Chemical elements
    Physical Properties
    Chemical Properties
      Formation of Water
      Purification of Water
      Hardness of Water
      Softening of Hard Water
      Permutit Process
      Sterilisation of Water
      Physical Properties of Ice
      Physical Properties of Water
      Physical Properties of Gaseous Water
      Chemical Properties of Water
        Water Decomposition
        Water as a Catalyst
        Dissociation of Steam
        Steam as Oxidising Agent
        Water Detection and Estimation
      Solubility of Gases in Water
      Solubility of Liquids in Water
      Solubility of Solids in Water
      Supersaturated Solutions
      Combined Water
      Water Analysis
    Hydrogen peroxide

Chemical Properties of Water

Chemical Properties of Water are very important for understanding. Water is readily decomposed at the ordinary temperature by alkali and alkaline earth metals in compact form. The temperatures at which water, in the form of ice, becomes measurably attacked by the alkali metals have been determined as follow:

Sodium-98° C
Potassium-105° C
Rubidium-108° C
Cesium-116° C.

Many other metals liberate the hydrogen on warming, particularly when in a finely divided condition. Thus pyrophoric iron rapidly decomposes water at 50° to 60° C., and its action is perceptible even below 10° C. Its reactivity appears to be independent of the presence of occluded gases or of carbon, and to be solely dependent upon its fine state of division. Boiling water is slowly decomposed by granulated lead.

Addition of magnesium powder to ten times its own weight of cold water, followed by a little palladous chloride, causes the evolution of hydrogen, which spontaneously ignites.

Although aluminium is not readily attacked by water at the ordinary temperature, in contact with iodine the hydroxide is formed, hydrogen being liberated. This appears to be due to the formation of a little aluminium iodide, AlI3, which is immediately hydrolysed to the hydroxide and hydrogen iodide. This latter then attacks the aluminium, liberating hydrogen, and yielding a further quantity of iodide, which immediately in its turn undergoes hydrolysis. Since the iodine does not enter into the final products, a very small quantity is sufficient to effect the oxidation of an indefinitely large quantity of aluminium; in other words, it is a catalyst.

Mention has already been made of the fact that the action of water is of considerable value in discriminating between acidic and basic oxides. A somewhat similar series of reactions takes place with chlorides. Thus the acid chlorides PCl5, PCl3, SiCl4, AsCl3, are converted into free hydrochloric acid, and the corresponding acid derived from the non-metal. In the case of chlorides derived from organic acids, analogous results obtain. Thus acetyl chloride, CH3.COCl, yields acetic acid, CH3.COOH, together with hydrochloric acid.

Metallic carbides are frequently decomposed by water, yielding hydrocarbons. One of the best known of these reactions is that with calcium carbide, which yields acetylene. Thus

CaC2 + 2H2O = Ca(OH)2 + C2H2.

Even combined water or " water of crystallisation " may induce this reaction. Sodium carbonate, Na2CO3.10H2O, is a useful salt to employ, and the reaction takes place at a more moderate temperature.

By the action of water on aluminium carbide, methane is obtained admixed, however, with a little hydrogen.

Al4C3 + 12H2O = 4Al(OH)3 + 3CH4.

Other carbides, such as those of thorium, uranium, and glucinum, likewise yield methane, but mixed with various hydrocarbons.

This behaviour of metallic carbides led Mendeleeff tentatively to suggest that the large natural reservoirs of petroleum in America have been formed by the action of water or steam on subterranean metallic carbides.

Phosphides and silicides frequently behave in an analogous manner. Thus, calcium phosphide, Ca2P2, is decomposed by water yielding phosphorus trihydride, PH3, and liquid phosphoretted hydrogen, P2H4, which is spontaneously inflammable.

3Ca2P2 + 12H2O = 6Ca(OH)2 + 4PH3 + 2P;
Ca2P2 + 4H2O = 2Ca(OH)2 + P2H4.

Even yellow phosphorus itself, when warmed with water, yields hydrogen phosphide.

Metallic nitrides and hydrides are decomposed by water either in the cold or on warming, yielding respectively ammonia and hydrogen.

Many organo-metallic derivatives are decomposed by water. Thus zinc methyl yields methane:

Zn(CH3)2 + 2H2O - Zn(OH)2 + 2CH4. Magnesium methyl iodide, Mg(CH3)I, behaves in an analogous manner: Mg(CH3)I + H2O = CH4 + Mg(OH)I.

Some metallic peroxides, such as sodium peroxide, are decomposed by water. An intimate mixture of powdered aluminium and sodium peroxide inflames when brought into contact with water.

A few metallic sulphides are decomposed by water. The majority, however, are stable in the presence of water, and this fact is made use of in routine methods of qualitative analysis. Water decomposes strontium sulphide, yielding a mixture of hydroxide and liydrosulphide, which can be readily separated on account of their widely differing solubilities, the latter substance being the more soluble:

2SrS + 2H2O = Sr(SH)2 + Sr(OH)2.

Hence, by extracting strontium sulphide with hot water and cooling the clear filtrate, pure crystalline strontium hydroxide is obtained.

With barium sulphide the reactions are more complex, and pure barium hydroxide cannot be obtained in the above manner.

When boiled with sulphur in the presence of oxygen in platinum vessels, water yields hydrogen sulphide and sulphuric acid.

Many salts are decomposed by water, particularly when their solutions are boiled, basic salts being produced.

Bismuth chloride almost immediately undergoes such " hydrolysis to the basic bismuthyl chloride,

BiCl3 + H2OBiOCl + 2HCl,

this reaction affording a convenient method of separating bismuth salts quantitatively from certain others. Solutions of ferric chloride, and indeed of many salts composed of a strong acid united with a feeble base, are acid in reaction from a similar cause:

FeCl3 + 3H2OFe(OH)3 + 3HCl.

For analogous reasons, salts containing feebly acidic radicles with strongly basic ones generally give alkaline solutions, e.g. potassium cyanide, sodium carbonate, etc.:

Na2CO3 + H2ONaHCO3 + NaOH.

The action is referable to the effect of the ionisation of the water; in the presence of one of its salts a weak acid such as hydrocyanic is dissociated to an extent so slight as to be comparable in dissociation with water itself. Under such conditions an appreciable competition will occur between the acid of the salt and the water for possession of the metallic radicle.

Chemical Properties of Water with ethereal salts or esters derived from the neutralisation of an acid by an alcohol also undergo hydrolysis by water and the fats, which are compounds of this type, are frequently decomposed in this way for the manufacture of glycerine, candles, or soap; the hydrolysis of such esters is catalytically accelerated by the addition of a mineral acid.

When an air-free solution of potassium cobalto cyanide, K4CO(CN)6, is boiled, hydrogen is evolved, the volume of which equals that of the oxygen absorbed if the solution is rapidly oxidised in air, but to twice the volume absorbed during slow oxidation in air. The excess of oxygen in the former case remains at the close of the reaction as hydrogen peroxide. Thus, in air absence,

2K4Co(CN)6 + 2H2O - 2K3Co(CN)6 + 2KOH + H2;

with rapid oxidation -

2K4Co(CN)6 + 2H2O + O2 = 2K3Co(CN)6 + 2KOH + H2O2;

but with slow oxidation -

2K4Co(CN)6 + H2O + O = 2K3Co(CN)6 + 2KOH.

The influence of water upon the direction of certain reactions, in consequencc of the heat liberated by the solution of one product, is beautifully illustrated in the case of sulphur, iodine, and their hydrides. The heats of reaction of hydrogen and solid iodine and sulphur in the dry are as follow:

(H2) + [I2] = 2(HI) - 12,072 calories
(H2) + [S] = (H2S) + 2,730 calories

hence sulphur will effect the decomposition of hydrogen iodide with marked heat evolution. On the other hand, if the reactions take place in the presence of liquid water, there is a considerable evolution of heat in both cases in consequence of the heats of solution of the hydrides. Thus

(H2) + [I2] + Aq. = 2HI.Aq. + 26,348 calories.
(H2) + [S] + Aq. = H2S. Aq. + 7,290 calories

Owing to the greater heat of solution of hydrogen iodide, the relative heat evolutions are now actually reversed, and solid iodine, in its turn, can decompose aqueous hydrogen sulphide, the reaction being markedly exothermic.

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