COMBUSTION (Lat. combustio, a burning, from combu•ere, to burn, probably for *cow urcre, but with b inserted on the analogy of ambustus, burned; less plausibly for *co-arab-tit-e•e, to burn, from corn-, together + amb, ambi, Gk. 40, amphi, around + vrcre, to burn), or BURNING. The process by whieh bodies combine with oxy gen, and are thus seemingly destroyed. The term is, in ordinary parlance, restricted to eases in which the process of combination takes place rapidly and is accompanied by heat and light, as the combustion of wood in a fireplace, the combustion of a candle, etc. In its more scientific usage, however, the term may designate any possible case of direct combination with oxygen, whether rapid or slow, whether accom panied by light or not. By analogy, the term is also sometimes applied to the rapid union of substances with 'supporters of combustion' other than oxygen, such as chlorine gas, in which a candle may burn almost as well as in air.
The light and heat of combustion are utilized for purposes of every-day life, the combustible material employed, i.e. the illuminant or the fuel, being usually sonic product containing car bon. Thus ordinary illuminating gas contains a number of gaseous chemical compounds of car bon. Coal and wood are mixtures of carbon com pounds, the former containing even a. consider able amount of free carbon, Hydrogen, too, is one of the chemical constituents of most fuels and illuminants. When combustion takes place. the carbon and hydrogen, combining with oxy gen, give, respectively, carbonic-acid gas (carbon dioxide) and water vapor. These are. therefore, the chief products of ordinary combustion.
The heat produced by the combustion of a given amount of material is ,independent of the rate at which the combustion takes place, but depends entirely upon the composition and chemical nature of the material burned. Every combustible chemical compound has, therefore, its own definite heat of combustion; that is to say, the number of heat units (say, gram-calo ries) produced by the combustion of one chemi cal equivalent (gram-molecule) of a compound, depends upon nothing but the nature of the com pound. The following are the heats of combus tion of a few well-known compounds of carbon: ordinary alcohol, 340.000 gram-ealories; acetic acid (the sour principle of vinegar), 210,000; ethyl-acetic ester, 544.000; cane-sugar. 1,355.000; cellulose, 630,000; urea, 152,000. The combus tion of a chemical compound may be conceived as taking place in two consecutive steps: first, the compound is decomposed, i.e. every one of its molecules is broken up into its constituent atoms—a process usually involving not evolu tion, but absorption of heat ; secondly, every single atom capable of so doing combines with oxygen (0) an atom of carbon (C), thus yield ing a molecule of carbonic acid and two atoms of hydrogen (H) yielding a molecule of water (H..0). This second step of the process is accompanied by the evolution of a quantity of heat depending upon the number of carbon and hydrogen atoms in a molecule of the com bustible compound. But, owing to the absorption
taking place during the first part of the process, a portion only of the heat produced during the second part actually appears in the form of sensible heat, and it is that portion which is called the heat of combustion. Au exact knowl edge of the heats of combustion of various sub stances is of great importance for theoretical as well as for immediate practical purposes. lts practical importance in comparing. for instance, different kinds of fuel, is self-evident and re quires no explanation. Its theoretical impor tance is mainly in the fact that with the aid of it the exact amount of heat evolved or absorbed during various chemical transforma tions can be readily calculated. According to the first law of thermodynamics, the amount of heat evolved or absorbed during any transfortna tion whatever, is independent of the manner in which the transformation takes place. For ex ample, the amount of heat produced by hurting one equivalent weight of ordinary alcohol and one equivalent weight of acetic acid, is the same whether we burn theta directly or first cause them to combine into ethyl-acetic ester, and then burn the latter. In the second case, the heat absorbed during the formation of the ester must, of course, be combined with that evolved during its combustion. But this suggests a simple way of obtaining the heat of formation of the ester by merely out two combustions. The total heat of combustion of free alcohol and acetic acid is 340,000 ± '210,000 = 550,000 gram calories (see above) ; that of ethyl-acetic ester is 554,000 gram-calories. The excess of 4000 gram calories must therefore represent the amount absorbed during the combination of alcohol and the acid into ethyl-acetic ester. In a similar manner, the heat of any chemical reaction may be determined, if the heats of combustion of the reacting substances and the heats of combustion of the products of the reaction are known. In ninny eases this is the only certain way of determining with sortie precision the heat of reactions, as direct measurement during a reaction would often involve very great experimental difficulties, while the direct measurement of the heat of com bustion is a comparatively simple matter. The heat of combustion is usually determined by chemists in the following manner: A known amount of the combustible substance is inclosed in an air-tight steel vessel filled with compressed oxygen and lined on the inside with platinum; the vessel is immersed in a calorimeter (see CALORIMETRY , and the substance is ignited with the aid of an iron wire heated by means of an electric current ; the observer measures the rise of temperature in the calorimeter, and from this calculates the amount of heat produced. The importance of knowing the heat of chemical reactions is dis cussed in the article TUERMOCHEMISTRY (q.v.).