The preparation of nitrogen may be also regarded as a displacement by means of oxygen. Ammonia burns In oxygen, thus: 3NH3 + $O2= 3H2O + N2, but at the same time some of the nitrogen unites with the oxygen and forms NO2, nitric peroxide: this gas interacts with the ammonia, forming ammonium nitrate and nitrite, NH4NO3 and NH4NO2. If, however, the oxygen be not free, but in combination with an easily reduced metal, such as copper, it will combine with the hydrogen of the ammonia at a red heat, setting free the nitrogen. Another method involves the mutual displacement of nitrogen from its oxide by means of hydrogen, and from its hydride, ammonia, by oxygen: 2NH3 + N2O3 = 3H2O + 2N2. This method is, however, usually represented by the equation NH4NO2 = 2H2O + N2; for ammonium nitrite, NH4NO2, may be regarded as a compound of N2O3 with 2NH3 and H2O. To obtain nitrogen by this method, since ammonium nitrite is not easily obtained, a solution of ammonium chloride may be warmed with one of sodium nitrite. The equation is then: NaNO2. Aq + NH4C1. Aq = 2H2O + N2 + NaCl. Aq. Another convenient method is to warm together solutions of sodium hypobromite and ammonium chloride ; the former loses oxygen readily, which combines with the hydrogen of the ammonia according to the equation: 3NaOBr.Aq. + 2NH4Cl.Aq. = 3NaBr.Aq. + 3H2O + 2HCl.Aq. + N2.
Although sulphur, selenium, and tellurium burn in oxygen, still they may be displaced from their hydrides, H28, H2Se, and H2Te, by means of oxygen at a red heat, provided the oxygen is present only in sufficient quantity to combine with the hydrogen, thus : 2H2S + O2-2H2O + S2. Aqueous solutions of these compounds, too, are decomposed on standing in contact with air, owing to similar displacement. Oxygen may displace mercury from its sulphide, cinnabar, HgS, which is the common ore of mercury; here the sulphide is roasted in air, when the sulphur combines with the oxygen to form sulphur dioxide, a gas at ordinary temperature; and mercury is liberated, also in the gaseous form, but condensing at temperatures below 3580.
Fluorine, chlorine, and bromine may also be employed as displacing agents for nitrogen and oxygen.
A current of fluorine led through water displaces the oxygen, forming hydrogen fluoride; but the oxygen is in an allotropic state (see Part i.), called "ozone." Again, if a stream of chlorine is passed through, or if bromine-water be added to, a solution of ammonia, the hydrogen and chlorine combine, while the nitrogen is set free: 2NH3.Aq + 3C12 = 6HC1 + N2; but as ammonia combines with hydrogen chloride, the reaction 6NH3 + 6HC1 = 6NH4C1 occurs simultaneously; the complete equation is the sum of these two : 8NH3.Aq + 3C12 = 6NH4Cl.Aq + N2.
Chlorine, added to a solution of bromide or iodide of a metal, displaces the bromine or iodine; here the non-ionised chlorine becomes ionised at the expense of the electron of the ionised bromine or iodine, while the latter lose their electrons, thus: 2K.Aq + 2Br-.Aq + CI2.Aq = 2K.Aq + 2Cl-.Aq + Br2.Aq; the electron (-) leaves the bromine to attach itself to the chlorine. Similarly, bromine displaces iodine from a soluble iodide. But iodine displaces chlorine from the nearly insoluble silver chloride. Here, the iodide is still less soluble than the chloride; and as chloride dissolves, the less soluble and therefore non-ionised iodide is formed.
(g) Many metals are able to displace others. Thus, iron placed in a solution of a copper salt displaces the copper; copper displaces silver; silver, gold. In all these cases the action is doubtless an electrical one, and dependent on the replacement of a metal of lower by one of higher electric potential; that of higher potential becomes ionised, while that of lower assumes the metallic state, thus: Cu.Aq + 2Cl-.Aq + -Fe= = -Fe.Aq + 2CI-. Aq + Cu= ; 2Ag.Aq + 2NOg-Aq + Cu= = Cu„Aq + 2Cl-.Aq+2Ag-.
(Z>) There are some plans of obtaining elements which, though they can be referred to one or other of the three general methods exemplified already, are, on account of their complexity, better treated separately. Among these are the methods of separating hydrogen. The metals of the alkalies and alkaline earths attack water, forming hydroxides and liberating hydrogen : 2Na + 2H2O = 2NaOH + H2; Ca + 2H2O = Ca(OH)2 + H2. Magnesium powder, boiled with water, gives off hydrogen slowly; but zinc requires the presence of an acid, and must not be pure, i.e. there must be a foreign metal present to serve as the anode. The impurity usually present in commercial zinc is lead; the acid, for instance, sulphuric acid, is present in dilute solution as ions of HH and SO4= ; the surface layer of the zinc, Zn =, transfers its electrons to the lead, which is in metallic contact with the zinc. They in their turn pass on the electrons to the hydrogen which escapes in a non-ionised state as 2H -. It may then be collected over water, in which it is very sparingly soluble. Hydrogen, while it is on the point of receiving electrons and is still in the ionised state, may be used to liberate certain elements from their oxides or chlorides. Zinc and hydrochloric acid, for instance, in a solution of stannous chloride, SnCl2. Aq, causes a deposition of tin owing to the exchange of electrons ; instead of the hydrogen receiving electrons from the lead or other impurity in the zinc, the tin takes them up in its stead. If zinc and hydrochloric acid are placed in contact with silver chloride, AgCl, which is an insoluble compound, the zinc parts with electrons to the silver and itself enters into solution with electrons attached. Lastly, if generated in a solution of ferric chloride, Fe.Aq+3CI-Aq, the zinc goes into solution as before; and it transfers an electron to each ferric ion, changing it to the ferrous ions of ferrous chloride, -Fe.Aq + 2Cl-.Aq. The valency of the iron is lowered. Such processes are generally termed reduction ; the hydrogen doubtless plays the part of a carrier of ions ; it is said to be in the " nascent state," and is named the " reducing agent."