About the Work
This volume presents Wilhelm Ostwald’s systematic attempt to lay bare the theoretical foundations upon which the entire practice of analytical chemistry rests. Written as an elementary treatment, the book nevertheless aspires to something far more ambitious than a mere collection of laboratory procedures. It seeks to transform analytical chemistry from a craft governed by accumulated empirical rules into a science whose operations can be understood, predicted, and justified through a relatively small number of underlying physical and chemical principles. The work is organised into two broad divisions, the first devoted to theory and the second to applications, with the theoretical portion occupying the greater share of the author’s attention and intellectual energy.
“It is only within the last few years, i.e. since the development of the general theory of chemical reactions and states of equilibrium, that it has become possible to elaborate a theory of analytical reactions.”
The book addresses itself to students and practitioners of chemistry who have already encountered the basic operations of analysis but who may never have paused to ask why those operations work as they do. Ostwald’s guiding conviction is that analytical chemistry, far from being a merely practical discipline, is in fact a particularly instructive branch of physical chemistry—one that illustrates, in concrete and tangible form, the laws of chemical equilibrium, the theory of solutions, and the behaviour of ions. Throughout the theoretical chapters, he draws upon the most recent developments in physical chemistry, particularly the theory of electrolytic dissociation associated with Arrhenius and the osmotic theory of solutions associated with van’t Hoff, to explain phenomena that had previously been understood only imperfectly or not at all.
The book is notable for its insistence that even the most routine analytical operations—washing a precipitate, shaking a solution with an immiscible solvent, titrating an acid against a base—are governed by precise quantitative laws that can be expressed mathematically and understood theoretically. Ostwald does not shrink from presenting these mathematical relationships, though he is careful to keep the treatment elementary and to omit derivations that would require advanced calculus. The result is a work that stands midway between a laboratory manual and a treatise on physical chemistry, drawing on both traditions to create something that was, at the time of its writing, genuinely new.
Context
The book emerges from a period of extraordinary ferment in the physical sciences, when chemistry was being transformed by a series of theoretical advances that promised to unify its scattered empirical observations. The closing decades of the nineteenth century had witnessed the development of the theory of electrolytic dissociation, the formulation of the law of mass-action in its modern form, and the recognition that dilute solutions obey laws strikingly analogous to those governing gases. These developments, associated with such figures as van’t Hoff, Arrhenius, and Ostwald himself, collectively constituted what might be called the physico-chemical revolution—a movement that sought to explain chemical phenomena in terms of the behaviour of molecules and ions governed by quantitative laws.
Ostwald was one of the principal architects of this revolution, and the present work represents his attempt to bring its insights to bear on the traditional discipline of analytical chemistry. The book reflects a moment when the older descriptive approach to chemistry—which classified substances and reactions according to their observable characteristics—was giving way to a more fundamental approach grounded in the behaviour of ions in solution and the equilibria they establish. The author writes with the confidence of someone who believes that the theoretical tools now available are adequate to explain, in principle, every operation the analyst performs.
The work also reflects the pedagogical concerns of its author. Ostwald was deeply interested in the teaching of chemistry and believed that students should be led to understand the reasons behind laboratory procedures rather than merely trained to execute them mechanically. The book repeatedly emphasises the importance of theory for intelligent practice, arguing that the analyst who understands the principles governing precipitation, washing, and titration will perform these operations more skillfully and interpret their results more correctly than one who merely follows rules of thumb. This conviction gives the book its distinctive character as a work that is simultaneously theoretical and practical, addressed to the student’s understanding as much as to the student’s hands.
Structure of the Book
The work is organised into two principal divisions, the first devoted to theory and the second to applications. The theoretical portion opens with a consideration of how different substances may be recognised, proceeds through the various methods by which substances may be separated from one another, and concludes with an extended treatment of the quantitative determination of substances. The second division applies the theoretical principles developed earlier to the systematic treatment of particular groups of elements and compounds.
The opening chapter concerns the recognition of different substances through their physical properties. Ostwald observes that any property of a substance can in principle serve to identify it, but that practical considerations favour properties whose differences can be established quickly and easily. The state of aggregation and colour receive particular attention as properties that can usually be assessed at a glance. The discussion of colour leads to an examination of the phenomena of absorption and surface reflection, and to the use of coloured flames and the spectroscope as means of recognising substances that volatilise when heated.
The second chapter takes up the far more difficult problem of separating substances from mixtures. Ostwald establishes the fundamental principle that separation always requires the creation of a surface of separation between the substances to be divided, and that such surfaces occur most readily between different states of aggregation. He then examines the separation of solids from solids, liquids from liquids, and gases from gases, treating such methods as elutriation, decantation, and the use of porous partitions. The chapter concludes with the important observation that when physical methods prove inadequate, the analyst must resort to chemical transformations that convert the substances to be separated into forms amenable to mechanical division.
The third chapter examines physical methods of separation in greater detail, focusing on distillation and solution. Ostwald develops the theory of distillation with considerable care, treating the vapour pressures of pure substances, the behaviour of mixtures of miscible and immiscible liquids, and the conditions under which constant-boiling mixtures arise. The treatment of solution covers the laws of solubility, the distinction between crystalline and amorphous forms, and the practical procedures for extracting soluble constituents from solids.
The fourth chapter, devoted to chemical separation, constitutes the theoretical heart of the book. Here Ostwald introduces the theory of electrolytic dissociation and shows how it illuminates virtually every operation the analyst performs. The chapter develops the law of mass-action in its general form, applies it to the behaviour of electrolytes, and derives from it the principles governing precipitation, the redissolving of precipitates, and the formation of complex compounds. The treatment of the solubility-product, the conditions for complete precipitation, and the factors that determine whether a given precipitate will dissolve in a given reagent are all presented as direct consequences of the fundamental laws of chemical equilibrium.
The fifth chapter treats the quantitative determination of substances, beginning with the general conditions that make a substance suitable for weighing and proceeding to the various methods by which quantities may be measured without complete separation from other constituents. The discussion covers the measurement of gases and liquids by volume, the analysis of binary mixtures through physical properties, and the methods of titration in which the quantity of a substance is inferred from the amount of a reagent required to react with it completely.
The second division of the book applies these theoretical principles to the systematic treatment of particular groups of elements. The chapters devoted to the metals of the alkalies, the iron group, the copper group, and the tin group, together with the treatment of the non-metals, demonstrate how the general principles developed earlier govern the behaviour of specific substances and guide the analyst in devising procedures for their separation and determination. The book concludes with a chapter on the calculation of analyses, which treats the numerical operations by which the results of analytical procedures are converted into statements about the composition of the original material.
Inside the Book
Recognition of Substances via Physical Properties
The text establishes a foundational principle: while any property of a substance’s condition can, in theory, be used to identify it, distinguishing between different kinds of matter always relies on observing quantitative differences in that property. The practical difficulty of measuring these differences varies greatly, and those properties whose variations can be established easily and quickly are naturally preferred for analytical work. Among these, the state of aggregation—whether a substance is solid, liquid, or gas—deserves first place, with colour taking second, as both can usually be assessed at a single glance.
Transition Stages Between States of Aggregation
While transition stages exist between the three primary states of matter, they are of little immediate concern. A gradual change from gas to liquid, for instance, only occurs under pressures higher than the substance’s critical pressure, which typically ranges between 25 and 100 atmospheres, a condition not encountered in ordinary analytical operations. The changes between solid and liquid are more relevant. These can be either sudden, as with the melting of ice, or gradual, as with the fusion of glass. Gradual fusion is characteristic of amorphous substances, whereas sudden melting is peculiar to crystalline ones. By simple observation, one can distinguish grades of fluidity—mobile, fluid, viscous, sticky, and solid—but without other aids, the eye cannot reliably discern more than four or five stages. For solids, deciding between amorphous and crystalline forms is often easy with large fragments: amorphous bodies show a conchoidal fracture with uneven surfaces, while crystalline bodies fracture along plane surfaces. Powders, however, require a pocket lens or microscope for such determination.
Colour and Light
Colour is a widely applicable characteristic for recognition. Small differences in the wave-lengths of reflected light are perceived by the eye as colour differences, effectively converting quantitative differences into a series of qualitatively distinct but continuous steps, allowing ten, twenty, or more grades to be distinguished. The light radiating from the surface of a coloured body is a mixture of two kinds: light coloured through absorption from the interior, and light thrown back by surface reflection, which is usually white. The relationship between these two depends on the degree of subdivision of the substance and the difference between its refraction-coefficient and that of the surrounding medium. Depending on the amount of white surface light, a substance’s colour can vary from white to a very dark tint approaching black, making it essential to specify the conditions of observation—whether the substance is compact, powdery, or floating in a liquid. Most analytical cases concern powders deposited from water, known as precipitates. Another important phenomenon is the coloured flame, where certain substances heated in a nearly non-luminous flame, such as a Bunsen burner, volatilise and emit light of a limited number of definite rays and colours. While the naked eye can observe this simply, the spectroscope separates the light into its constituent parts, providing an infinitely higher degree of discrimination and constituting one of the most thorough and certain aids for recognizing substances that yield coloured flames. Many other properties exist, but as they are slower and more difficult to assess, they are not considered practically here.
The General Problem of Separation
Recognizing a single substance presupposes a practical system for choosing and tabulating properties to reduce labour and attain accuracy. The problem becomes far more complicated with a mixture, where separation must precede recognition, and separation is naturally the more difficult operation. To separate one substance from others, it must be brought into a condition where it is detached by a surface of separation. Such surfaces occur chiefly between different states of aggregation, though they are not excluded when the state is the same. Methods of separation are therefore referred to the different states of aggregation: solids from solids, solids from liquids, liquids from liquids, solids or liquids from gases, and gases from gases. Crucially, separation is always a mechanical operation; a “chemical separation” consists of transforming substances by chemical means into other substances that can then be separated mechanically.
Separation of Solids from Solids
The principle for separating solids is to allow forces to act on one constituent, transporting it to a spot from whence it can be removed. Specific gravity differences are the most widely applicable property for this. If a mixture of two solids is stirred in a liquid whose specific gravity lies between theirs, the lighter rises and the heavier sinks. If the specific gravities are known, the liquid’s density is regulated accordingly; if unknown, one begins with a liquid denser than both and lowers its density by adding a lighter liquid until separation occurs. With more than two solids, the density is gradually lowered in short steps, causing the heaviest to fall first for removal, followed by the next heaviest, and so on. Suitable liquids for water-insoluble substances include aqueous solutions of potassium-mercury iodide, barium-mercury iodide, and cadmium borotungstate. For soluble substances, methylene iodide or acetylene tetrabromide, diluted with light liquids such as xylene, are used. As these liquids’ specific gravity does not much exceed 3, denser solids cannot be separated by them, though molten substances of higher specific gravity may be used in some cases. Elutriation, or “washing,” is a similar but less perfect procedure where solids in fine division sink faster the denser they are, and a stream of liquid carries away the less dense constituents. However, the rapidity of sinking depends not only on specific gravity but greatly on particle size, as smaller particles subside more slowly. This complex relationship makes the process unsuited to exact separations, and it is practical only when specific gravity differences are marked. Beyond hydrostatic forces, none are of general applicability for separating solids, though in particular cases, magnetic forces can be used, such as removing particles of iron with a magnet.
Washing Precipitates
The discussion turns to the calculations governing the washing of precipitates. When the washing liquid amounts to nine times the original moistening solution, and one gramme of foreign substance is initially present, after four washings only 0.0001 gramme of impurity remains. The question of optimal washing with a fixed total liquid volume requires differential calculus, which the author omits, but the result is stated clearly: washing many times with small quantities is superior to few washings with large quantities.
Adsorption Phenomena
These calculations, first made by Bunsen, disagree with observed facts. Theory suggests four washings with ten times the adhering liquid volume should suffice, but experience shows precipitates remain far from pure. The error lies in assuming the impurity is reduced to the m-th part by mixing with (m−1) times water and filtering. What is omitted is adsorption, the surface condensation on solid bodies. At contact-surfaces between solids and liquids, the concentration of a dissolved substance differs from—and usually exceeds—that in the bulk solution. Consequently, more dissolved substance remains after draining than the adhering liquid alone would contain, and each washing removes less than theory predicts. Almost nothing is known about the laws of adsorption, but the adsorbed amount is probably proportional to the surface area and depends on the nature of the solid and dissolved bodies and the solution’s concentration. The author presents a calculation assuming adsorbed impurity is proportional to concentration, yielding a formula that resembles the earlier one but with the solvent’s action only partially effective. This coefficient probably diminishes rapidly as the adsorbed substance decreases, explaining the practical difficulty of washing out the final traces. The rule holds: add wash-water in small portions, allowing thorough draining. Adsorption also occurs on filter paper cellulose, so when a muddy liquid is partially filtered through a dry filter, the first drops must be discarded as they are less concentrated due to adsorption by the paper. The filter soon reaches equilibrium, and subsequent runnings match the original concentration. Alkaline liquids show this phenomenon markedly, while acids and neutral salt solutions show it less.
Enlargement of Crystalline Grains
Fine crystalline powders convert to larger grains during digestion due to surface tension at solid-liquid boundaries. Smaller crystals are somewhat more soluble than larger ones, a fact that follows necessarily from surface-tension energy considerations. The liquid remains supersaturated regarding larger crystals, so small crystals dissolve while large ones grow. No substance is absolutely insoluble; even silver chloride, bromide, and iodide have measurable solubility. The rate of change increases with solubility and with temperature. Coarse crystals are preferable as they filter quickly, are purer, and wash more easily since adsorbed impurity is proportional to surface area.
Colloidal Precipitates
Many amorphous bodies show indeterminate solubility, forming pseudo-solutions between true solutions and suspensions. They precipitate upon warming, adding foreign bodies, or evaporation. Ferric oxide, alumina, silicic acid, and most precipitated metallic sulphides belong to this class, occurring as gelatinous or flocculent precipitates that stop up filters or pass through after washing. Salt solutions precipitate all such substances, and acids and bases act similarly but more strongly. When salt is removed or diluted, many colloids re-dissolve, which is why a precipitate “goes through” a filter during washing.
Salt Ionisation and the Classification of Electrolytes
Ostwald begins this section by establishing a clear hierarchy of ionisation among different classes of substances. Neutral salts are the most strongly ionised, and in the medium-concentration aqueous solutions typically used in analytical work, much more than half of the total salt is usually present as free ions. The degree of dissociation is highest for salts with monovalent ions, such as potassium chloride, silver nitrate, and ammonium bromide, and it decreases as the valency of the ions rises. The specific nature of the metal or acid radicle otherwise has little influence. Notable exceptions to this rule are the mercury halogen compounds, which are ionised very slightly, and cadmium compounds, which are somewhat more so; zinc compounds stand midway between cadmium and the general run of salts. Among the halides themselves, iodine compounds show the smallest ionisation and chlorine compounds the greatest.
Acids and bases display degrees of ionisation that correspond directly to their familiar “strength.” The strongest acids—the halogen hydrides (except hydrofluoric acid), nitric, chloric, perchloric, sulphuric, and polythionic acids—are ionised to the same extent as neutral salts. Moderately strong acids like phosphoric, sulphurous, and acetic acids are not ionised beyond about ten percent under ordinary conditions, while weak acids such as carbonic acid, sulphuretted hydrogen, hydrocyanic, silicic, and boracic acids are ionised to less than one percent, with the last two being scarcely measurable. A parallel classification holds for bases: the hydroxides of the alkali and alkaline earth metals, thallium hydroxide, and quaternary ammonium compounds are strong; ammonia, the fatty-series amines, silver oxide, and magnesia are moderate; and the hydroxides of di- and trivalent metals, the aromatic-series amines, and most alkaloids are weak.
The Law of Mass-Action and Its Application to Electrolytes
Ostwald then introduces the fundamental law governing chemical equilibrium. For a reversible reaction where substances A combine to form substances B, the product of the concentrations of the reacting substances, each raised to the power of its molecular coefficient, is proportional to the same product for the products of the reaction. This law, first propounded by Wenzel over a century earlier, states that the chemical action of any substance is proportional to its acting mass or concentration. The law is only approximate when concentration is defined as amount divided by total volume, becoming truly accurate only for dilute solutions. A crucial refinement is that ions must be counted as separate chemical substances in these equations, so a completely dissociated salt like potassium chloride must be written as its constituent ions rather than as an intact molecule. The apparent exceptions to the mass-action law were reconciled with it through the dissociation theory, with one limitation being that positive and negative ions must always be present in equivalent quantities.
For a binary electrolyte that splits into two ions, the relationship between ionisation and dilution can be expressed mathematically. If α represents the fraction of the electrolyte that is ionised, then as dilution increases, α becomes larger, approaching complete ionisation at infinite dilution. The equilibrium constant is very high and similar for all neutral salts, but it varies enormously for acids—being high for strong acids and low for weak ones. Crucially, the differences in ionisation among acids diminish with increasing dilution, so that infinitely dilute solutions of different acids possess the same strength, since the acid is then completely ionised. The same principle applies to bases.
Gradual and Complex Dissociation
Salt ions are not absolutely stable entities. They may undergo hydrolysis or further dissociation. Complex ions, such as the metallo-ammonium ion Ag·4NH₃, dissociate into a metallic ion and free ammonia, while the complex ion Ag(CN)₂ of potassium-silver cyanide dissociates electrolytically into a silver ion and two cyanide ions. These dissociations follow the same laws as other chemical equilibria. For electrolytes with ions of unequal valency, such as a dibasic acid, dissociation does not occur in a single step. A dibasic acid H₂A first dissociates into a hydrogen ion and the acid ion HA′, which then further dissociates into another hydrogen ion and the anion A″. The dissociation constant for the second step is always much smaller than that for the first. This means that the hydrogen atoms of a polybasic acid have different “strength” values—the first is always more strongly acidic than the second, and so on.
Alkaline and Acid Reactions of Salts
When hydrogen ions are present in a solution of a neutral salt of a weak acid, the effect is that both free anions and a corresponding quantity of non-ionised acid molecules exist. For a strong acid with a high constant, the concentration of non-ionised acid remains very small, but for a weak acid with a low constant, it becomes measurable. Potassium cyanide exemplifies this, as its aqueous solution contains a perceptible quantity of hydrocyanic acid, detectable by its odour. Such salts characteristically show an alkaline reaction, which depends on the presence of hydroxyl ions. The formation of non-ionised acid takes hydrogen ions from water. Since the product of the hydroxyl and hydrogen ion concentrations in water must remain constant, a diminution of hydrogen ions causes a proportionate increase in hydroxyl ions, which can then become measurable. Identical considerations apply to salts of weak bases, which show an acid reaction. When both the acid and the base are weak, these processes mutually support one another, producing appreciable quantities of both non-ionised acid and base.
Heterogeneous Equilibrium and the Law of Distribution
Ostwald then turns to heterogeneous equilibrium, which exists when a system is divided into portions by surfaces of separation. In two contiguous phases, each substance present in both must bear a constant ratio of concentrations, a ratio that depends on the nature of the substance and the temperature. This law applies to gas absorption, the solubility of solids in liquids, and vapour pressure. In these cases, one phase—the solid or the pure liquid—maintains a constant concentration while altering only in amount. This justifies distinguishing states of constant concentration (solids and simple liquids) from those of variable concentration (gases and dissolved substances).
Velocity of Reaction, Temperature, and Catalysis
The analyst must often follow the progress of a reaction, although most analytical methods use ion-reactions that complete immeasurably quickly. The velocity of a reaction is directly proportional to the concentration of each reacting constituent. All reaction courses begin at their greatest velocity and decrease progressively, theoretically completing only after infinite time, though practically the residual action becomes non-measurable after a relatively short period. Temperature markedly affects reaction velocity, with measured instances showing that velocity doubles with approximately a ten-degree rise. Catalytic actions involve substances that influence velocity without apparent participation. Acids act as general catalysers, hastening slow reactions in proportion to their free hydrogen ion concentration, while specific catalysers, such as iron salts in oxidations, and gaseous catalytic actions involving finely divided metals with large surfaces, are also noted.
Precipitation, Super-saturation, and the Solubility-Product
Precipitation occurs when the constituents of an incompletely soluble substance meet in solution, but it is always preceded by a state of super-saturation, and no precipitation is ever perfect. Super-saturated solutions may persist indefinitely without contact with a solid, and deposition requires either the presence of the solid or a sufficient degree of super-saturation. More soluble solids show super-saturation more readily, as demonstrated by the sulphates of barium, strontium, and calcium: barium precipitates almost instantaneously, strontium requires measurable time, and calcium sulphate may remain supersaturated for months. Precipitates are almost always electrolytic, and their solutions contain mainly ions. For a saturated electrolyte solution, the solid equilibrates with a minimal non-ionised portion, which in turn equilibrates with its dissociated ions. The product of the concentrations of these ions is a constant, known as the solubility-product. When this product is exceeded, the liquid is super-saturated; when it is not reached, the liquid dissolves the solid. Adding an excess of the precipitant reduces the concentration of the precipitable ion, though never to zero, and a greater excess is needed for more soluble precipitates. Washing precipitates should therefore use solutions containing a common ion—lead sulphate washes better with dilute sulphuric acid, and mercurous chromate with a nitrate solution.
Precipitation with Acids and Bases
The conditions for precipitation become more complex when acids and bases are involved, as any degree of ionisation may occur. Calcium salts are precipitated by all soluble carbonates but not by free carbonic acid, because soluble carbonates ionise normally, producing a sufficient concentration of carbonic acid ions, whereas carbonic acid itself, being excessively weak, yields only minute proportions of these ions. Lead salts present a more complicated case. Since lead carbonate is less soluble than calcium carbonate, carbon dioxide can precipitate lead from a moderately concentrated solution. As the reaction proceeds, the hydrogen ions liberated from the carbonic acid accumulate, hindering further ionisation of the acid until precipitation ceases.
The Law of Successive Reactions and Its Practical Consequences
The discussion opens with a deceptively simple question: if a system is forced into its least stable state and then left to its own devices, which state will it finally reach? One might naturally assume that the most stable state—the last in the series—would be the ultimate product. Yet experience, gathered by various investigators in special cases and generalised by the author, shows otherwise. It is not the most stable state that is first produced, but rather the least stable one that is still possible, meaning the second in the series of stability.
This law is illustrated by numerous familiar phenomena. In precipitation reactions, a supersaturated solution is always first produced, which later—sometimes after a considerable interval—yields the solid form, as when potassium salts are precipitated by tartaric acid. When several solid forms are possible, the more unstable and soluble ones appear first; nearly every precipitate comes down amorphous at the moment of precipitation, and if filtered immediately, an appreciable amount of what would ultimately precipitate remains in solution. When mercuric chloride is precipitated by stannous chloride, calomel is thrown down first, however much reducing salt is added, with metallic mercury appearing only as the final product. Similarly, oxidising agents acting on oxidisable substances do not immediately form complete oxidation products but intermediate-stage compounds, even though these intermediates seem more readily oxidisable. The oxidation of alcohol by chromic acid produces aldehyde first, though aldehyde is more easily oxidised than alcohol itself, taking up oxygen directly from air while alcohol requires a ferment or catalyser.
The practical application of this principle is significant for analytical chemistry. Analytical precipitations, though mostly ion reactions in solution, require appreciable time for quantitative accuracy. Teachers must impress upon pupils the need to allow sufficient time for reactions to complete, and the author hopes these paragraphs help teachers justify this rule and help students understand its basis.
Chapter V: The Quantitative Determination of Substances
Qualitative analysis ends with the recognition of a substance, often involving the separation of constituents. The further question—how much of each substance is present—requires the measurement of quantities. Measurement need not always be preceded by separation from everything else; common practice estimates one constituent in a complex sample without heeding the others. Methods are thus considered separately: those requiring previous separation and those not requiring it.
For weighing, a substance must be unalterable in air, non-deliquescent, and preferably able to withstand red heat without chemical change. Few substances satisfy these conditions, so the law of constant mass-proportions is extensively used: a compound unsuited for measurement is converted into one with the desired properties, and the quantity is calculated from the measured quantity of the latter. This ratio follows from the law of combining weights—all chemical combinations occur in ratios of definite relative numbers, the atomic weights. The sum of atomic weights in a compound is its combining weight, and the relation between the combining weights of two substances, one preparable from the other, gives the conversion ratio. Given atomic weights and the chemical equation, one can calculate the coefficient reducing the original substance’s amount to the final product’s amount, and vice versa.
The author provides a concrete example: hydrochloric acid in solution is estimated by acting on excess carbonate of lime, passing the carbon dioxide through baryta solution, and weighing the precipitated barium carbonate. From the formula-weights, the coefficient is 0.3701, which multiplies the barium carbonate weight to give the original hydrochloric acid. What holds for weighing applies to other estimation modes; transformations must be easy and complete, and the theoretical factor should be checked empirically against weighed original substance. Methods with secondary reactions are suspect and rejected or used only in default of better ones.
Pure Substances and Methods of Measurement
After separation, the simplest reliable determination is weighing. The balance determines gravitational force, which is proportional to mass, and other amount-varying properties such as volume and chemical energy are likewise proportional to mass. For gases, volume measurement is usually preferable to weighing, since a gas’s weight is a small fraction of its container’s weight, magnifying weighing errors. Temperature and pressure effects are eliminated by reducing to normal conditions using the standard formula, though rigorous calculations must note that the 760-millimetre mercury column is not a perfect definition of normal pressure, since it depends on gravity’s intensity, which varies with geographical latitude and height above sea-level. The author suggests abandoning the old normal-pressure definition for absolute units.
For liquids, only temperature need be allowed for, as compressibility and atmospheric pressure variations are negligible. No general expansion law exists, so expansion must be determined for each substance, and usually only apparent expansion—liquid expansion minus vessel expansion—matters. Volume measurement of solids is never applied in analysis for weight determination, since solids cannot fill empty spaces as liquids or gases do. Attempts to determine precipitate amount without washing, from the mean specific gravity of precipitate plus liquid and the liquid’s specific gravity alone, have proven inaccurate, as precipitate density varies and is influenced uncontrollably by adsorption.
Physical Methods and Their Limitations
Cases where physical methods are practically applicable are not common and are not very exact, because the special property being measured is rarely independent of outside substances. Examples include the determination of cane sugar from its optical rotatory power and various colorimetric analyses. Serious errors have repeatedly been found in colorimetric analyses, arising from the assumption that foreign bodies had no influence when their influence was material.
Methods of Titration
Chemical methods are more reliable and numerous than physical ones. They are based on subjecting the substance to a chemical reaction with a suitable reagent, where either the complete conversion of the original substance or the slightest excess of reagent can be recognised by a striking sign. It is easier to judge whether the condition is infringed—that no other substance present shall exert an influence—in chemical methods, making them widely applicable. The quantitative estimation depends on measuring the quantity of reagent needed to completely transform the substance, usually by noting the volume of the reagent-solution used, though for exceptionally accurate determinations, weighing the reagent is preferred as it is unaffected by temperature variations.
Titration methods are divided into two groups: those where the disappearance of the original substance furnishes the end-phenomenon, and those where the excess of reagent does. The second group is further divided into cases where the excess is directly apparent and those where it is only rendered visible through an added medium—the indicator. Iodometric analysis exemplifies the first group, where the reaction transforms dark-coloured free iodine into colourless sodium iodide, with thiosulphate solution added until the colour vanishes. The estimation of iron by permanganate of potash exemplifies the first section of the second group, where the pink permanganate colour remains permanent once all ferrous salt is used up. The typical example of the indicator method is acidimetry or alkalimetry, where a dye like litmus is added, giving a blue salt with alkali that is decomposed by the slightest excess of acid.
Since the time of Friedrich Mohr, the strength of volumetric reagents has been regulated so that an equivalent in grammes is in a litre or a multiple of a litre, and the number of cubic centimetres used gives the amount of substance in milligramme-equivalents.
Applications
The author states his object is to give a living interest to the laws detailed in Part I by discussing the analytical properties of different substances. He is not attempting to teach analytical chemistry to beginners, as he knows time must elapse before the new views are accepted. Students should ponder the scientific principles underlying what they have been taught. A footnote records that since the first German edition of 1894, the author has received letters announcing that introducing these new views into elementary courses has proved highly successful. The division of the subject follows the usual analytical groups, with the treatment of the ionic state being essentially new, since analytical reactions are, with very few exceptions, reactions of ions.
Chapter VI: The Hydrogen and Hydroxyl Ions
Compounds whose aqueous solutions contain the hydrogen ion are defined as acids, and those containing the hydroxyl ion as bases. They are recognised qualitatively by colour reactions in dyes used as indicators. For a dye to be an indicator, it must be acid or basic, have a different colour when non-dissociated from its ionic condition, and must not be a strong acid or base, since a strong acid would break up into ions in the free state, showing no colour change on neutralisation. A weak acid exists mostly as undissociated molecules, with ions produced only after neutralisation into a neutral salt. The properties of an indicator depend on its dissociability; a very weak acid indicator will be sensitive, as acids of moderate or even small ionisation will give up hydrogen to it. Phenol-phthalein is a very weakly acid indicator, colourless in the molecular state.
Chapter VII: The Alkali Metals
Potassium, rubidium, caesium, sodium, and lithium always exist in solution as monovalent positive ions, never as complex ions, consistently showing reactions of those simple ions without anomalous reactions. Their hydroxides are readily soluble, highly ionised, and constitute the strongest known bases. With ordinary precipitating reagents they form only soluble salts, remaining in solution after other metals precipitate—a property exploited for separation.
Permanganate Titration and Detection
The intense color of permanganate solutions renders indicators unnecessary, making this one of the rare instances of titration without one. The permanganate ion can be detected with even greater delicacy using a spectroscope than with the naked eye. Its spectrum, which is identical for solutions of every salt of the acid since the same colored ion is present in each, displays five dark bands in the yellow and green regions and remains distinctly visible at dilutions where the color is imperceptible to the eye alone.
Cobalt and Nickel
Cobalt and nickel lack the capacity to form trivalent ions. They do form higher oxides, but these are peroxide in character—insoluble in dilute acids and liberating chlorine when treated with hydrochloric acid. Only divalent ions are dealt with in analytical practice, along with complex compounds possessing special reaction properties. The cobalt ion is red, while the nickel ion is emerald green.
Non-ionised cobalt salts are mostly dark blue; as solutions become concentrated, the diminishing ionisation causes a colour change from red to blue. Warming and adding strongly ionised salts with a like ion produce this effect, best observed by adding concentrated hydrochloric acid to cobalt chloride. Contrary to prevailing views, there is no contradiction in ionisation being reduced by warming, as this has been proven for ionic dissociation.
Cobalt and nickel exhibit a peculiarity: their salts are not precipitated by sulphuretted hydrogen from acid solution, yet once precipitated, the sulphides are no longer soluble in dilute acid. This anomaly remains unexplained. Two conjectures are offered: first, that the sulphides change into a less soluble form immediately after precipitation; second, that they exist only in a sparingly soluble form, but persistent supersaturation phenomena occur in acid solution. The latter is considered less probable since the sulphides precipitate readily from acetic acid solution. A footnote cites the work of A. Villiers, which favours the first explanation.
Cobalt salts show a greater capacity for forming complex ions than nickel, which forms the basis for separation methods. The most convenient method treats mixed solutions with potassium nitrite in acetic acid solution, forming potassium-cobalt nitrite, which is insoluble in excess potassium nitrite. Precipitation proceeds slowly, requiring the liquid to stand for several hours for practically complete separation—proof of complex salt formation rather than an ordinary ion reaction. Nickel gives no such insoluble compound.
Another separation method employs complex cyanogen compounds: cobalt’s is extremely stable, not decomposed by acids even on boiling, while the nickel compound precipitates sparingly soluble nickelous cyanide under those conditions. Ammonium compounds show similar stability differences: both metals precipitate as hydroxides with ammonia, redissolving in excess. Nickel-ammonium compounds decompose readily, losing ammonia in air when solid, whereas cobalt compounds oxidise to stable complexes unaffected even by warming with alkali, belonging to quite different types.
Zinc
Zinc forms only divalent ions; higher oxidation products are unknown. Zinc hydroxide may give up hydrogen as an ion, producing the faintly acid ion Zn0₂″, and zinc also appears in complex ions with cyanogen, ammonia, and similar agents. Zinc oxide, insoluble in water, dissolves in potash, soda, and ammonia—but for different reasons: in the first case through formation of Zn0₂″ ions, and in the second through complex zinc-ammonium ions. These latter are fairly stable, so the hydroxide is not split up hydrolytically; zinc oxide dissolves in ammonia without excess ammonium salt present, though this explanation is noted as not quite certain.
Zinc sulphide is less soluble than other sulphides of this group. Most zinc separates when sulphuretted hydrogen passes through neutral zinc salt solutions of strong acids. How much zinc remains depends on the acid’s nature: equilibrium with solid zinc sulphide is determined by the product of zinc and sulphur ion concentrations. Since sulphuretted hydrogen is a very weak acid, the sulphur ion concentration is inversely proportional to the free hydrogen ion concentration from liberated acid. The total sulphuretted hydrogen concentration may be considered constant. The less the acid ionises and the more concentrated the zinc salt solution, the less zinc escapes precipitation. Since weak acid ionisation can be lowered by adding neutral salts, the old method of precipitating zinc with excess sodium acetate is theoretically justified. To completely prevent precipitation, sufficient strong acid must be added—free hydrogen ions lower hydrogen sulphide ionisation so the solubility-product value is never reached. The added acid must be approximately proportional to the amount, or rather the square root of the amount, of zinc salt.
Chapter X: Metals of the Copper Group—General Properties
These metals differ from the iron group by their sulphides’ insolubility in dilute strong acids—merely a difference of degree, with intermediate stages in cadmium and lead. The metals differ considerably among themselves. Some are “noble” metals; others may be called “nobler” than iron group metals. This word expresses a definite property: ionisation tendency, measured by the free energy given out when an equivalent of metal transitions into the ionic state. A greater tendency means the metal passes more easily and quickly into the ionic condition. Potassium has the greatest value; aluminium, zinc, tin, and cadmium are smaller; lead is almost zero; copper, antimony, bismuth, silver, and gold are negative—free energy is given out when their ions change into the metallic state. This applies only to measurable ion concentrations; at very small concentrations, every metal is displaced toward the less noble side. Where noble metals begin is determined not by this but by whether the metal is oxidised by free oxygen. Metals with positive ionisation tendency dissolve in acids with hydrogen evolution, since hydrogen’s ionisation tendency is almost zero.
Silver Complex Ions and Reactions
Silver forms a complex ion with thiosulphate anions that is somewhat more stable than its ammonia compound, with silver uniting directly with sulphur. Sodium thiosulphate dissolves all silver compounds soluble in ammonia plus some that are not soluble or only very sparingly soluble, such as silver bromide. The most stable complex silver compound is the silver cyanide ion, Ag(CN)₂⁻. This forms so quickly and easily that the reaction enables titration of cyanogen: the liquid is made alkaline, silver nitrate solution is run in; while cyanogen ions remain in excess the solution stays clear, but once the ratio of one silver atom to two cyanogen molecules is exceeded, silver cyanide precipitates.
Potassium cyanide dissolves all silver salts except the sulphide, which is—next to mercury sulphide—the most insoluble sulphide of this group. Silver’s affinity for sulphur is demonstrated by its behaviour with thiosulphates. For the same reason, metallic silver decomposes sulphuretted hydrogen with liberation of hydrogen, just as zinc breaks up hydrochloric acid; the “noble” metal character entirely disappears. Silver also dissolves in potassium cyanide solution with hydrogen evolution; experiments confirm distinct solubility. Silver salts are always used to test for halogens, though they react only when halogens are present as ions. Slow precipitation by silver salts with organic halogen compounds suggests such compounds can be dissociated in minute degree.
Mercury
Mercury forms mono- and divalent ions—the former resembling silver ions, the latter cadmium ions. Many similarities exist between mercury and copper. Mercury characteristically forms compounds of low ionisation, producing numerous “abnormal” reactions. The monovalent (mercurous) ion yields sparingly soluble halogen compounds whose solubilities follow the same order as the silver series. Unlike silver salts, mercurous compounds form black insoluble ammonia compounds; silver chloride dissolves in aqueous ammonia, but mercurous chloride turns black. Resemblances to copper exist in mono-/divalent transformations, with striking differences: mercurous oxide passes readily into mercuric, yet mercuric halides are far more stable than mercurous ones, contrary to copper compounds. Mercurous sulphide has no existence, decomposing immediately into mercuric sulphide and metallic mercury. Mercuric sulphide is exceptionally stable—the only sulphide not dissolving in nitric acid. It resembles next-group sulphides in dissolving in potassium sulphide, but only in concentrated solution with caustic alkali present; dilution reprecipitates it. Ammonium sulphide does not dissolve it.
The divalent mercuric ion in oxygen salts, normally dissociated, is very weakly basic; its salts hydrolyse extensively in aqueous solution, requiring excess free acid for clarity. Halogen compounds dissolve perfectly stably but deviate from oxygen salts in their reactions. Electric conductivity shows mercury halides ionise extremely slightly, so solutions contain very few mercuric ions. Compounds of slight ionisation result when constituent ions meet; thus mercury oxygen salts assume halide reactions when contacting soluble halogen compounds. Considerable heat evolves—contrasting with ordinary neutral salt interactions where thermal neutrality holds. Mercuric oxide acts as a very weak base toward most acids, but added to soluble halogen compounds the liquid becomes strongly alkaline. The action is weakest with chlorides and strongest with iodides—potassium iodide is decomposed by mercuric oxide to 90 per cent.
The Tri-Metallic Phosphates and the Limits of Aqueous Chemistry
Ostwald turns to a striking anomaly: the third hydrogen atom of phosphoric acid is almost impossible to replace in aqueous solution. While soluble tri-metallic phosphates of the alkali metals and ammonium exist, they are not what they seem. In water, they are almost completely hydrolyzed, breaking down into bi-metallic phosphate ions and free alkali. A solution of Na₃PO₄, therefore, contains sodium ions, the divalent HPO₄″ ion, and hydroxyl ions—but practically none of the trivalent PO₄‴ ion. The reason is that the tendency of the third hydrogen to ionize is far weaker than that of water itself. The moment a PO₄‴ ion forms, it reacts with water, pulling a hydrogen ion from it to become HPO₄″ and leaving behind a hydroxyl ion.
This difficulty vanishes entirely when we consider solid, sparingly soluble salts. The explanation offered is that developing a negative charge on an already negatively charged complex requires far more work than doing so on a neutral one, and this is even more true for a third charge. In a solid, non-ionised salt, this condition does not apply, so normal tri-metallic phosphates are perfectly stable in the solid state. Indeed, they are the only phosphates found in nature.
The precipitation of ordinary sodium phosphate (Na₂HPO₄) with a silver solution illustrates the complexity. The simple equation—3AgNO₃ + Na₂HPO₄ = Ag₃PO₄ + 2NaNO₃ + HNO₃—is imperfect, because the liberated nitric acid would dissolve some of the silver phosphate. A more accurate representation is 3AgNO₃ + 2Na₂HPO₄ = Ag₃PO₄ + 3NaNO₃ + H₂NaPO₄, yielding acid-reacting dihydrogen-sodium phosphate instead of free acid. The acidity of this product arises from the ionization of the monovalent H₂PO₄′ ion into a hydrogen ion and a divalent HPO₄″ ion.
The Law of Equal Valency and Complex Phosphates
From these observations, Ostwald draws a general law: compounds built from ions of equal valency tend to form sparingly soluble salts. Monovalent halogens are precipitated by monovalent silver, mercury, and copper cations; divalent alkaline earth metals by divalent sulphate, oxalate, and carbonate ions; and trivalent iron, chromium, and aluminium phosphates are insoluble in acetic acid. This law, however, cannot be reversed, as many equi-valent salts dissolve readily, indicating that another unknown condition influences solubility.
Phosphoric acid also forms complex compounds with tungstic and molybdic acids, where the basicity remains constant while the proportion of trioxide varies. Phospho-molybdic acid yields very sparingly soluble yellow salts with alkali metals and ammonium, dissolving little even in free acid, especially when one ion is in excess. This property is exploited to detect and quantitatively separate phosphoric acid from nitric acid solutions by adding an excess of molybdic acid and ammonium nitrate in nitric acid. The mixture must be warmed gently and allowed to stand for a considerable time, as this non-ionic reaction requires measurable time to complete. The yellow precipitate dissolves in ammonia, after which the phosphoric acid can be precipitated as ammonium-magnesium phosphate—a process used in analysis and for recovering molybdenum residues.
Pyro-, Meta-, Phosphorous, and Hypophosphorous Acids
Ortho-phosphoric acid transforms into pyro-phosphoric (H₄P₂O₇) and meta-phosphoric (HPO₃) acids by losing water. The meta-acid is not a true analogue of nitric acid but a condensed acid of higher molecular weight, and several meta-phosphoric acids exist with different molecular weights. Meta-acid, found in fused or vitreous phosphoric acid, precipitates albumen and gives a white silver salt. Pyro-phosphoric acid does not precipitate albumen but does give a precipitate with barium chloride, which ortho-acid does not. Neither meta- nor pyro-acid shows the reactions of ortho-acid with magnesia mixture or ammonium molybdate. These derivatives convert back to ortho-phosphate when fused with excess alkaline carbonate or warmed in strongly acid solution. For quantitative estimation, they are converted to the ortho-compound and precipitated as magnesium-ammonium phosphate.
Phosphorous and hypophosphorous acids are di- and monobasic, respectively, and are always transformed into phosphoric acid for estimation. Their characteristic behaviour is that heating them or their salts liberates spontaneously inflammable phosphuretted hydrogen with a separation of red phosphorus. They act as reducing agents, precipitating calomel from acid corrosive sublimate solution. Nascent hydrogen reduces them to phosphuretted hydrogen, but not phosphoric acid—a behaviour shared with sulphuric acid versus the lower sulphur acids, presumably connected to how hydrogen is linked. Phosphoric acid has the formula OP(OH)₃, with no hydrogen linked to phosphorus, while the others have OP(OH)₂H and OP(OH)H₂. Pure hypophosphorous acid shows a sharp neutralisation point with an indicator, characteristic of a fairly strong monobasic acid, whereas phosphorous acid resembles phosphoric acid in that its second hydrogen is much more difficult to substitute, so its neutral salts hydrolyze and give an alkaline reaction.
Publisher’s Catalogue
The volume concludes with a publisher’s advertisement for medical books, listing titles, authors, and press reviews. Among the works advertised are “Methods of Pathological Histology” by C. von Kahlden, translated by H. Morley Fletcher, which the British Medical Journal deemed valuable for presenting principal histological methods clearly. Also listed is “A Course of Elementary Practical Bacteriology” by A. A. Kanthack and J. H. Drysdale, and E. Klein’s “The Bacteria in Asiatic Cholera,” which Nature called a valuable criticism of Koch’s argument. Further entries include Ernst Ziegler’s “Text-Book of Pathological Anatomy,” T. Lauder Brunton’s works on pharmacology and therapeutics, and W. Hale White’s “Text-Book of General Therapeutics,” each commended by contemporary journals for their contributions to medical science.
Main Arguments
The central argument of the book is that analytical chemistry, properly understood, is not a collection of empirical recipes but a branch of physical chemistry governed by precise quantitative laws. Ostwald maintains that every operation the analyst performs—precipitation, washing, filtration, extraction, titration—can be understood as a particular application of the general principles of chemical equilibrium and electrolytic dissociation. The analyst who grasps these principles will not merely execute procedures more intelligently but will be able to devise new procedures and to judge the reliability of existing ones.
A second major argument concerns the nature of separation. Ostwald insists that separation is always, in the last analysis, a mechanical operation. Even when the analyst speaks of chemical separation, what is meant is that chemical reactions have been used to transform the substances present into new substances that can then be separated by purely mechanical means—filtration, decantation, distillation, or the like. This apparently simple observation has far-reaching consequences, for it implies that the entire art of qualitative analysis consists in finding chemical transformations that convert the constituents of a mixture into forms whose physical properties permit their mechanical division.
The book argues further that the theory of electrolytic dissociation provides the key to understanding why analytical reactions proceed as they do. The properties of a dilute salt solution, Ostwald maintains, depend on its free ions rather than on the undissociated salt. This insight simplifies the analyst’s task enormously, for instead of having to learn the reactions of the thousands of possible salts that could be formed from the fifty or so common cations and anions, one need only know the properties of the hundred or so individual ions. The theory also explains why certain reactions fail to occur even though the substances involved might seem capable of reacting: the reaction may require the presence of an ion that is not produced in sufficient concentration under the prevailing conditions.
The book advances a distinctive view of the conditions necessary for complete precipitation and for the effective washing of precipitates. Ostwald shows that no precipitation is ever perfect and that the completeness of precipitation is governed by the solubility-product—the constant product of the concentrations of the ions of a sparingly soluble substance in a saturated solution. From this principle follow such practical rules as the desirability of adding an excess of the precipitating reagent and the advantage of washing precipitates with solutions containing an ion common to the precipitate rather than with pure water.
A further argument concerns the role of time in analytical operations. Ostwald observes that many precipitates initially form in an amorphous or supersaturated state and only gradually pass into their final crystalline condition. The law of successive reactions, which he formulates as the principle that the least stable state possible is the one first produced, implies that the analyst must allow sufficient time for reactions to reach completion. The book urges teachers to impress this necessity upon their students and provides the theoretical justification for what might otherwise appear to be mere procedural caution.
The book also argues for the superiority of chemical methods of quantitative determination over physical ones. While physical methods based on such properties as optical rotation or colour may be convenient in special cases, they are rarely independent of the influence of foreign substances and have repeatedly led to serious errors. Chemical methods, by contrast, can be checked more readily for the influence of other substances present and are therefore more reliable and more widely applicable.
Key Concepts
The theory of electrolytic dissociation stands at the centre of the book’s conceptual apparatus. According to this theory, salts in aqueous solution are dissociated into their constituent ions, which behave as independent chemical individuals. The degree of dissociation varies with the nature of the salt and with the dilution of the solution, increasing as the solution becomes more dilute. Neutral salts are the most strongly dissociated, with salts of monovalent ions showing the greatest dissociation. Acids and bases exhibit degrees of dissociation corresponding to their strength as ordinarily understood: the strongest acids and bases are the most completely ionised, while weak acids such as carbonic acid and weak bases such as the hydroxides of the di- and trivalent metals are ionised to only a very slight extent.
The law of mass-action, formulated in its general mathematical form, governs all chemical equilibria. For a reversible reaction involving several substances, the product of the concentrations of the reacting substances, each raised to the power corresponding to the number of molecules participating in the reaction, bears a constant ratio to the corresponding product for the products of the reaction. This law, first propounded by Wenzel more than a century before the book was written, had been reconciled with apparent exceptions through the dissociation theory and the treatment of ions as separate chemical individuals.
The solubility-product is a particular application of the law of mass-action to saturated solutions of sparingly soluble electrolytes. For such a substance, the product of the concentrations of its ions in a saturated solution is constant at a given temperature. When this product is exceeded, the solution is supersaturated and precipitation will occur; when it is not reached, the solution can dissolve more of the solid. This principle governs both the precipitation of substances and their redissolving, and it explains why the completeness of precipitation depends on the addition of an excess of the precipitating reagent.
Adsorption is treated as a phenomenon of great practical importance for the analyst. At the contact surfaces between solids and liquids, the concentration of dissolved substances differs from—and usually exceeds—that in the bulk of the solution. This surface condensation means that precipitates retain impurities far more tenaciously than the simple theory of washing would predict. The book presents calculations showing that the effectiveness of washing is only partial because of adsorption and derives the practical rule that wash-water should be added in small portions with thorough draining between each addition.
The law of successive reactions states that when a system can pass through several states of increasing stability, the least stable state possible is the one first produced. This principle explains why precipitates initially form in amorphous or supersaturated conditions and only gradually pass into their final crystalline form, why calomel rather than metallic mercury is produced when mercuric chloride is reduced, and why oxidation reactions produce intermediate products rather than proceeding directly to the final stage of oxidation.
The distribution coefficient governs the partition of a substance between two immiscible solvents. When a solution is shaken with a second solvent, the substance distributes itself between the two layers in a constant ratio, and repeated extractions with fresh portions of the second solvent remove the substance according to a geometrical progression. The book shows that many small extractions are more efficient than a few large ones and that the efficiency of the operation depends on the numerical value of the distribution coefficient.
The electro-chemical series arranges the metals in order of their tendency to pass into the ionic state. Metals at one end of the series, such as potassium, pass readily into ionic condition, while those at the other end, such as silver and mercury, require an expenditure of energy for this conversion. This series governs the behaviour of metals in electrolytic separations and explains why some metals can be deposited from their solutions by electric currents while others cannot.
Themes
A recurring theme of the book is the unity of chemical knowledge. Ostwald continually shows that phenomena which might appear unrelated—the washing of a precipitate, the extraction of a substance by shaking with an immiscible solvent, the distillation of a mixture of liquids—are in fact governed by the same mathematical laws. The same geometrical progression that describes the removal of an impurity from a precipitate by repeated washing also describes the removal of a dissolved substance by repeated extraction. The same principles of equilibrium that govern the vapour pressure of a liquid govern the solubility of a solid and the distribution of a substance between two solvents. This insistence on underlying unity gives the book its intellectual coherence and its persuasive power.
The relation between theory and practice is another pervasive theme. Ostwald does not present theory as an ornament to practice but as its essential foundation. The analyst who understands why a procedure works will perform it more reliably and will be able to judge when it is applicable and when it is not. Conversely, the book repeatedly shows how practical problems have stimulated theoretical developments and how theoretical insights have suggested new practical procedures. The theory of electrolytic dissociation, for example, not only explains why certain reactions occur but also indicates how conditions may be adjusted to make them occur more completely.
The theme of simplification through theory appears in the book’s treatment of the enormous variety of chemical substances. The analyst who must deal with the reactions of thousands of possible salts is relieved of this burden by the recognition that the properties of dilute solutions depend on their ions rather than on their undissociated molecules. Instead of learning the reactions of all possible combinations of fifty cations with fifty anions, one need only know the properties of the hundred individual ions. The theory thus reduces an apparently unmanageable multiplicity to a comprehensible unity.
The book also sounds a cautionary note about the limits of analytical methods. No precipitation is ever perfect; no washing ever removes all impurities; no extraction ever completely transfers a substance from one solvent to another. The analyst must therefore understand not only how to perform operations but also how completely they can be performed and what errors remain inherent in them. This theme of the inevitable imperfection of analytical operations gives the book a note of scientific honesty and intellectual maturity.
The Author’s Conclusions
Ostwald concludes that analytical chemistry has been placed on a truly scientific footing by the theoretical developments of recent decades. The theory of electrolytic dissociation, the law of mass-action, and the recognition that dilute solutions obey laws analogous to those of gases have together provided a unified explanation of the phenomena with which the analyst deals. The author expresses confidence that the remaining difficulties are matters of detail rather than of fundamental principle and that continued theoretical progress will further illuminate the practice of analysis.
The book concludes its theoretical portion with the observation that the principles developed earlier find their application in the systematic treatment of particular groups of substances. The author states that his object in the second division of the work is to give a living interest to the theoretical principles by showing how they govern the behaviour of the metals and non-metals in actual analytical practice. The applications portion demonstrates that the separation and determination of the various groups of elements proceed according to the general laws already established and that the analyst who understands these laws can approach each new analytical problem with confidence in the rationality of the procedures employed.
The final chapter on the calculation of analyses treats the numerical operations by which the results of analytical procedures are converted into statements about the composition of the original material. The book emphasises that these calculations rest on the law of constant mass-proportions and the related law of combining weights, which specify the definite numerical ratios in which substances combine. The conversion of a weighed quantity of one substance into the corresponding quantity of another is accomplished through the ratio of their combining weights, a ratio that follows directly from the chemical equation describing the transformation. The book insists that the theoretical factors used in such calculations should be checked empirically against weighed quantities of the original substance and that methods involving secondary reactions should be regarded with suspicion.
Significance
The book occupies an important position in the history of analytical chemistry as one of the first sustained attempts to place the discipline on a firm theoretical foundation. Its significance lies not in the novelty of the individual procedures it describes—most of these were already familiar to practising analysts—but in its demonstration that these procedures are governed by general laws that can be stated precisely and applied systematically. The book thus contributed to the transformation of analytical chemistry from an empirical craft into a branch of physical chemistry, a transformation that had profound consequences for both the teaching and the practice of the discipline.
The work is also significant for its pedagogical philosophy. Ostwald believed that students should understand the reasons behind laboratory procedures and that this understanding would make them better analysts than mere mechanical execution of recipes could. The book is organised to lead the student from general principles to particular applications, always showing how the former illuminate the latter. This approach influenced the teaching of analytical chemistry for generations and helped to establish the expectation that a course in analysis should include a substantial theoretical component.
The book’s treatment of the theory of electrolytic dissociation and its applications to analytical problems represents an important moment in the assimilation of that theory into the mainstream of chemical practice. By showing how the theory explained phenomena that had previously been understood only imperfectly, and by deriving from it practical rules of great utility, the book helped to establish the theory’s credentials among chemists who might otherwise have regarded it as an abstract speculation with little bearing on their work.
The work also stands as a notable example of the scientific monograph addressed to students rather than to specialists. Ostwald writes with clarity and a concern for the reader’s understanding, taking care to explain the reasoning behind each step and to connect each topic to those that precede and follow it. The book thus exemplifies a style of scientific writing that seeks not merely to inform but to educate in the fullest sense—to develop in the reader the habit of understanding phenomena through their underlying principles rather than through memorised facts.
