Chemical elements
    Physical Properties
    Chemical Properties
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Element Sodium Na, Alkali Metal

About Sodium

The chemical relations of sodium are very similar to those of potassium, so that for chemical purposes the one metal can in most cases replace the other. This holds good especially for those reactions in which the ions come into account. The reason of this is that natrion also represents a far more stable state than metallic sodium, and the reactions of this element, as in the case of potassium, are therefore chiefly characterised by the fact that the ion is formed with especial readiness from the metal, but the metal only with difficulty from the ion. Since, further, the state of the salts in the solid form approaches more nearly to that of the ions than to that of the metal, sodium, like potassium, will be easily transformed from one of its salts into another, but will be converted only with difficulty from a salt into the metal or a compound closely related to this.

Metallic sodium does not occur in nature, since it would everywhere have an opportunity of exercising its tendency to pass into natrion. Natrion, however, has an extensive distribution, and, along with chloridion, with which it occurs in sea-water, it may be regarded as the most abundant ion in those parts of the earth's surface which are accessible to us.

In more remote times, the compounds of the two elements potassium and sodium were confused with one another. When it was learned how to distinguish them, caustic potash was known as the vegetable, and caustic soda as the mineral alkali, because the former was obtained chiefly from the ash of plants, the latter from common salt. It was later found by Klaproth that both elements are present in the mineral kingdom. So far as the vegetable kingdom is concerned, an essential difference does certainly exist between the two elements, for compounds of potassium must be present in considerable amount in plants in order that these may develop normally. Sodium compounds, it is true, are never wanting in plants, but they are more chance constituents which pass into the plants from the soil, in which they are always present, and seem not to play any particular part in them.

Although, therefore, normal vegetation may be hindered by an entire exclusion of sodium compounds (although no indubitable evidence on this point exists), it is certain that the quantities of sodium which may possibly be necessary for a plant are incomparably smaller than the amounts of potassium which are indispensable.

The cause of this difference may be found in the following circumstance. Whereas the soil in which plants thrive has the remarkable property of withdrawing dissolved potassium compounds from solution, and retaining them in such a way that they can be taken up only in a very slight degree by water, the behaviour is quite different with respect to the sodium compounds. These are not retained by the soil, but filter through without difficulty. Whereas, therefore, the amount of potassium compounds in the soil is considerable and almost independent of chance conditions, the amount of the sodium compounds is subject to variation and to chance. On the principle of the survival of the fittest, it is intelligible that the chemical requirements of the plants, the satisfaction of which is effected by an alkali metal (or its ion), should be supplied by the constantly present potassium, since organisms whose life depended on utilising sodium compounds would die out by reason of the readily occurring lack of these.

The accumulation of natrion in sea-water is due to the same cause. When in the decomposition of the rocks by water and carbonic acid, the alkali metals pass into solution in the form of their ions, they follow, in the first instance, the general movement of the water towards the ocean. The potassium, however, is mostly retained on the way, because it is seized hold of by the soil; the sodium, however, passes on unhindered to the sea, and is deposited again in the solid form only in rare cases, viz. when the sea-water is concentrated by evaporation until the solid salt forms.

Cases of this have occurred, especially in former geological periods, and have given rise to beds of rock-salt or sodium chloride, the two ions which are present in greatest abundance in sea-water having been deposited together as solid salt.

Metallic Sodium

We have already, on several occasions, become acquainted with metallic sodium as a silver white, soft, and readily fusible metal, which reacts energetically in contact with water, and just as readily forms compounds with many other substances. It behaves, in general, quite similarly to potassium, from which it is distinguished by the somewhat inferior violence of its reactions.

Thus, sodium does not take fire when thrown on water; it does so, however, if its motion and the cooling which is thereby effected is prevented. This happens when the metal is placed on wet paper or on an aqueous jelly of glue or of starch. The evolved hydrogen as well as a portion of the metal then burns with a bright yellow flame, and all the flames in a room in which such a combustion has occurred burn distinctly yellow for a considerable time. This is due to the fact that the dissipated sodium compounds colour the flames yellow, even when present in the minutest quantities.

Sodium melts at 97.5°, and boils at about 740°. The accurate determination of its vapour density is difficult, but the experiments which have been made agree in showing that the molar weight of sodium vapour is 23, or equal to the combining weight. This identity is a general property of the metals, so far as these are known in the vapour form, and less doubtful cases of the confirmation of this rule will be given later.

By mixture with other metals, the melting point of sodium is lowered. This is especially well seen on adding potassium; in this way alloys can easily be obtained which are liquid at the room temperature.

This phenomenon is by no means to be explained as a consequence of chemical combination between the metals; on the contrary, it is the simple consequence of the perfectly general fact that the melting point of every substance is lowered by the addition of such substances as are soluble in the liquid form of the first substance. If the melting point of the pure substance is not too high above room temperature, it may be lowered to below this temperature, and the phenomenon in question makes its appearance.

The first preparation of metallic sodium was effected by Davy by means of the voltaic pile at the same time as that of potassium. Shortly afterwards the method of obtaining it by distillation of sodium carbonate with charcoal was discovered, corresponding to the method mentioned under potassium. Since about 1860, sodium has been obtained on the large scale by this method, the metal being used for preparing aluminium. Recently, however, the electrical method has again been adopted, and sodium is obtained by the decomposition of sodium hydroxide by means of the electric current. By reason of the comparatively small cost of electrical energy, combined with the good yield obtained, sodium can be obtained more cheaply by this than by the old method. It is remarkable that this method is identical with that by which sodium was first prepared, for on that occasion also, sodium hydroxide was the original substance.

The electrolysis is carried out in iron pots divided by permeable partitions. At the anode oxygen escapes, at the cathode sodium and hydrogen are formed. The separated metal is lighter than the liquid hydroxide, and therefore floats to the surface; it is skimmed off from time to time.

Sodium can also be obtained by the electrolysis of fused sodium chloride. Much difficulty, however, is caused by the high melting point of this salt. The melting point can be lowered by mixing it with potassium chloride; mixtures of sodium with a little potassium are then obtained, but not the pure metal. For this reason, successful attempts have also been made to employ the readily fusible sodium nitrate for the electrolysis.

Metallic sodium is largely used in the arts and in the laboratory. Its former importance for obtaining other difficultly reducible metals has been lost, since the object can generally be attained more readily by means of magnesium or aluminium, or by the electrolytic method. It is used, however, as a powerful reducing agent in many reactions in organic chemistry, and for obtaining reactive intermediate products.

For these purposes, the metal is best employed in a condition in which it offers a large surface. Since, on account of the softness of the metal, it cannot be reduced to small pieces by blows or by filing, it is forced, by means of an iron screw press, through narrow openings, and is thus obtained in the form of wire or of ribbon, according to the shape of the opening. Since in this state the metal very rapidly oxidises in the air, the wire is allowed to fall directly into the liquid on which it is to act, or it is collected in a liquid which does not contain oxygen. Petroleum, which is usually employed for this purpose, has the disadvantage that it is difficult to remove; for chemical purposes, therefore, it is better to use readily volatile hydrocarbons obtained from the low-boiling portions of petroleum (so-called petroleum benzine or petroleum ether).

Sodium Occurrence

Although sodium in the free state is not found in nature, it is present in combination in most minerals. Soda-felspar or albite is a double silicate of sodium and aluminium, 3Na2O,Al2O3,6SiO2. Sea-water contains 2.6 to 2.9 per cent, of sodium chloride, NaCl, the deposits left by the evaporation of inland seas being known as rock-salt Both the carbonates and the sulphate of sodium occur dissolved in the water of many mineral springs, while the sulphate is a constituent of certain double salts, such as glauberite or sodium calcium sulphate, and blodite or sodium magnesium sulphate. Great deposits of Chile saltpetre or sodium nitrate, NaNO3, are present in Chile. Cryolite or sodium aluminium fluoride, 3NaF,AlF3, is an important mineral found in Greenland. Sodium carbonate occurs in South America and Egypt, and is also found as gaylussite, a double carbonate of sodium and calcium. Tincal or disodium tetraborate, Na2B4O7,10H2O, is native to Thibet, India, and California, and a double borate of sodium and calcium called cryptomorphite is also found.

Sodium History

The knowledge of sodium carbonate or "soda" is of great antiquity, as indicated by two references in the Bible. The word translated " nitre " in the Authorized Version means " natron " or " soda," and is correctly rendered "lye" in Jer. II. 22 in the Revised Version. No alteration has been made in the other reference. The confusion of terms evidently originated in the resemblance between the Greek vcrpov employed by Dioscorides and the Latin nitrum used by Pliny to denote sodium carbonate, and the word " nitre," loosely employed in early English as synonymous with "natron" or "soda," but now reserved for potassium nitrate.

In the sixteenth century Biringuccio seems to have appreciated the distinction between "nitrum" or soda and "sal nitri" or saltpetre. Somewhat earlier the Arabs introduced into Europe the words "natrun," "natrum," and "natron," signifying soda, and " nitrum," meaning saltpetre. They also introduced the word "alkali", but drew no distinction between soda, derived from the ashes of sea-plants, and potash, obtained from the ashes of land-plants. These substances were denominated " fixed alkali" in contradistinction to the volatile ammonium carbonate.

In 1736 Duhamel de Monceau noted the difference between the "mineral alkali" or soda obtained from rock-salt and the "vegetable alkali" or potash extracted from plant-ashes. Marggraf recorded the difference in flame-coloration produced by the two substances, and in 1796 Klaproth discovered the presence of "vegetable alkali" in the mineral world as leucite. At an earlier date it was noticed that the so-called "mild alkali" or sodium carbonate is converted into "caustic alkali," or sodium hydroxide (in modern parlance), by the action of slaked lime, and in 1756 Black proved the presence of "fixed air" (carbon dioxide) in the "mild alkali."

In 1807 Davy isolated the alkali-metals by electrolysis of the fused hydroxides, thus proving the invalidity of Lavoisier's conception of the oxides as elementary substances.


Chemical Elements


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