$44 PRINCIPLES OF CHEMISTRY 



'Iron forms one other oxide besides the ferric and ferrous oxides ; 

 this contains twice as much oxygen as the former, but is so very 

 unstable that it can neither be obtained in the free state nor as a 

 hydrate. Whenever sjjch conditions of double decomposition occur as 

 should allow of its serration in the free state, it decomposes into 

 Oxygen and ferric oxide. It is known in the state of salts, and is only 

 stable in the presence of alkalis, and forms salts with them which have 

 a decidedly alkaline reaction ; it is therefore a feebly acid oxide. Thus 

 when small pieces of iron are heated with nitre or potassium chlorate 

 a potassium salt of the composition K 2 Fe0 4 is formed, and therefore 

 the hydrate corresponding with this salt should have the composition 

 H 2 FeO 4 . It is called ferric acid. Its anhydride ought to contain 

 Fe0 3 or Fe 2 O 6 twice as much oxygen as ferric oxide. If a solution 

 of potassium ferrate be mixed with acid, the free hydrate ought 

 to be formed, but it immediately decomposes (2K 2 FeO 4 -f SH.^SC^ 

 = 2K 2 S0 4 + Fe 2 (SO 4 ) 3 + 5H 2 O + O 3 ), oxygen being evolved. If a 

 small quantity of acid be taken, or il; a solution of potassium ferrate 

 be heated with solutions of other metallic salts, ferric oxide is sepa- 

 rated for instance : 



2CuS0 4 + 2K 2 Fe0 4 = 2K 2 S0 4 + 3 -f Fe 2 3 + 2CuO. 



Both these oxides are of course deposited in the form of hydrates. 

 This shows that not only the hydrate H 2 FeO 4 , but also the salts of the 

 heavy metals corresponding with this higher oxide of iron, are not 

 formed by reactions of double decomposition. The solution of potas- 

 sium ferrate naturally acts as a powerful oxidising agent ; for instance, 

 it transforms manganous oxide into the dioxide, sulphurous into 

 sulphuric acid, oxalic acid into carbonic anhydride and water, &c. 26 

 Iron thus combines with oxygen in three proportions : RO, R 2 3 , 



gross of the reaction was expressed by the amount of liberated iodine in percentages 



of the theoretical amount. For instance, the following amount of iodide of potassium 

 was decomposed when Fe 2 (S0 4 ) 5 +2nKI was taken: 



> 1 28 6 10 20 



After 15' 11-4 26'3 40'6 78'5 01'8 96'0 



p 80* 14-0 85-8 47-8 78'5 94'3 97'4 



Ihour 19-0 42'7 66'0 84'0 95'7 97'6 



10 ' 82-6 56-0 75-7 93'2 96'5 97'6 



48. 89'4 C7'7, 82'6 93'4 96'6 97*6 



Similar results were obtained for FeCl 3 , but then the amount of iodine liberated was 

 eomewhat greater. Similar results .were also obtained by increasing the mass of FeX$ 

 per KI, and by replacing it by HI (see Chapter XXI., Note 26). 



|,a 26 If chlorine be passed through a strong solutidn of potassium hydroxide in which 

 hydrated ferric oxide is suspended, the turbid liquid acquires a dark pomegranaie-red 

 colour and contains potassium ferrate : 10KHO + Fe 3 3 + 8C1 3 -' 2KjFeO 4 + 6KC1 + SlLjO. 

 The chlorine must not be in excess, otherwise the salt is again decomposed, although the 

 mode of decomposition is unknown ; however, ferric chloride and potassium chlora ,0 

 are probably formed. Another way in which the above-described salt is formed is also 



