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Bible Encyclopedias
Solution
1911 Encyclopedia Britannica
(from Lat. solvere, to loosen, dissolve). When a solid such as salt or sugar dissolves in contact with water to form a uniform substance from which the components may be regained by evaporation the substance is called a solution. Gases too dissolve in liquids, while mixtures of various liquids show similar properties. Certain solids also consist of two or more components which are united so as to show similar effects. All these cases of solution are to be distinguished from chemical compounds on the one hand, and from simple mixtures on the other. When a substance contains its components in definite proportions which can only change, if at all, by sudden steps, it may be classed as a chemical compound. When the relative quantities of the components can vary continuously within certain limits, the substance is either a solution or a mixture. The distinction between these two classes is not sharp; though when the properties of the resultant are sensibly the sum of those of the pure components, as is nearly the case for a complex gas such as air, it is usual to class it as a mixture. When the properties of the resultant substance are different from those of the components and it is not a chemical compound we define it as a solution.
Historical
Solutions were not distinguished from definite chemical compounds till John Dalton discovered the laws of definite and multiple proportions, but many earlier observations on the solubility of solids in water and the density of the resulting solutions had been made. As early as 1788 Sir Charles Blagden (1748-1820) made measurements of the freezing points of salt solutions, and showed that the depression of freezing point was. roughly proportional to the amount of salt dissolved. About 1850 Thomas Graham published his famous experiments on diffusion, both with and without a separating membrane. In 1867 botanical investigations by M. Traube, and in 1877 others by W. Pfeffer, made known the phenomena of the osmotic pressure which is set up by the passage of solvent through a membrane impermeable to the dissolved substance or solute. The importance of these experiments from the physical point of view was recognized by J. H. van't Hoff in 1885, who showed that Pfeffer's results indicated that osmotic pressure of a dilute solution conformed to the well-known laws of gas pressure, and had the same absolute value as the same number of molecules would exert as a gas filling a space equal to the value of the solvent. The conception of a semi-permeable membrane, permeable to the solvent only, was used by van't Hoff as a means of applying the principles of thermodynamics to the theory of solution.
Another method of applying the same principles is due to J. Willard Gibbs, who considered the whole problem of physical and chemical equilibrium in papers published in 1877, though the application of his principles only began to make extensive progress about twenty years after the publication of his purely theoretical investigations. The phenomena of solution and of vapour pressure constitute cases of equilibrium, and conform to the laws deduced by Gibbs, which thus yield a valuable method of investigating and classifying the equilibria of solutions.
Solubility
Some pairs of liquids are soluble in each other in all proportions, but, in general, when dealing with solutions of solids or gases in liquids, a definite limit is reached to the amount which will go into solution when the liquid is in contact with excess of the solid or gas. This limit depends on the nature of the two components, on the temperature and on the pressure. When the limit is reached the solution is said to be saturated, and the system is in equilibrium. If the solution of a solid more soluble when hot be cooled below the saturation point, the whole of the solid sometimes remains in solution. The liquid is then said to be supersaturated. But here the conditions are different owing to the absence of solid. If a crystal of the solid be added, the condition of supersaturation is destroyed, and the ordinary equilibrium of saturation is reached by precipitation of solid from solution.
The quantity of substance, or solute, which a given quantity of liquid or solvent will dissolve in presence of excess of the solute measures the solubility of the solute in the given solvent in the conditions of temperature and pressure. The solubilities of solids may be expressed in terms of the mass of solute which will dissolve in loo grammes of water.
The following may be taken as examples: - When dealing with gases it is usually more convenient to express the solubility as the ratio of the volume of the gas absorbed to the volume of the absorbing liquid. For gases such as oxygen and nitrogen dissolved in water the solubility as thus defined is independent of the pressure, or the mass of gas dissolved is proportional to the pressure. This relation does not hold for very soluble gases, such as ammonia, at low temperatures. As a general rule gases are less soluble at high than at low temperatures - unlike the majority of solids. Thus oxygen, 4.89 volumes of which dissolve at atmospheric pressure in I volume of water at o° C., only dissolves to the extent of 3Io volumes at 20° and 170 volumes at 100°.
Cause of Solubility
At the outset of the subject we are met by a fundamental problem, to which no complete answer can be given: Why do certain substances dissolve in certain other substances and not in different substances? Why are some pairs of liquids miscible in each other in all proportions, while other pairs do not mix at all, or only to a limited extent? No satisfactory correlation of solubility with chemical or other properties has been made. It is possible to state the conditions of solubility in terms of the theory of available energy, but the result comes to little more than a re-statement of the problem in other terms. Nevertheless, such a re-statement is in itself sometimes an advance in knowledge. It is certain then that when dissolution occurs the available energy of the whole system is decreased by the process, while when equilibrium is reached and the solution is saturated the available energy is a minimum. When a variable quantity is at a minimum a slight change in the system does not affect its value, and therefore, when a solution is saturated, the increase in the available energy of the liquid phase produced by dissolving in it some of the solid must be equal to the decrease in the available energy of the solid phase, caused by the abstraction from the bulk of that part dissolved. The general theory of such equilibria will be studied later under the head of the phase rule.
It is possible that a correlation may be made between solubility and the energy of surface tension. If a solid is immersed in a liquid a certain part of the energy of the system depends on, and is proportional to, the area of contact between solid and liquid. Similarly with two liquids like oil and water, which do not mix, we have surface energy proportional to the area of contact. Equilibrium requires that the available energy and therefore the area of contact should be a minimum, as is demonstrated in Plateau's beautiful experiment, where a large drop of oil is placed in a liquid of equal density and a perfect sphere is formed. If, however, the energy of surface tension between the two substances were negative the surface would tend to a maximum, and complete mixture would follow. From this point of view the natural solubility of two substances involves a negative energy of surface tension between them.
Gibbs's
A saturated solution is a system in equilibrium, and exhibits the thermodynamic relations which hold for all such systems. Just as two electrified bodies are in equilibrium when their electric potentials are equal, so two parts of a chemical and physical system are in equilibrium when there is equality between the chemical potentials of each component present in the two parts. Thus water and steam are in equilibrium with each other when the chemical potential of water substance is the same in the liquid as in the vapour. The chemical potentials are clearly functions of the composition of the system, and of its temperature and pressure. It is usual to call each part of the system of uniform composition throughout a phase; in the example given, water substance, the only component is present in two phases - a liquid phase and a vapour phase, and when the potentials of the component are the same in each phase equilibrium exists.
If in unit mass of any phase we have n components instead of one we must know the amount of n - I components present in that unit mass before we know the exact composition of it. Thus if in one gramme of a mixture of water, alcohol and salt we are told the amount of water and salt, we can tell the amount of alcohol. If, instead of one phase, we have r phases, we must find out the values of r(n - I) quantities before we know the composition of the whole system. Thus, to investigate the composition of the system we must be able to calculate the value of r (n-1) unknown quantities. To these must be added the external variables of temperature and pressure, and then as the total number of variables, we have r (n+I) + 2.
To determine these variables we may form equations between the chemical potentials of the different components - quantities which are functions of the variables to be determined. If µ i and µ2 denote the potentials of any one component in two phases in contact, when there is equilibrium, we know that µ i =P2If a third phase is in equilibrium with the other two we have also =123. These two equations involve the third relation µ2 =As, which therefore is not an independent equation. Hence with three phases we can form two independent equations for each component. With r phases we can form r - I equations for each component, and with n components and r phases we obtain n (r - 1) equations.
Now by elementary algebra we know that if the number of independent equations be equal to the number of unknown quantities all the unknown quantities can be determined, and can possess each one value only. Thus we shall be able to specify the system completely when the number of variables, viz. r (n - I) + 2, is equal to the number P of equations, viz. n(r - I); that is when r=n + 2. Thus, when a system possesses two more phases than the number of its components, all the phases will be in equilibrium with each other at one definite composition, one definite temperature and one definite pressure, and in no other conditions. To take the simplest case of a one component system water substance has its three phases of solid ice, liquid water and gaseous vapour in equilibrium with each other at the freezing point of water under the pressure of its own vapour. If we attempt to change either the temperature or the pressure ice will melt, water will evaporate or vapour condense until one or other of the phases has vanished. We then have in equilibrium two phases only, and the temperature and pressure may change. Thus, if we supply heat to the mixture of ice, water and steam ice will melt and eventually vanish. We then have water and vapour in equilibrium, and, as more heat enters, the temperature rises and the vapour-pressure rises with it. But, if we fix arbitrarily the temperature the pressure of equilibrium can have one value only. Thus by fixing one variable we fix the state of the whole system. This condition is represented in the algebraic theory when we have one more unknown quantity than the number of equations; i.e. when r(n-- t) + 2 = n (r - I) + I or r =n+I, and the number of phases is one more than the number of components. Similarly if we have F more unknowns than we have equations to determine them, we must fix arbitrarily F coordinates before we fix the state of the whole system. The number F is called the number of degrees of freedom of the system, and is measured by the excess of the number of unknowns over the number of variables. Thus F = r(n - I) + 2 - n(r - t) = n - r + 2, a result which was deduced by J. Willard Gibbs (1839-1903) and is known as Gibbs's Phase-Rule (see Energetics).
The phenomena of equilibrium can be represented on diagrams. Thus, if we take our co-ordinates to represent pressure and temperature, the state of the systems p with ice, water and vapour in equilibrium is represented by the point 0 where the pressure is that of the vapour of water at the freezing point and the temperature is the freezing point under that pressure. If all the ice be melted, we pass along the vapour pressure curve of water OA. If all the water be frozen, we have the vapour pressure curve of ice OB; while, if the pressure be raised, so that all the vapour vanishes, we get the curve OC of equilibrium between the pressure and the freezing point of water. The slope of these curves is determined by the so-called "latent heat equation" FIG. I.
(see Thermodynamics), dp/dt = X/t(v2 - v i), where p and t denote the pressure and temperature, X the heat required to change unit mass of the systems from one phase to the other, and v2 - v1 the resulting change in volume. The phase rule combined with the latent heat equation contains the whole theory of chemical and physical equilibrium.
Application to Solutions
In a system containing a solution we have to deal with two components at least. The simplest case is that of water and a salt, such as sodium chloride, which crystallizes without water. To obtain a non-variant system, we must assemble four phases - two more than the number of components. The four phases are (I) crystals of salt, (2) crystals of ice, (3) a saturated solution of the salt in water, and (4) the vapour, which is that practically of water alone, since the salt is non-volatile at the temperature in question. Equilibrium between these phases is obtained at the freezing point of the saturated solution under the pressure of the vapour. At that pressure and temperature the four phases can co-exist, and, as long as all of them are present, the pressure and temperature will remain steady. Thus a mixture of ice, salt and the saturated solution has a constant freezing point, and the composition of the solution is constant and the same as that of the mixed solids which freeze out on the abstraction of heat. This constancy both in freezing point and composition formerly was considered as a characteristic of a pure chemical compound, and hence these mixtures were described as components and given the name of "cryohydrates." In representing on a diagram the phenomena of equilibrium in a two-component system we require a third axis along which p to plot the composition of a variable phase. It is usual to take three axes at right angles to each other to represent pressure, temperature and the composition of the variable phase.. On a plane figure this solid diagram must be drawn in perspective, the third axis C being imagined to lie out of the plane of the paper. The phase-rule diagram that we FIG. 2. construct is then a sketch of a solid model, the lines of which do not really lie in the plane of the paper.
Let us return to the case of the system of salt and water. At the cryohydric point 0 we have four phases in equilibrium at a definite pressure, temperature and composition of the liquid phase. The condition of the system is represented by a single point on the diagram. If heat be added to the mixture ice will melt and salt dissolve in the water so formed. If the supply of ice fails first the temperature will rise, and, since solid salt remains, we pass along a curve OA giving the relation between temperature and the vapour pressure of the saturated solution. If, on the other hand, the salt of the cryohydrate fails before the ice the water given by the continued fusion dilutes the solution, and we pass along the curve OB which shows the freezing points of a series of solutions of constantly increasing dilution. If the process be continued till a very large quantity of ice be melted the resulting solution is so dilute that its freezing point B is identical with that of the pure solvent. Again, starting from 0, by the abstraction of heat we can remove all the liquid and travel along the curve OD of equilibrium between the two solids (salt and ice) and the vapour. Or, by increasing the pressure, we eliminate the vapour and obtain the curve OF giving the relation between pressure, freezing point and composition when a saturated solution is in contact with ice and salt.
If the salt crystallizes with a certain amount of water as well as with none, we get a second point of equilibrium between four phases. Sodium sulphate, for instance, crystallizes below 32.6° as Na 2 SO 4 IoH 2 O, and above that temperature as the anhydrous solid Na2S04. Taking the point 0 to denote the state of equilibrium between ice, hydrate; saturated solution and vapour, we pass along OA till a new solid phase, that of Na2S04, appears at 32.6°; from this point arise four curves, analogous to those diverging from the point O.
For the quantitative study of such systems in detail it is convenient to draw plane diagrams which are theoretically projections of the curves of the solid phase rule diagram on one or other of these planes. Experiments on the relation between temperature and concentration are illustrated by projecting the curve OA of fig. 2 on the tc-plane. The pressure at each point should be that of the vapour, but since the solubility of a solid does not change much with pressure, measurements under the constant atmospheric pressure give a curve practically identical with the theoretical one.
Fig. 3 gives the equilibrium between sodium sulphate and water in this way. B is the freezing point of pure water, 0 that 2345 FIG. 3.
of a saturated solution of Na2S04IoH 2 O. The curve OP represents the varying solubility of the hydrate as the temperature rises from the cryohydric point to 32.6°. At that temperature crystals of the anhydrous Na 2 SO 4 appear, and a new fixed equilibrium exists between the four phases - hydrate, anhydrous salt, solution and vapour. As heat is supplied, the hydrate is transformed gradually into the anhydrous salt and water. When this process is complete the temperature rises, and we pass along a new curve giving the equilibrium between anhydrous crystals, solution and vapour. In this particular case the solubility decreases with rise of temperature. This behaviour is exceptional.
Two Liquid Components
The more complete phenomena of mutual solubility are illustrated by the case of phenol and water. In fig. 4 A represents the freezing point of pure water, and AB the freezing point curve showing the depression of the freezing point as phenol is added. At B is a nonvariant system made up of ice, solid phenol, saturated solution and vapour. BCD c is the solubility curve of phenol in water. At C a new o° A C liquid phase appears - the B solution of water in liquid ' r q Ivater 50 Phenol phenol, the solubility of which FIG. 4.