen and hydrogen are two gases which may be mixed in all proportions at ordinary temperatures, and it is easy to show that the properties of the products are simply those of mixtures of the two free gases. If, however, an electric spark be passed through the mixtures, powerful chemical union ensues, with its concomitants, great evolution of heat and consequent rise of temperature, and a compound, water, is formed which presents physical and chemical properties entirely different from those of its constituents.

In general, powerful chemical forces give rise to the evolution of large quantities of heat, and the properties of the resulting substance differ vastly more from those of its components than is the case with simple mixtures. This constitutes a valuable criterion as to whether mere mixture is involved on the one hand, or strong chemical union on the other. When, however, the chemical forces are weak and the reaction, being incomplete, leads to a state of chemical equilibrium, in which all the reacting substances are present side by side, this criterion vanishes. For example, the question whether a salt combines with water molecules when dissolved in water cannot be said even yet to be fully settled, and, although there can be no doubt that solution is, in many cases, attended by chemical processes, still we possess as yet no means of deciding, with certainty, how many molecules of water have bound themselves to a single molecule of the dissolved substance (solute). On the other hand, we possess exact methods of testing whether gases or solutes in dilute solution react one with another and of determining the equilibrium state which is attained. For if one solute react with another on adding the latter to its solution, then corresponding to the decrease of its concentration there must also be a decrease of vapour pressure, and of solubility in other solvents; further, in the case of a mixture of gases, the concentration of each single constituent follows from its solubility in some suitable solvent. We thus obtain the answer to the question: whether the concentration of a certain constituent has decreased during mixing, i.e. whether it has reacted chemically.

When a compound can be obtained in a pure state, analysis affords us an important criterion of its chemical nature, for unlike mixtures, the compositions of which are always variable within wider or narrower limits, chemical compounds present definite and characteristic mass-relations, which find full expression in the atomic theory propounded by Dalton (see Atom). According to this theory a mixture is the result of the mutual interpenetration of the molecules of substances, which remain unchanged as such, whilst chemical union involves changes more deeply seated, inasmuch as new molecular species appear. These new substances, if well-defined chemical compounds, have a perfectly definite composition and contain a definite, generally small, number of elementary atoms, and therefore the law of constant proportions follows at once, and the fact that only an integral number of atoms of any element may enter into the composition of any molecule determines the law of multiple proportions.

These considerations bring us face to face with the task of more closely investigating the nature of chemical forces, in other words, of answering the question: Nature of chemical forces. what forces guide the atoms in the formation of a new molecular species? This problem is still far from being completely answered, so that a few general remarks must suffice here.

It is remarkable that among the most stable chemical compounds, we find combinations of atoms of one and the same element. Thus, the stability of the di-atomic molecule N2 is so great, that no trace of dissociation has yet been proved even at the highest temperatures, and as the constituent atoms of the molecule N2 must be regarded as absolutely identical, it is clear that “polar” forces cannot be the cause of all chemical action. On the other hand, especially powerful affinities are also at work when so-called electro-positive and electro-negative elements react. The forces which here come into play appear to be considerably greater than those just mentioned; for instance, potassium fluoride is perhaps the most stable of all known compounds.

It is also to be noticed that the combinations of the electro-negative elements (metalloids) with one another exhibit a metalloid character, and also we find, in the mutual combinations of metals, all the characteristics of the metallic state; but in the formation of a salt from a metal and a metalloid we have an entirely new substance, quite different from its components; and at the same time, the product is seen to be an electrolyte, i.e. to have the power of splitting up into a positively and a negatively charged constituent when dissolved in some solvent. These considerations lead to the conviction that forces of a “polar” origin play an important part here, and indeed we may make the general surmise that in the act of chemical combination forces of both a non-polar and polar nature play a part, and that the latter are in all probability identical with the electric forces.

It now remains to be asked—what are the laws which govern the action of these forces? This question is of fundamental importance, since it leads directly to those laws which regulate the chemical process. Besides the already mentioned fundamental law of chemical combination, that of constant and multiple proportions, there is the law of chemical mass-action, discovered by Guldberg and Waage in 1867, which we will now develop from a kinetic standpoint.

Kinetic Basis of the Law of Chemical Mass-action.—We will assume that the molecular species A1, A2, ... A′1, A′2, ... are present in a homogeneous system, where they can react on each other only according to the scheme

A1 + A2 + ... ↔ A′1 + A′2 + ...;

this is a special case of the general equation

n1A1 + n2A2 + ... ↔ n′1A′1 + n′2A′2 + ...,

in which only one molecule of each substance takes part in the reaction. The reacting substances may be either gaseous or form a liquid mixture, or be dissolved in some selected solvent; but in each case we may state the following considerations regarding the course of the reaction. For a transformation to take place from left to right in the sense of the reaction equation, all the molecules A1, A2, ... must clearly collide at one point; otherwise no reaction is possible, since we shall not consider side-reactions. Such a collision need not of course bring about that transposition of the atoms of the single molecules which constitutes the above reaction. Much rather must it be of such a kind as is favourable to that loosening of the bonds that bind the atoms in the separate molecules, which must precede this transposition. Of a large number of such collisions, therefore, only a certain smaller number will involve a transposition from left to right in the sense of the equation. But this number will be the same under the same external conditions, and the greater the more numerous the collisions; in fact a direct ratio must exist between the two. Bearing in mind now, that the number of collisions must be proportional to each of the concentrations of the bodies A1, A2, ..., and therefore, on the whole, to the product of all these concentrations, we arrive at the conclusion that the velocity v of the transposition from left to right in the sense of the reaction equation is v = kc1c2 ..., in which c1, c2, ... represent the spatial concentrations, i.e. the number of gram-molecules of the substances A1, A2, ... present in one litre, and k is, at a given temperature, a constant which may be called the velocity-coefficient.

Exactly the same consideration applies to the molecules A′1, A′2.... Here the velocity of the change from right to left in the sense of the reaction-equation increases with the number of collisions of all these molecules at one point, and this is proportional to the product of all the concentrations. If k’ denotes the corresponding proportionality-factor, then the velocity v’ of the change from right to left in the sense of the reaction-equation is v′ = k′c′1c′2.... These spatial concentrations are often called the “active masses” of the reacting components. Hence the reaction-velocity in the sense of the reaction-equation from left to right, or the reverse, is proportional to the product of the “active-masses” of the left-hand or right-hand components respectively.

Neither v nor v′ can be separately investigated, and the measurements of the course of a reaction always furnish only the difference of these two quantities. The reaction-velocity Law of chemical statics. actually observed represents the difference of these two partial reaction-velocities, whilst the amount of change observed during any period of time is equal to the change in the one direction, minus the change in the opposite direction. It must not be assumed, however, that on the attainment of equilibrium all action has ceased, but rather that the velocity of change in one direction has become equal to that in the opposite direction, with the result that no further total change can be observed, i.e. the system has reached equilibrium, for which the relation v - v′ = 0 must therefore hold, or what is the same thing

kc1c2 ... = k’c’1c’2 ...,

this is the fundamental law of chemical statics.

The conception that the equilibrium is not to be attributed to absolute indifference between the reacting bodies, but that these continue to exert their mutual actions undiminished and the opposing changes now balance, is of fundamental significance in the interpretation of changes of matter in general. This is generally expressed in the form: the equilibrium in this and other analogous cases is not static but dynamic. This conception was a direct result of the kinetic-molecular considerations, and was applied with special success to the development of the kinetic theory of gases. Thus with Clausius, we conceive the equilibrium of water-vapour with water, not as if neither water vaporized nor vapour condensed, but rather as though the two processes went on unhindered in the equilibrium state, i.e. during contact of saturated vapour with water, in a given time, as many water molecules passed through the water surface in one direction as in the opposite direction. This view, as applied to chemical changes, was first advanced by A.W. Williamson (1851), and further developed by C.M. Guldberg and P. Waage and others.

From the previous considerations it follows that the reaction-velocity at every moment, i.e. the velocity with which the chemical process advances towards the Law of chemical kinetics. equilibrium state, is given by the equation

V = v - v′ = kc1c2 ... - k′c′1c′2 ...;

this states the fundamental law of chemical kinetics.

The equilibrium equation is simply a special case of this more general one, and results when the total velocity is written zero, just as in analytical mechanics the equilibrium conditions follow at once by specialization of the general equations of motion.

No difficulty presents itself in the generalization of the previous equations for the reaction which proceeds after the scheme

n1A1 + n2A2 + ... = n′1A′1 + n′2A′2 + ...,

where n1, n2, ..., n′1, n′2, ... denote the numbers of molecules of the separate substances which take part in the reaction, and are therefore whole, mostly small, numbers (generally one or two, seldom three or more). Here as before, v and v′ are to be regarded as proportional to the number of collisions at one point of all molecules necessary to the respective reaction, but now n1 molecules of A1, n2 molecules of A2, &c., must collide for the reaction to advance from left to right in the sense of the equation; and similarly n′1 molecules of A′1, n′2 molecules of A′2, &c., must collide for the reaction to proceed in the opposite direction. If we consider the path of a single, arbitrarily chosen molecule over a certain time, then the number of its collisions with other similar molecules will be proportional to the concentration C of that kind of molecule to which it belongs. The number of encounters between two molecules of the kind in question, during the same time, will be in general C times as many, i.e. the number of encounters of two of the same molecules is proportional to the square of the concentration C; and generally, the number of encounters of n molecules of one kind must be regarded as proportional to the nth power of C, i.e. Cn.

The number of collisions of n1 molecules of A1, n2 molecules of A2 ... is accordingly proportional to C1n1C2n2 ..., and the reaction-velocity corresponding to it is therefore

v = kC1n1C2n2 ...,

and similarly the opposed reaction-velocity is

v′ = k′C′1n′1C′2n′2 ...;

the resultant reaction-velocity, being the difference of these two partial velocities, is therefore

V = v - v′ = kC1n1C2n2 ... - k′C′1n′1C′2n′2 ...

This is the most general expression of the law of chemical mass-action, for the case of homogeneous systems.

Equating V to zero, we obtain the equation for the equilibrium state, viz.

C1n1C2n2 ... / C′1n′1C′2n′2 ... = k / k′ = K;

K is called the “equilibrium-constant.”

These formulae hold for gases and for dilute solutions, but assume the system to be homogeneous, i.e. to be either a homogeneous gas-mixture or a homogeneous dilute solution. Limitations and applications of the laws. The case in which other states of matter share in the equilibrium permits of simple treatment when the substances in question may be regarded as pure, and consequently as possessing definite vapour-pressures or solubilities at a given temperature. In this case the molecular species in question, which is, at the same time, present in excess and is hence usually, called a Bodenkörper, must possess a constant concentration in the gas-space or solution. But since the left-hand side of the last equation contains only variable quantities, it is simplest and most convenient to absorb these constant concentrations into the equilibrium-constant; whence we have the rule: leave the molecular species present as Bodenkörper out of account, when determining the concentration-product. Guldberg and Waage expressed this in the form “the active mass of a solid substance is constant.” The same is true of liquids when these participate in the pure state in the equilibrium, and possess therefore a definite vapour-pressure or solubility. When, finally, we are not dealing with a dilute solution but with any kind of mixture whatever, it is simplest to apply the law of mass-action to the gaseous mixture in equilibrium with this. The composition of the liquid mixture is then determinable when the vapour-pressures of the separate components are known. This, however, is not often the case; but in principle this consideration is important, since it involves the possibility of extending the law of chemical mass-action from ideal gas-mixtures and dilute solutions, for which it primarily holds, to any other system whatever.

The more recent development of theoretical chemistry, as well as the detailed study of many chemical processes which have found technical application, leads more and more convincingly to the recognition that in the law of chemical mass-action we have a law of as fundamental significance as the law of constant and multiple proportions. It is therefore not without interest to briefly touch upon the development of the doctrine of chemical affinity.

Historical Development of the Law of Mass-action.—The theory developed by Torbern Olof Bergman in 1775 must be regarded as the first attempt of importance to account for the mode of action of chemical forces. The essential principle of this may be stated as follows:—The magnitude of chemical affinity may be expressed by a definite number; if the affinity of the substance A is greater for the substance B than for the substance C, then the latter (C) will be completely expelled by B from its compound with A, in the sense of the equation A·C + B = A·B + C. This theory fails, however, to take account of the influence of the relative masses of the reacting substances, and had to be abandoned as soon as such an influence was noticed. An attempt to consider this factor was made by Claude Louis Berthollet (1801), who introduced the conception of chemical equilibrium. The views of this French chemist may be summed up in the following sentence:—Different substances have different affinities for each other, which only come into play on immediate contact. The condition of equilibrium depends not only upon the chemical affinity, but also essentially upon the relative masses of the reacting substances.

Essentially, Berthollet’s idea is to-day the guiding principle of the doctrine of affinity. This is especially true of our conceptions of many reactions which, in the sense of Bergman’s idea, proceed to completion, i.e. until the reacting substances are all used up; but only for this reason, viz. that one or more of the products of the reaction is removed from the reaction mixture (either by crystallization, evaporation or some other process), and hence the reverse reaction becomes impossible. Following Berthollet’s idea, two Norwegian investigators, C.M. Guldberg and Peter Waage, succeeded in formulating the influence of the reacting masses in a simple law—the law of chemical mass-action already defined. The results of their theoretical and experimental studies were published at Christiania in 1867 (Études sur les affinités chimiques); this work marks a new epoch in the history of chemistry. Even before this, formulae to describe the progress of certain chemical reactions, which must be regarded as applications of the law of mass-action, had been put forward by Ludwig Wilhelmy (1850), and by A.G. Vernon-Harcourt and William Esson (1856), but the service of Guldberg and Waage in having grasped the law in its full significance and logically applied it in all directions, remains of course undiminished. Their treatise remained quite unknown; and so it happened that John Hewitt Jellett (1873), J.H. van’t Hoff (1877), and others independently developed the same law. The thermodynamic basis of the law of mass-action is primarily due to Horstmann, J. Willard Gibbs and van’t Hoff.

Applications.—Let us consider, as an example of the application of the law of mass-action, the case of the dissociation of water-vapour, which takes place at high temperatures in the sense of the equation 2H2O = 2H2 + O2. Representing the concentrations of the corresponding molecular species by [H2], &c., the expression [H2]² [O2] / [H2O]² must be constant at any given temperature. This shows that the dissociation is set back by increasing the pressure; for if the concentrations of all three kinds of molecules be increased by strong compression, say to ten times the former amounts, then the numerator is increased one thousand, the denominator only one hundred times. Hence if the original equilibrium-constant is to hold, the dissociation must go back, and, what is more, by an exactly determinable amount. At 2000° C. water-vapour is only dissociated to the extent of a few per cent; therefore, even when only a small excess of oxygen or hydrogen be present, the numerator in the foregoing expression is much increased, and it is obvious that in order to restore the equilibrium state, the concentration of the other component, hydrogen or oxygen as the case may be, must diminish. In the case of slightly dissociated substances, therefore, even a relatively small excess of one component is sufficient to set back the dissociation substantially.

Chemical Kinetics.—It has been already mentioned that the law of chemical mass-action not only defines the conditions for chemical equilibrium, but contains at the same time the principles of chemical kinetics. The previous considerations show indeed that the actual progress of the reaction is determined by the difference of the reaction-velocities in the one and the other (opposed) direction, in the sense of the corresponding reaction-equation. Since the reaction-velocity is given by the amount of chemical change in a small interval of time, the law of chemical mass-action supplies a differential equation, which, when integrated, provides formulae which, as numerous experiments have shown, very happily summarize the course of the reaction. For the simplest case, in which a single species of molecule undergoes almost complete decomposition, so that the reaction-velocity in the reverse direction may be neglected, we have the simple equation

dx/dt = k(a-x)

and if x = 0 when t = 0 we have by integration

k = t-1 log {a / (a-x) }.

We will now apply these conclusions to the theory of the ignition of an explosive gas-mixture, and in particular to the combustion of “knallgas” (a mixture of hydrogen Theory of explosive combustion. and oxygen) to water-vapour. At ordinary temperatures knallgas undergoes practically no change, and it might be supposed that the two gases, oxygen and hydrogen, have no affinity for each other. This conclusion, however, is shown to be incorrect by the observation that it is only necessary to add some suitable catalyst such as platinum-black in order to immediately start the reaction. We must therefore conclude that even at ordinary temperatures strong chemical affinity is exerted between oxygen and hydrogen, but that at low temperatures this encounters great frictional resistances, or in other words that the reaction-velocity is very small. It is a matter of general experience that the resistances which the chemical forces have to overcome diminish with rising temperature, i.e. the reaction-velocity increases with temperature. Therefore, when we warm the knallgas, the number of collisions of oxygen and hydrogen molecules favourable to the formation of water becomes greater and greater, until at about 500° the gradual formation of water is observed, while at still higher temperatures the reaction-velocity becomes enormous. We are now in a position to understand what is the result of a strong local heating of the knallgas, as, for example, by an electric spark. The strongly heated parts of the knallgas combine to form water-vapour with great velocity and the evolution of large amounts of heat, whereby the adjacent parts are brought to a high temperature and into a state of rapid reaction, i.e. we observe an ignition of the whole mixture. If we suppose the knallgas to be at a very high temperature, then its combustion will be no longer complete owing to the dissociation of water-vapour, whilst at extremely high temperatures it would practically disappear. Hence it is clear that knallgas appears to be stable at low temperatures only because the reaction-velocity is very small, but that at very high temperatures it is really stable, since no chemical forces are then active, or, in other words, the chemical affinity is very small.

The determination of the question whether the failure of some reaction is due to an inappreciable reaction-velocity or to absence of chemical affinity, is of fundamental importance, and only in the first case can the reaction be hastened by catalysts.

Many chemical compounds behave like knallgas. Acetylene is stable at ordinary temperatures, inasmuch as it only decomposes slowly; but at the same time it is explosive, for the decomposition when once started is rapidly propagated, on account of the heat evolved by the splitting up of the gas into carbon and hydrogen. At very high temperatures, however, acetylene acquires real stability, since carbon and hydrogen then react to form acetylene.

Many researches have shown that the combustion of an inflammable gas-mixture which is started at a point, e.g. by an electric spark, may be propagated in two essentially Explosion-waves. different ways. The characteristic of the slower combustion consists in this, viz. that the high temperature of the previously ignited layer spreads by conduction, thereby bringing the adjacent layers to the ignition-temperature; the velocity of the propagation is therefore conditioned in the first place by the magnitude of the conductivity for heat, and more particularly, in the second place, by the velocity with which a moderately heated layer begins to react chemically, and so to rise gradually in temperature, i.e. essentially by the change of reaction-velocity with temperature. A second entirely independent mode of propagation of the combustion lies at the basis of the phenomenon that an explosive gas-mixture can be ignited by strong compression or—more correctly—by the rise of temperature thereby produced. The increase of the concentrations of the reacting substances consequent upon this increase of pressure raises the reaction-velocity in accordance with the law of chemical mass-action, and so enormously favours the rapid evolution of the heat of combustion.

It is therefore clear that such a powerful compression-wave can not only initiate the combustion, but also propagate it with extremely high velocity. Indeed a compression-wave of this kind passes through the gas-mixture, heated by the combustion to a very high temperature. It must, however, be propagated considerably faster than an ordinary compression-wave, for the result of ignition in the compressed (still unburnt) layer is the production of a very high pressure, which must in accordance with the principles of wave-motion increase the velocity of propagation. The absolute velocity of the explosion-wave would seem, in the light of these considerations, to be susceptible of accurate calculation. It is at least clear that it must be considerably higher than the velocity of sound in the mass of gas strongly heated by the explosion, and this is confirmed by actual measurements (see below) which show that the velocity of the explosion-wave is from one and a half times to double that of sound-waves at the combustion temperature.

We are now in a position to form the following picture of the processes which follow upon the ignition of a combustible gas-mixture contained in a long tube. First we have the condition of slow combustion; the heat is conveyed by conduction to the adjacent layers, and there follows a velocity of propagation of a few metres per second. But since the combustion is accompanied by a high increase of pressure, the adjacent, still unburnt layers are simultaneously compressed, whereby the reaction-velocity increases, and the ignition proceeds faster. This involves still greater compression of the next layers, and so if the mixture be capable of sufficiently rapid combustion, the velocity of propagation of the ignition must continually increase. As soon as the compression in the still unburnt layers becomes so great that spontaneous ignition results, the now much more pronounced compression-waves excited with simultaneous combustion must be propagated with very great velocity, i.e. we have spontaneous development of an “explosion-wave.” M.P.E. Berthelot, who discovered the presence of such explosion-waves, proved their velocity of propagation to be independent of the pressure, the cross-section of the tubes in which the explosive gas-mixture is contained, as well as of the material of which these are made, and concluded that this velocity is a constant, characteristic of the particular mixture. The determination of this velocity is naturally of the highest interest.

In the following table Berthelot’s results are given along with the later (1891) concordant ones of H.B. Dixon, the velocities of propagation of explosions being given in metres per second.

Reacting Mixture. Velocity of Wave in Metres per second.

Berthelot. Dixon.

Hydrogen and oxygen, H2+O 2810 2821

Hydrogen and nitrous oxide, H2+N2O 2284 2305

Methane and oxygen, CH4+4O 2287 2322

Ethylene and oxygen, C2H4+6O 2210 2364

Acetylene and oxygen, C2H2+5O 2482 2391

Cyanogen and oxygen, C2N2+4O 2195 2321

Hydrogen and chlorine, H2+Cl2 .. 1730

Hydrogen and chlorine, 2H2+Cl2 .. 1849

The maximum pressure of the explosion-wave possesses very high values; it appears that a compression of from 1 to 30-40 atmospheres is necessary to produce spontaneous ignition of mixtures of oxygen and hydrogen. But since the heat evolved in the path of the explosion causes a rise of temperature of 2000°-3000°, i.e. a rise of absolute temperature about four times that directly following upon the initial compression, we are here concerned with pressures amounting to considerably more than 100 atmospheres. Both the magnitude of this pressure and the circumstance that it so suddenly arises are peculiar to the very powerful forces which distinguish the explosion-wave from the slow combustion-wave.

Nascent State.—The great reactive power of freshly formed or nascent substances (status nascens)may be very simply referred to the principles of mass-action. As is well known, this phenomenon is specially striking in the case of hydrogen, which may therefore be taken as a typical example. The law of mass-action affirms the action of a substance to be the greater the higher its concentration, or, for a gas, the higher its partial-pressure. Now experience teaches that those metals which liberate hydrogen from acids are able to supply the latter under extremely high pressure, and we may therefore assume that the hydrogen which results, for example, from the action of zinc upon sulphuric acid is initially under very high pressures which are then afterwards relieved. Hence the hydrogen during liberation exhibits much more active powers of reduction than the ordinary gas.

A deeper insight into the relations prevailing here is offered from the atomistic point of view. From this we are bound to conclude that the hydrogen is in the first instance evolved in the form of free atoms, and since the velocity of the reaction H + H = H2 at ordinary temperatures, though doubtless very great, is not practically instantaneous, the freshly generated hydrogen will contain a remnant of free atoms, which are able to react both more actively and more rapidly. Similar considerations are of course applicable to other cases.

Ion-reactions.—The application of the law of chemical mass action is much simplified in the case in which the reaction-velocity is enormously great, when practically an instantaneous adjustment of the equilibrium results. Only in this case can the state of the system, which pertains after mixing the different components, be determined merely from knowledge of the equilibrium-constant. This case is realized in the reactions between gases at very high temperatures, which have, however, been little investigated, and especially by the reactions between electrolytes, the so-called ion-reactions. In this latter case, which has been thoroughly studied on account of its fundamental importance for inorganic qualitative and quantitative analysis, the degrees of dissociation of the various electrolytes (acids, bases and salts) are for the most part easily determined by the aid of the freezing-point apparatus, or of measurements of the electric conductivity; and from these data the equilibrium-constant K may be calculated. Moreover, it can be shown that the state of the system can be determined when the equilibrium constants of all the electrolytes which are present in the common solution are known. If this be coupled with the law that the solubility of solid substances, as with vapour-pressures, is independent of the presence of other electrolytes, it is sufficient to know the solubilities of the electrolytes in question, in order to be able to determine which substances must participate in the equilibrium in the solid state, i.e. we arrive at the theory of the formation and solution of precipitates.

As an illustration of the application of these principles, we shall deal with a problem of the doctrine of affinity, namely, that of the relative strengths of acids and bases. It Strength of acids and bases. was quite an early and often repeated observation that the various acids and bases take part with very varying intensity or avidity in those reactions in which their acid or basic nature comes into play. No success attended the early attempts at giving numerical expression to the strengths of acids and bases, i.e. of finding a numerical coefficient for each acid and base, which should be the quantitative expression of the degree of its participation in those specific reactions characteristic of acids and bases respectively. Julius Thomsen and W. Ostwald attacked the problem in a far-seeing and comprehensive manner, and arrived at indisputable proof that the property of acids and bases of exerting their effects according to definite numerical coefficients finds expression not only in salt-formation but also in a large number of other, and indeed very miscellaneous, reactions.

When Ostwald compared the order of the strengths of acids deduced from their competition for the same base, as determined by Thomsen’s thermo-chemical or his own volumetric method, with that order in which the acids arrange themselves according to their capacity to bring calcium oxalate into solution, or to convert acetamide into ammonium acetate, or to split up methyl acetate into methyl alcohol and acetic acid catalytically, or to invert cane-sugar, or to accelerate the mutual action of hydriodic on bromic acid, he found that in all these well-investigated and very miscellaneous cases the same succession of acids in the order of their strengths is obtained, whichever one of the above chemical processes be chosen as measure of these strengths. It is to be noticed that all these chemical changes cited took place in dilute aqueous solution, consequently the above order of acids refers only to the power to react under these circumstances. The order of acids proved to be fairly independent of temperature. While therefore the above investigations afforded a definite qualitative solution of the order of acids according to strengths, the determination of the quantitative relations offered great difficulties, and the numerical coefficients, determined from the separate reactions, often displayed great variations, though occasionally also surprising agreement. Especially great were the variations of the coefficients with the concentration, and in those cases in which the concentration of the acid changed considerably during the reaction, the calculation was naturally quite uncertain. Similar relations were found in the investigation of bases, the scope of which, however, was much more limited.

These apparently rather complicated relations were now cleared up at one stroke, by the application of the law of chemical mass-action on the lines indicated by S. Arrhenius in 1887, when he put forward the theory of electrolytic dissociation to explain that peculiar behaviour of substances in aqueous solution first recognized by van’t Hoff in 1885. The formulae which must be made use of here in the calculation of the equilibrium-relations follow naturally by simple application of the law of mass-action to the corresponding ion-concentrations.

The peculiarities which the behaviour of acids and bases presents, and, according to the theory of Arrhenius, must present—peculiarities which found expression in the very early distinction between neutral solutions on the one hand, and acid or basic ones on the other, as well as in the belief in a polar antithesis between the two last—must now, in the light of the theory of electrolytic dissociation, be conceived as follows:—

The reactions characteristic of acids in aqueous solution, which are common to and can only be brought about by acids, find their explanation in the fact that this class of bodies gives rise on dissociation to a common molecular species, namely, the positively charged hydrogen-ion (H+). The specific chemical actions peculiar to acids are therefore to be attributed to the hydrogen-ion just as the actions common to all chlorides are to be regarded as those of the free chlorine-ions. In like manner, the reactions characteristic of bases in solution are to be attributed to the negatively charged hydroxyl-ions (OH-), which result from the dissociation of this class of bodies.

A solution has an acid reaction when it contains an excess of hydrogen-ions, and a basic reaction when it contains an excess of hydroxyl-ions. If an acid and an alkaline solution be brought together mutual neutralization must result, since the positive H-ions and the negative OH-ions cannot exist together in view of the extremely weak conductivity of pure water and its consequent slight electrolytic dissociation, and therefore they must at once combine to form electrically neutral molecules, in the sense of the equation

H+ + OH- = H2O.

In this lies the simple explanation of the “polar” difference between acid and basic solutions. This rests essentially upon the fact that the ion peculiar to acids and the ion peculiar to bases form the two constituents of water, i.e. of that solvent in which we usually study the course of the reaction. The idea of the “strength” of an acid or base at once arises. If we compare equivalent solutions of various acids, the intensity of those actions characteristic of them will be the greater the more free hydrogen-ions they contain; this is an immediate consequence of the law of chemical mass-action. The degree of electrolytic dissociation determines, therefore, the strength of acids, and a similar consideration leads to the same result for bases.

Now the degree of electrolytic dissociation changes with concentration in a regular manner, which is given by the law of mass-action. For if C denote the concentration of the electrolyte and a its degree of dissociation, the above law states that

C²a²/C(1-a) = Ca²/(1-a) = K.

At very great dilutions the dissociation is complete, and equivalent solutions of the most various acids then contain the same number of hydrogen-ions, or, in other words, are equally strong; and the same is true of the hydroxyl-ions of bases. The dissociation also decreases with increasing concentration, but at different rates for different substances, and the relative “strengths” of acids and bases must hence change with concentration, as was indeed found experimentally. The dissociation-constant K is the measure of the variation of the degree of dissociation with concentration, and must therefore be regarded as the measure of the strengths of acids and bases. So that in this special case we are again brought to the result which was stated in general terms above, viz. that the dissociation-coefficient forms the measure of the reactivity of a dissolved electrolyte. Ostwald’s series of acids, based upon the investigation of the most various reactions, should therefore correspond with the order of their dissociation-constants, and further with the order of their freezing-point depressions in equivalent solutions, since the depression of the freezing-point increases with the degree of electrolytic dissociation. Experience confirms this conclusion completely. The degree of dissociation of an acid, at a given concentration, for which its molecular conductivity is Λ, is shown by the theory of electrolytic dissociation to be a = Λ / Λ∞; Λ∞, the molecular conductivity at very great dilution in accordance with the law of Kohlrausch, is u + v, where u and v are the ionic-mobilities (see Conduction, Electric). Since u, the ionic-mobility of the hydrogen ion, is generally more than ten times as great as v, the ionic-mobility of the negative acid-radical, Λ∞ has approximately the same value (generally within less than 10%) for the different acids, and the molecular-conductivity of the acids in equivalent concentration is at least approximately proportional to the degree of electrolytic dissociation, i.e. to the strength.

In general, therefore, the order of conductivities is identical with that in which the acids exert their specific powers. This remarkable parallelism, first perceived by Arrhenius and Ostwald in 1885, was the happy development which led to the discovery of electrolytic dissociation (see Conduction, Electric; and Solution).

Catalysis.—We have already mentioned the fact, early known to chemists, that many reactions proceed with a marked increase of velocity in presence of many foreign substances. With Berzelius we call this phenomenon “catalysis,” by which we understand that general acceleration of reactions which also progress when left to themselves, in the presence of certain bodies which do not change in amount (or only slightly) during the course of the reaction. Acids and bases appear to act catalytically upon all reactions involving consumption or liberation of water, and indeed that action is proportional to the concentration of the hydrogen or hydroxyl-ions. Further, the decomposition of hydrogen peroxide is “catalysed” by iodine-ions, the condensation of two molecules of benzaldehyde to benzoin by cyanogen-ions. One of the earliest known and technically most important instances of catalysis is that of the oxidation of sulphur dioxide to sulphuric acid by oxygen in the presence of oxides of nitrogen. Other well-known and remarkable examples are the catalysis of the combustion of hydrogen and of sulphur dioxide in oxygen by finely-divided platinum. We may also mention the interesting work of Dixon and Baker, which led to the discovery that a large number of gas-reactions, e.g. the combustion of carbon monoxide, the dissociation of sal-ammoniac vapour, and the action of sulphuretted hydrogen upon the salts of heavy metals, cease when water-vapour is absent, or at least proceed with greatly diminished velocity.

“Negative catalysis,” i.e. the retardation of a reaction by addition of some substance, which is occasionally observed, appears to depend upon the destruction of a “positive catalyte” by the body added.

A catalyte can have no influence, however, upon the affinity of a process, since that would be contrary to the second law of thermodynamics, according to which affinity of an isothermal process, which is measured by the maximum work, only depends upon the initial and final states. The effect of a catalyte is therefore limited to the resistances opposing the progress of a reaction, and does not influence its driving-force or affinity. Since the catalyte takes no part in the reaction its presence has no effect on the equilibrium-constant. This, in accordance with the law of mass-action, is the ratio of the separate reaction-velocities in the two contrary directions. A catalyte must therefore always accelerate the reverse-reaction. If the velocity of formation of a body be increased by addition of some substance then its velocity of decomposition must likewise increase. We have an example of this in the well-known fact that the formation, and no less the saponification, of esters, proceeds with increased velocity in the presence of acids, while the observation that in absence of water-vapour neither gaseous ammonium chloride dissociates nor dry ammonia combines with hydrogen chloride becomes clear on the same grounds.

A general theory of catalytic phenomena does not at present exist. The formation of intermediate products by the action of the reacting substance upon the catalyte has often been thought to be the cause of these. These intervening products, whose existence in many cases has been proved, then split up into the catalyte and the reaction-product. Thus chemists have sought to ascribe the influence of oxides of nitrogen on the formation of sulphuric acid to the initial formation of nitrosyl-sulphuric acid, SO2(OH)(NO2), from the mixture of sulphur dioxide, oxides of nitrogen and air, which then reacted with water to form sulphuric and nitrous acids. When the velocity of such intermediate reactions is greater than that of the total change, such an explanation may suffice, but a more certain proof of this theory of catalysis has only been reached in a few cases, though in many others it appears very plausible. Hence it is hardly possible to interpret all catalytic processes on these lines.

In regard to catalysis in heterogeneous systems, especially the hastening of gas-reactions by platinum, it is very probable that it is closely connected with the solution or absorption of the gases on the part of the metal. From the experiments of G. Bredig it seems that colloidal solutions of a metal act like the metal itself. The action of a colloidal-platinum solution on the decomposition of hydrogen peroxide is still sensible even at a dilution of 1/70,000,000 grm.-mol. per litre; indeed the activity of this colloidal-platinum solution calls to mind in many ways that of organic ferments, hence Bredig has called it an “inorganic ferment.” This analogy is especially striking in the change of their activity with time and temperature, and in the possibility, by means of bodies like sulphuretted hydrogen, hydrocyanic acid, &c., which act as strong poisons upon the latter, of “poisoning” the former also, i.e. of rendering it inactive. In the case of the catalytic action of water-vapour upon many processes of combustion already mentioned, a part of the effect is probably due to the circumstance, disclosed by numerous experiments, that the union of hydrogen and oxygen proceeds, between certain temperature limits at least, after the equation H2 + O2 = H2O2, that is, with the preliminary formation of hydrogen peroxide, which then breaks down into water and oxygen, and further, above all, to the fact that this substance results from oxygen and water at high temperatures with great velocity, though indeed only in small quantities.

The view now suggests itself, that, for example, in the combustion of carbon monoxide at moderately high temperatures, the reaction

(I.) 2CO + O2 = 2CO2

advances with imperceptible speed, but that on the contrary the two stages

(II.) 2H2O + O2 = 2H2O2,

(III.) 2CO + 2H2O2 = 2CO2 + 2H2O,

which together result in (I.), proceed rapidly even at moderate temperatures.

Temperature and Reaction-Velocity.—There are few natural constants which undergo so marked a change with temperature as those of the velocities of chemical changes. As a rule a rise of temperature of 10° causes a twofold or threefold rise of reaction-velocity.

If the reaction-coefficient k, in the sense of the equation derived above, viz. k = t-1 log {a / (a-x)}, be determined for the inversion of cane-sugar by an acid of given concentration, the following values are obtained:—

Temperature =25°40°45°50°55°

k =9.773139268491;

here a rise of temperature of only 30° suffices to raise the speed of inversion fifty times.

We possess no adequate explanation of this remarkable temperature influence; but some account of it is given by the molecular theory, according to which the energy of that motion of substances in homogeneous gaseous or liquid systems which constitutes heat increases with the temperature, and hence also the frequency of collision of the reacting substances. When we reflect that the velocity of motion of the molecules of gases, and in all probability those of liquids also, are proportional to the square root of the absolute temperature, and therefore rise by only 1/6% per degree at room-temperature, and that we must assume the number of collisions proportional to the velocity of the molecules, we cannot regard the actually observed increase of reaction-velocity, which often amounts to 10 or 12% per degree, as exclusively due to the quickening of the molecular motion by heat. It is more probable that the increase of the kinetic energy of the atomic motions within the molecule itself is of significance here, as the rise of the specific heat of gases with temperature seems to show. The change of the reaction-coefficient k with temperature may be represented by the empirical equation log k = -AT-1 + B + CT, where A, B, C are positive constants. For low temperatures the influence of the last term is as a rule negligible, whilst for high temperatures the first term on the right side plays a vanishingly small part.

Definition of Chemical Affinity.—We have still to discuss the question of what is to be regarded as the measure of chemical affinity. Since we are not in a position to measure directly the intensity of chemical forces, the idea suggests itself to determine the strength of chemical affinity from the amount of the work which the corresponding reaction is able to do. To a certain extent the evolution of heat accompanying the reaction is a measure of this work, and attempts have been made to measure chemical affinities thermo-chemically, though it may be easily shown that this definition was not well chosen. For when, as is clearly most convenient, affinity is so defined that it determines under all circumstances the direction of chemical change, the above definition fails in so far as chemical processes often take place with absorption of heat, that is, contrary to affinities so defined. But even in those cases in which the course of the reaction at first proceeds in the sense of the evolution of heat, it is often observed that the reaction advances not to completion but to a certain equilibrium, or, in other words, stops before the evolution of heat is complete.

A definition free from this objection is supplied by the second law of thermodynamics, in accordance with which all processes must take place in so far as they are able to do external work. When therefore we identify chemical affinity with the maximum work which can be gained from the process in question, we reach such a definition that the direction of the process is under all conditions determined by the affinity. Further, this definition has proved serviceable in so far as the maximum work in many cases may be experimentally measured, and moreover it stands in a simple relation to the equilibrium constant K. Thermodynamics teaches that the maximum work A may be expressed as A = RT log K, when R denotes the gas-constant, T the absolute temperature. In this it is further assumed that both the molecular species produced as well as those that disappear are present in unit concentration. The simplest experimental method of directly determining chemical affinity consists in the measurement of electromotive force. The latter at once gives us the work which can be gained when the corresponding galvanic element supplies the electricity, and, since the chemical exchange of one gram-equivalent from Faraday’s law requires 96,540 coulombs, we obtain from the product of this number and the electromotive force the work per gram-equivalent in watt-seconds, and this quantity when multiplied by O.23872 is obtained in terms of the usual unit, the gram-calorie. Experience teaches that, especially when we have to deal with strong affinities, the affinity so determined is for the most part almost the same as the heat-evolution, whilst in the case in which only solid or liquid substances in the pure state take part in the reaction at low temperatures, heat-evolution and affinity appear to possess a practically identical value.

Hence it seems possible to calculate equilibria for low temperatures from heats of reaction, by the aid of the two equations

A = Q, A = RT log K;

and since the change of A with temperature, as required by the principles of thermodynamics, follows from the specific heats of the reacting substances, it seems further possible to calculate chemical equilibria from heats of reaction and specific heats. The circumstance that chemical affinity and heat-evolution so nearly coincide at low temperatures may be derived from the hypothesis that chemical processes are the result of forces of attraction between the atoms of the different elements. If we may disregard the kinetic energy of the atoms, and this is legitimate for low temperatures, it follows that both heat-evolution and chemical affinity are merely equal to the decrease of the potential energy of the above-mentioned forces, and it is at once clear that the evolution of heat during a reaction between only pure solid or pure liquid substances possesses special importance.

More complicated is the case in which gases or dissolved substances take part. This is simplified if we first consider the mixing of two mutually chemically indifferent gases. Thermodynamics teaches that external work may be gained by the mere mixing of two such gases (see Diffusion), and these amounts of work, which assume very considerable proportions at high temperatures, naturally affect the value of the maximum work and so also of the affinity, in that they always come into play when gases or solutions react. While therefore we regard as chemical affinity in the strictest sense the decrease of potential energy of the forces acting between the atoms, it is clear that the quantities here involved exhibit the simplest relations under the experimental conditions just given, for when only substances in a pure state take part in a reaction, all mixing of different kinds of molecules is excluded; moreover, the circumstance that the respective substances are considered at very low temperatures reduces the quantities of energy absorbed as kinetic energy by their molecules to the smallest possible amount.

Chemical Resistance.—When we know the chemical affinity of a reaction, we are in a position to decide in which direction the process must advance, but, unless we know the reaction-velocity also, we can in many cases say nothing as to whether or not the reaction in question will progress with a practically inappreciable velocity so that apparent chemical indifference is the result. This question may be stated in the light of the law of mass-action briefly as follows:—From a knowledge of the chemical affinity we can calculate the equilibrium, i.e. the numerical value of the constant K = k / k′; but to be completely informed of the process we must know not only the ratio of the two velocity-constants k and k′, but also the separate absolute values of the same.

In many respects the following view is more comprehensive, though naturally in harmony with the one just expressed. Since the chemical equilibrium is periodically attained, it follows that, as in the case of the motion of a body or of the diffusion of a dissolved substance, it must be opposed by very great friction. In all these cases the velocity of the process at every instant is directly proportional to the driving-force and inversely proportional to the frictional resistance. We hence arrive at the result that an equation of the form

reaction-velocity = chemical force/chemical resistance

must also hold for chemical change; here we have an analogy with Ohm’s law. The “chemical force” at every instant may be calculated from the maximum work (affinity); as yet little is known about “chemical resistance,” but it is not improbable that it may be directly measured or theoretically deduced. The problem of the calculation of chemical reaction-velocity in absolute measure would then be solved; so far we possess indeed only a few general facts concerning the magnitude of chemical resistance. It is immeasurably small at ordinary temperatures for ion-reactions, and, on the other hand, fairly large for nearly all reactions in which carbon-bonds must be loosened (so-called “inertia of the carbon-bond”) and possesses very high values for most gas-reactions also. With rising temperature it always strongly diminishes; on the other hand, at very low temperatures its values are always enormous, and at the absolute zero of temperature may be infinitely great. Therefore at that temperature all reactions cease, since the denominator in the above expression assumes enormous values.

It is a very remarkable phenomenon that the chemical resistance is often small in the case of precisely those reactions in which the affinity is also small; to this circumstance is to be traced the fact that in many chemical changes the most stable condition is not at once reached, but is preceded by the formation of more or less unstable intermediate products. Thus the unstable ozone is very often first formed on the evolution of oxygen, whilst in the reaction between oxygen and hydrogen water is often not at once formed, but first the unstable hydrogen peroxide as an intermediate product.

Let us now consider the chemical process in the light of the equation

reaction-velocity = chemical force/chemical resistance.

Thermodynamics shows that at very low temperatures, i.e. in the immediate vicinity of the absolute zero, there is no equilibrium, but every chemical process advances to completion in the one or the other direction. The chemical forces therefore act in the one direction towards complete consumption of the reacting substance. But since the chemical resistance is now immensely great, they can produce practically no appreciable result.

At higher temperatures the reaction always proceeds, at least in homogeneous systems, to a certain equilibrium, and as the chemical resistance now has finite values this equilibrium will always finally be reached after a longer or shorter time. Finally, at very high temperatures the chemical resistance is in every case very small, and the equilibrium is almost instantaneously reached; at the same time, the affinity of the reaction, as in the case of the mutual affinity between oxygen and hydrogen, may very strongly diminish, and we have then chemical indifference again, not because, as at low temperatures, the denominator of the previous expression becomes very great, but because the numerator now assumes vanishingly small values.

(W. N.)

CHEMISTRY (formerly “chymistry”; Gr. χυμεία; for derivation see Alchemy), the natural science which has for its province the study of the composition of substances. In common with physics it includes the determination of properties or characters which serve to distinguish one substance from another, but while the physicist is concerned with properties possessed by all substances and with processes in which the molecules remain intact, the chemist is restricted to those processes in which the molecules undergo some change. For example, the physicist determines the density, elasticity, hardness, electrical and thermal conductivity, thermal expansion, &c.; the chemist, on the other hand, investigates changes in composition, such as may be effected by an electric current, by heat, or when two or more substances are mixed. A further differentiation of the provinces of chemistry and physics is shown by the classifications of matter. To the physicist matter is presented in three leading forms—solids, liquids and gases; and although further subdivisions have been rendered necessary with the growth of knowledge the same principle is retained, namely, a classification based on properties having no relation to composition. The fundamental chemical classification of matter, on the other hand, recognizes two groups of substances, namely, elements, which are substances not admitting of analysis into other substances, and compounds, which do admit of analysis into simpler substances and also of synthesis from simpler substances. Chemistry and physics, however, meet on common ground in a well-defined branch of science, named physical chemistry, which is primarily concerned with the correlation of physical properties and chemical composition, and, more generally, with the elucidation of natural phenomena on the molecular theory.

It may be convenient here to state how the whole subject of chemistry is treated in this edition of the Encyclopaedia Britannica. The present article includes the following sections:—

I. History.—This section is confined to tracing the general trend of the science from its infancy to the foundations of the modern theory. The history of the alchemical period is treated in more detail in the article Alchemy, and of the iatrochemical in the article Medicine. The evolution of the notion of elements is treated under Element; the molecular hypothesis of matter under Molecule; and the genesis of, and deductions from, the atomic theory of Dalton receive detailed analysis in the article Atom.

II. Principles.—This section treats of such subjects as nomenclature, formulae, chemical equations, chemical change and similar subjects. It is intended to provide an introduction, necessarily brief, to the terminology and machinery of the chemist.

III. Inorganic Chemistry.—Here is treated the history of descriptive inorganic chemistry; reference should be made to the articles on the separate elements for an account of their preparation, properties, &c.

IV. Organic Chemistry.—This section includes a brief history of the subject, and proceeds to treat of the principles underlying the structure and interrelations of organic compounds.

V. Analytical Chemistry.—This section treats of the qualitative detection and separation of the metals, and the commoner methods employed in quantitative analysis. The analysis of organic compounds is also noticed.

VI. Physical Chemistry.—This section is restricted to an account of the relations existing between physical properties and chemical composition. Other branches of this subject are treated in the articles Chemical Action; Energetics; Solution; Alloys; Thermochemistry.

I. History

Although chemical actions must have been observed by man in the most remote times, and also utilized in such processes as the extraction of metals from their ores and in the arts of tanning and dyeing, there is no evidence to show that, beyond an unordered accumulation of facts, the early developments of these industries were attended by any real knowledge of the nature of the processes involved. All observations were the result of accident or chance, or possibly in some cases of experimental trial, but there is no record of a theory or even a general classification of the phenomena involved, although there is no doubt that the ancients had a fair knowledge of the properties and uses of the commoner substances. The origin of chemistry is intimately bound up with the arts which we have indicated; in this respect it is essentially an experimental science. A unifying principle of chemical and physical changes was provided by metaphysical conceptions of the structure of matter. We find the notion of “elements,” or primary qualities, which confer upon all species of matter their distinctive qualities by appropriate combination, and also the doctrine that Greek philosophy. matter is composed of minute discrete particles, prevailing in the Greek schools. These “elements,” however, had not the significance of the elements of to-day; they connoted physical appearances or qualities rather than chemical relations; and the atomic theory of the ancients is a speculation based upon metaphysical considerations, having, in its origin, nothing in common with the modern molecular theory, which was based upon experimentally observed properties of gases (see Element; Molecule).

Although such hypotheses could contribute nothing directly to the development of a science which laid especial claim to experimental investigations, yet indirectly they stimulated inquiry into the nature of the “essence” with which the four “elements” were associated. This quinta essentia had been speculated upon by the Greeks, some regarding it as immaterial or aethereal, and others as material; and a school of philosophers termed alchemists arose who attempted the isolation of this essence. The existence of a fundamental principle, unalterable and indestructible, prevailing alike through physical and chemical changes, was generally accepted. Any change which a substance may chance to undergo was simply due to the discarding or taking up of some proportion of the primary “elements” or qualities: of these coverings “water,” “air,” “earth” and “fire” were regarded as clinging most tenaciously to the essence, while “cold,” “heat,” “moistness” and “dryness” were more easily cast aside or assumed. Several origins have been Alchemy. suggested for the word alchemy, and there seems to have been some doubt as to the exact nature and import of the alchemical doctrines. According to M.P.E. Berthelot, “alchemy rested partly on the industrial processes of the ancient Egyptians, partly on the speculative theories of the Greek philosophers, and partly on the mystical reveries of the Gnostics and Alexandrians.” The search for this essence subsequently resolved itself into the desire to effect the transmutation of metals, more especially the base metals, into silver and gold. It seems that this secondary principle became the dominant idea in alchemy, and in this sense the word is used in Byzantine literature of the 4th century; Suidas, writing in the 11th century, defines chemistry as the “preparation of silver and gold” (see Alchemy).

From the Alexandrians the science passed to the Arabs, who made discoveries and improved various methods of separating substances, and afterwards, from the 11th century, became seated in Europe, where the alchemical doctrines were assiduously studied until the 15th and 16th centuries. It is readily understood why men imbued with the authority of tradition should prosecute the search for a substance which would confer unlimited wealth upon the fortunate discoverer. Some alchemists honestly laboured to effect the transmutation and to discover the “philosopher’s stone,” and in many cases believed that they had achieved success, if we may rely upon writings assigned to them. The period, however, is one of literary forgeries; most of the MSS. are of uncertain date and authorship, and moreover are often so vague and mystical that they are of doubtful scientific value, beyond reflecting the tendencies of the age. The retaining of alchemists at various courts shows the high opinion which the doctrines had gained. It is really not extraordinary that Isaac Hollandus was able to indicate the method of the preparation of the “philosopher’s stone” from “adamic” or “virgin” earth, and its action when medicinally employed; that in the writings assigned to Roger Bacon, Raimon Lull, Basil Valentine and others are to be found the exact quantities of it to be used in transmutation; and that George Ripley, in the 15th century, had grounds for regarding its action as similar to that of a ferment.

In the view of some alchemists, the ultimate principles of matter were Aristotle’s four elements; the proximate constituents were a “sulphur” and a “mercury,” the father and mother of the metals; gold was supposed to have attained to the perfection of its nature by passing in succession through the forms of lead, brass and silver; gold and silver were held to contain very pure red sulphur and white quicksilver, whereas in the other metals these materials were coarser and of a different colour. From an analogy instituted between the healthy human being and gold, the most perfect of the metals, silver, mercury, copper, iron, lead and tin, were regarded in the light of lepers that required to be healed.

Notwithstanding the false idea which prompted the researches of the alchemists, many advances were made in descriptive chemistry, the metals and their salts receiving much Iatrochemistry. attention, and several of our important acids being discovered. Towards the 16th century the failure of the alchemists to achieve their cherished purpose, and the general increase of medical knowledge, caused attention to be given to the utilization of chemical preparations as medicines. As early as the 15th century the alchemist Basil Valentine had suggested this application, but the great exponent of this doctrine was Paracelsus, who set up a new definition: “The true use of chemistry is not to make gold but to prepare medicines.” This relation of chemistry to medicine prevailed until the 17th century, and what in the history of chemistry is termed the iatrochemical period (see Medicine) was mainly fruitful in increasing the knowledge of compounds; the contributions to chemical theory are of little value, the most important controversies ranging over the nature of the “elements,” which were generally akin to those of Aristotle, modified so as to be more in accord with current observations. At the same time, however, there were many who, opposed to the Paracelsian definition of chemistry, still laboured at the problem of the alchemists, while others gave much attention to the chemical industries. Metallurgical operations, such as smelting, roasting and refining, were scientifically investigated, and in some degree explained, by Georg Agricola and Carlo Biringuiccio; ceramics was studied by Bernard Palissy, who is also to be remembered as an early worker in agricultural chemistry, having made experiments on the effect of manures on soils and crops; while general technical chemistry was enriched by Johann Rudolf Glauber.1

The second half of the 17th century witnessed remarkable transitions and developments in all branches of natural science, and the facts accumulated by preceding generations Boyle. during their generally unordered researches were replaced by a co-ordination of experiment and deduction. From the mazy and incoherent alchemical and iatrochemical doctrines, the former based on false conceptions of matter, the latter on erroneous views of life processes and physiology, a new science arose—the study of the composition of substances. The formulation of this definition of chemistry was due to Robert Boyle. In his Sceptical Chemist (1662) he freely criticized the prevailing scientific views and methods, with the object of showing that true knowledge could only be gained by the logical application of the principles of experiment and deduction. Boyle’s masterly exposition of this method is his most important contribution to scientific progress. At the same time he clarified the conception of elements and compounds, rejecting the older notions, the four elements of the “vulgar Peripateticks” and the three principles of the “vulgar Stagyrists,” and defining an element as a substance incapable of decomposition, and a compound as composed of two or more elements. He explained chemical combination on the hypotheses that matter consisted of minute corpuscles, that by the coalescence of corpuscles of different substances distinctly new corpuscles of a compound were formed, and that each corpuscle had a certain affinity for other corpuscles.

Although Boyle practised the methods which he expounded, he was unable to gain general acceptance of his doctrine of elements; and, strangely enough, the theory which next dominated chemical thought was an alchemical Phlogistic theory. invention, and lacked the lucidity and perspicuity of Boyle’s views. This theory, named the phlogistic theory, was primarily based upon certain experiments on combustion and calcination, and in effect reduced the number of the alchemical principles, while setting up a new one, a principle of combustibility, named phlogiston (from φλοιστός, burnt). Much discussion had centred about fire or the “igneous principle.” On the one hand, it had been held that when a substance was burned or calcined, it combined with an “air”; on the other hand, the operation was supposed to be attended by the destruction or loss of the igneous principle. Georg Ernst Stahl, following in some measure the views held by Johann Joachim Becher, as, for instance, that all combustibles contain a “sulphur” (which notion is itself of older date than Becher’s terra pinguis), regarded all substances as capable of resolution into two components, the inflammable principle phlogiston, and another element—“water,” “acid” or “earth.” The violence or completeness of combustion was proportional to the amount of phlogiston present. Combustion meant the liberation of phlogiston. Metals on calcination gave calces from which the metals could be recovered by adding phlogiston, and experiment showed that this could generally be effected by the action of coal or carbon, which was therefore regarded as practically pure phlogiston; the other constituent being regarded as an acid. At the hands of Stahl and his school, the phlogistic theory, by exhibiting a fundamental similarity between all processes of combustion and by its remarkable flexibility, came to be a general theory of chemical action. The objections of the antiphlogistonists, such as the fact that calces weigh more than the original metals instead of less as the theory suggests, were answered by postulating that phlogiston was a principle of levity, or even completely ignored as an accident, the change of qualities being regarded as the only matter of importance. It is remarkable that this theory should have gained the esteem of the notable chemists who flourished in the 18th century. Henry Cavendish, a careful and accurate experimenter, was a phlogistonist, as were J. Black, K. W. Scheele, A. S. Marggraf, J. Priestley and many others who might be mentioned.