Rabu, 01 Juni 2011


Occurrence. As has been stated in a former chapter, nitrogen constitutes a large fraction of the atmosphere. The compounds of nitrogen, however, cannot readily be obtained from this source, since at any ordinary temperature nitrogen is able to combine directly with very few of the elements.

In certain forms of combination nitrogen occurs in the soil from which it is taken up by plants and built into complex substances composed chiefly of carbon, hydrogen, oxygen, and nitrogen. Animals feeding on these plants assimilate the nitrogenous matter, so that this element is an essential constituent of both plants and animals.

Decomposition of organic matter by bacteria. When living matter dies and undergoes decay complicated chemical reactions take place, one result of which is that the nitrogen of the organic matter is set free either as the element nitrogen, or in the form of simple compounds, such as ammonia (NH3) or oxides of nitrogen. Experiment has shown that all such processes of decay are due to the action of different kinds of bacteria, each particular kind effecting a different change.

Decomposition of organic matter by heat. When organic matter is strongly heated decomposition into simpler substances takes place in much the same way as in the case of bacterial decomposition. Coal is a complex substance of vegetable origin, consisting largely of carbon, but also containing hydrogen, oxygen, and nitrogen. When this is heated in a closed vessel so that air is excluded, about one seventh of the nitrogen is converted into ammonia, and this is the chief source from which ammonia and its compounds are obtained.


Ammonia (NH3). Several compounds consisting exclusively of nitrogen and hydrogen are known, but only one, ammonia, need be considered here.

Preparation of ammonia. Ammonia is prepared in the laboratory by a different method from the one which is used commercially.

1. Laboratory method. In the laboratory ammonia is prepared from ammonium chloride, a compound having the formula NH4Cl, and obtained in the manufacture of coal gas. As will be shown later in the chapter, the group NH4 in this compound acts as a univalent radical and is known as ammonium. When ammonium chloride is warmed with sodium hydroxide, the ammonium and sodium change places, the reaction being expressed in the following equation.

NH4Cl + NaOH = NaCl + NH4OH. 

The ammonium hydroxide (NH4OH) so formed is unstable and breaks down into water and ammonia.
NH4OH = NH3 + H2O. 

Calcium hydroxide (Ca(OH)2) is frequently used in place of the more expensive sodium hydroxide, the equations being

2NH4Cl + Ca(OH)2 = CaCl2 + 2NH4OH,
2NH4OH = 2H2O + 2NH3

In the preparation, the ammonium chloride and calcium hydroxide are mixed together and placed in a flask arranged as shown in Fig. 35. The mixture is gently warmed, when ammonia is evolved as a gas and is collected by displacement of air.

2. Commercial method. Nearly all the ammonia of commerce comes from the gasworks. Ordinary illuminating gas is made by distilling coal, as will be explained later, and among the products of this distillation a solution of ammonia in water is obtained. This solution, known as gas liquor, contains not only ammonia but other soluble substances. Most of these combine chemically with lime, while ammonia does not; if then lime is added to the gas liquor and the liquor is heated, the ammonia is driven out from the mixture. It may be dissolved again in pure, cold water, forming aqua ammonia, or the ammonia water of commerce.

Preparation from hydrogen and nitrogen. When electric sparks are passed for some time through a mixture of hydrogen and nitrogen, a small percentage of the two elements in the mixture is changed into ammonia. The action soon ceases, however, for the reason that ammonia is decomposed by the electric discharge. The reaction expressed in the equation
N + 3H = NH3
can therefore go in either direction depending upon the relative quantities of the substances present. This recalls the similar change from oxygen into ozone, which soon ceases because the ozone is in turn decomposed into oxygen.

Physical properties. Under ordinary conditions ammonia is a gas whose density is 0.59. It is therefore little more than half as heavy as air. It is easily condensed into a colorless liquid, and can now be purchased in liquid form in steel cylinders. The gas is colorless and has a strong, suffocating odor. It is extremely soluble in water, 1 l. of water at 0° and 760 mm. pressure dissolving 1148 l. of the gas. In dissolving this large volume of gas the water expands considerably, so that the density of the solution is less than that of water, the strongest solutions having a density of 0.88.

Chemical properties. Ammonia will not support combustion, nor will it burn under ordinary conditions. In an atmosphere of oxygen it burns with a feeble, yellowish flame. When quite dry it is not a very active substance, but when moist it combines with a great many substances, particularly with acids.

Uses. It has been stated that ammonia can be condensed to a liquid by the application of pressure. If the pressure is removed from the liquid so obtained, it rapidly passes again into the gaseous state and in so doing absorbs a large amount of heat. Advantage is taken of this fact in the preparation of artificial ice. Large quantities of ammonia are also used in the preparation of ammonium compounds.

The manufacture of artificial ice. Fig. 36 illustrates the method of preparing artificial ice. The ammonia gas is liquefied in the pipes X by means of the pump Y. The heat generated is absorbed by water flowing over the pipes. The pipes lead into a large brine tank, a cross section of which is shown in the figure. Into the brine (concentrated solution of common salt) contained in this tank are dipped the vessels A, B, C, filled with pure water. The pressure is removed from the liquid ammonia as it passes into the pipes immersed in the[Pg 126] brine, and the heat absorbed by the rapid evaporation of the liquid lowers the temperature of the brine below zero. The water in A, B, C is thereby frozen into cakes of ice. The gaseous ammonia resulting from the evaporation of the liquid ammonia is again condensed, so that the process is continuous.

Ammonium hydroxide (NH4OH). The solution of ammonia in water is found to have strong basic properties and therefore contains hydroxyl ions. It turns red litmus blue; it has a soapy feel; it neutralizes acids, forming salts with them. It seems probable, therefore, that when ammonia dissolves in water it combines chemically with it according to the equation

NH3 + H2O = NH4OH, 

and that it is the substance NH4OH, called ammonium hydroxide, which has the basic properties, dissociating into the ions NH4 and OH. Ammonium hydroxide has never been obtained in a pure state. At every attempt to isolate it the substance breaks up into water and ammonia,—
NH4OH = NH3 + H2O. 

The ammonium radical. The radical NH4 plays the part of a metal in many chemical reactions and is called ammonium. The ending -ium is given to the name to indicate the metallic properties of the substance, since the names[Pg 127] of the metals in general have that ending. The salts formed by the action of the base ammonium hydroxide on acids are called ammonium salts. Thus, with hydrochloric acid, ammonium chloride is formed in accordance with the equation
NH4OH + HCl = NH4Cl + H2O.
Similarly, with nitric acid, ammonium nitrate (NH4NO3) is formed, and with sulphuric acid, ammonium sulphate ((NH4)2S04).

It will be noticed that in the neutralization of ammonium hydroxide by acids the group NH4 replaces one hydrogen atom of the acid, just as sodium does. The group therefore acts as a univalent metal.

Combination of nitrogen with hydrogen by volume. Under suitable conditions ammonia can be decomposed into nitrogen and hydrogen by passing electric sparks through the gas. Accurate measurement has shown that when ammonia is decomposed, two volumes of the gas yield one volume of nitrogen and three volumes of hydrogen. Consequently, if the two elements were to combine directly, one volume of nitrogen would combine with three volumes of hydrogen to form two volumes of ammonia. Here, as in the formation of steam from hydrogen and oxygen, small whole numbers serve to indicate the relation between the volumes of combining gases and that of the gaseous product.


In addition to ammonium hydroxide, nitrogen forms several compounds with hydrogen and oxygen, of which nitric acid (HNO3) and nitrous acid [Pg 128](HNO2) are the most familiar.

Nitric acid (HNO3). Nitric acid is not found to any extent in nature, but some of its salts, especially sodium nitrate (NaNO3) and potassium nitrate (KNO3) are found in large quantities. From these salts nitric acid can be obtained.

Preparation of nitric acid. When sodium nitrate is treated with concentrated cold sulphuric acid, no chemical action seems to take place. If, however, the mixture is heated in a retort, nitric acid is given off as a vapor and may be easily condensed to a liquid by passing the vapor into a tube surrounded by cold water, as shown in Fig. 37. An examination of the liquid left in the retort shows that it contains sodium acid sulphate (NaHSO4), so that the reaction may be represented by the equation

NaNO3 + H2SO4 = NaHSO4 + HNO3

If a smaller quantity of sulphuric acid is taken and the mixture is heated to a high temperature, normal sodium sulphate is formed:

2NaNO3 + H2SO4 = Na2SO4 + 2HNO3.

In this case, however, the higher temperature required decomposes a part of the nitric acid.

The commercial preparation of nitric acid. Fig. 38 illustrates a form of apparatus used in the preparation of nitric acid on a large scale. Sodium nitrate and sulphuric acid are heated in the iron retort A. The resulting acid vapors pass in the direction indicated by the arrows, and are condensed in the glass tubes B, which are covered with cloth kept cool by streams of water. These tubes are inclined so that the liquid resulting from the condensation of the vapors runs back into C and is drawn off into large vessels (D).

Physical properties of nitric acid. Pure nitric acid is a colorless liquid, which boils at about 86° and has a density of 1.56. The concentrated acid of commerce contains about 68% of the acid, the remainder being water. Such a mixture has a density of 1.4. The concentrated acid fumes somewhat in moist air, and has a sharp choking odor.

Chemical properties. The most important chemical properties of nitric acid are the following.

1. Acid properties. As the name indicates, this substance is an acid, and has all the properties of that class of substances. It changes blue litmus red and has a sour taste in dilute solutions. It forms hydrogen ions in solution and neutralizes bases forming salts. It also acts upon the oxides of most metals, forming a salt and water. It is one of the strongest acids.

2. Decomposition on heating. When boiled, or exposed for some time to sunlight, it suffers a partial decomposition according to the equation
2HNO3 = H2O + 2NO2 + O.
The substance NO2, called nitrogen peroxide, is a brownish gas, which is readily soluble in water and in nitric acid. It therefore dissolves in the undecomposed acid, and imparts a yellowish or reddish color to it. Concentrated[Pg 130] nitric acid highly charged with this substance is called fuming nitric acid.

3. Oxidizing action. According to its formula, nitric acid contains a large percentage of oxygen, and the reaction just mentioned shows that the compound is not a very stable one, easily undergoing decomposition. These properties should make it a good oxidizing agent, and we find that this is the case. Under ordinary circumstances, when acting as an oxidizing agent, it is decomposed according to the equation

2HNO3 = H2O + 2NO + 3O. 

The oxygen is taken up by the substance oxidized, and not set free, as is indicated in the equation. Thus, if carbon is oxidized by nitric acid, the oxygen combines with carbon, forming carbon dioxide (CO2):

C + 2O = CO2

4. Action on metals. We have seen that when an acid acts upon a metal hydrogen is set free. Accordingly, when nitric acid acts upon a metal, such as copper, we should expect the reaction to take place which is expressed in the equation

Cu + 2HNO3 = Cu(NO3)2 + 2H. 

This reaction does take place, but the hydrogen set free is immediately oxidized to water by another portion of the nitric acid according to the equation

HNO3 + 3H = 2H2O + NO. 

As these two equations are written, two atoms of hydrogen are given off in the first equation, while three are used up in the second. In order that the hydrogen may be equal in[Pg 131] the two equations, we must multiply the first by 3 and the second by 2. We shall then have

3Cu + 6HNO3 = 3Cu(NO3)2 + 6H, 

2HNO3 + 6H = 4H2O + 2NO. 

The two equations may now be combined into one by adding the quantities on each side of the equality sign, canceling the hydrogen which is given off in the one reaction and used up in the other. We shall then have the equation

3Cu + 8HNO3 = 3Cu(NO3)2 + 2NO + 4H2O. 

A number of other reactions may take place when nitric acid acts upon metals, resulting in the formation of other oxides of nitrogen, free nitrogen, or even ammonia. The reaction just given is, however, the usual one.

Importance of steps in a reaction. This complete equation has the advantage of making it possible to calculate very easily the proportions in which the various substances enter into the reaction or are formed in it. It is unsatisfactory in that it does not give full information about the way in which the reaction takes place. For example, it does not suggest that hydrogen is at first formed, and subsequently transformed into water. It is always much more important to remember the steps in a chemical reaction than to remember the equation expressing the complete action; for if these steps in the reaction are understood, the complete equation is easily obtained in the manner just described.

Salts of nitric acid,—nitrates. The salts of nitric acid are called nitrates. Many of these salts will be described in the study of the metals. They are all soluble in water, and when heated to a high temperature undergo decomposition. In a few cases a nitrate on being heated evolves oxygen, forming a nitrite:
NaNO3 = NaNO2 + O.
In other cases the decomposition goes further, and the metal is left as oxide:
Cu(NO3)2 = CuO + 2NO2 + O. 

Nitrous acid (HNO2). It is an easy matter to obtain sodium nitrite (NaNO2), as the reaction given on the previous page indicates. Instead of merely heating the nitrate, it is better to heat it together with a mild reducing agent, such as lead, when the reaction takes place which is expressed by the equation
NaNO3 + Pb = PbO + NaNO2.
When sodium nitrite is treated with an acid, such as sulphuric acid, it is decomposed and nitrous acid is set free:
NaNO2 + H2SO4 = NaHSO4 + HNO2.
The acid is very unstable, however, and decomposes readily into water and nitrogen trioxide (N2O3):
2HNO2 = H2O + N2O3.
Dilute solutions of the acid, however, can be obtained.


Nitrogen combines with oxygen to form five different oxides. The formulas and names of these are as follows:
nitrous oxide.
nitric oxide.
nitrogen peroxide.
nitrogen trioxide, or nitrous anhydride.
nitrogen pentoxide, or nitric anhydride.
These will now be briefly discussed.

Nitrous oxide (laughing gas) (N2O). Ammonium nitrate, like all nitrates, undergoes decomposition when heated; and owing to the fact that it contains no metal, but does[Pg 133] contain both oxygen and hydrogen, the reaction is a peculiar one. It is represented by the equation
NH4NO3 = 2H2O + N2O. 

The oxide of nitrogen so formed is called nitrous oxide or laughing gas. It is a colorless gas having a slight odor. It is somewhat soluble in water, and in solution has a slightly sweetish taste. It is easily converted into a liquid and can be purchased in this form. When inhaled it produces a kind of hysteria (hence the name "laughing gas"), and even unconsciousness and insensibility to pain if taken in large amounts. It has long been used as an anæsthetic for minor surgical operations, such as those of dentistry, but owing to its unpleasant after effects it is not so much in use now as formerly.

Chemically, nitrous oxide is remarkable for the fact that it is a very energetic oxidizing agent. Substances such as carbon, sulphur, iron, and phosphorus burn in it almost as brilliantly as in oxygen, forming oxides and setting free nitrogen. Evidently the oxygen in nitrous oxide cannot be held in very firm combination by the nitrogen.
Nitric oxide (NO). We have seen that when nitric acid acts upon metals, such as copper, the reaction represented by the following equation takes place:
3Cu + 8HNO3 = 3Cu(NO3)3 + 2NO + 4H2O.
Nitric oxide is most conveniently prepared in this way. The metal is placed in the flask A (Fig. 39) and the acid added slowly through the funnel tube B. The gas escapes through C and is collected over water.[Pg 134]
Pure nitric oxide is a colorless gas, slightly heavier than air, and is practically insoluble in water. It is a difficult gas to liquefy. Unlike nitrous oxide, nitric oxide does not part with its oxygen easily, and burning substances introduced into this gas are usually extinguished. A few substances like phosphorus, which have a very strong affinity for oxygen and which are burning energetically in the air, will continue to burn in an atmosphere of nitric oxide. In this case the nitric oxide loses all of its oxygen and the nitrogen is set free as gas.

Action of nitric oxide with oxygen. When nitric oxide comes into contact with oxygen or with the air, it at once combines with the oxygen even at ordinary temperatures, forming a reddish-yellow gas of the formula NO2, which is called nitrogen peroxide. This action is not energetic enough to produce a flame, though considerable heat is set free.

Nitrogen peroxide (NO2). This gas, as we have just seen, is formed by allowing nitric oxide to come into contact with oxygen. It can also be made by heating certain nitrates, such as lead nitrate:
Pb(NO3)2 = PbO + 2NO2 + O. 

It is a reddish-yellow gas of unpleasant odor, which is quite poisonous when inhaled. It is heavier than air and is easily condensed to a liquid. It dissolves in water, but this solution is not a mere physical solution; the nitrogen peroxide is decomposed, forming a mixture of nitric and nitrous acids:
2NO2 + H2O = HNO2 + HNO3

Nitrogen peroxide will not combine with more oxygen; it will, however, give up a part of its oxygen to burning substances, acting as an oxidizing agent:
NO2 = NO + O. 

Acid anhydrides. The oxides N2O3 (nitrogen trioxide) and N2O5 (nitrogen pentoxide) are rarely prepared and need not be separately described. They bear a very interesting relation to the acids of nitrogen. When dissolved in water they combine with the water, forming acids:

N2O3 + H2O = 2HNO2,
N2O5 + H2O = 2HNO3

On the other hand, nitrous acid very easily decomposes, yielding water and nitrogen trioxide, and by suitable means nitric acid likewise may be decomposed into water and nitrogen pentoxide:

2HNO2 = H2O + N2O3,
2HNO3 = H2O + N2O5

In view of the close relation between these oxides and the corresponding acids, they are called anhydrides of the acids, N2O3 being nitrous anhydride and N2O5 nitric anhydride.

DEFINITION: Any oxide which will combine with water to form an acid, or which together with water is formed by the decomposition of an acid, is called an anhydride of that acid.

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