Aluminum and its reaction with water

Alu­minum was first ob­tained in the ear­ly 19th cen­tu­ry, by the physi­cist Hans Chris­tian Ørst­ed. He con­duct­ed his ex­per­i­ment with an potas­si­um amal­gam, alu­minum chlo­ride, and with the sub­se­quent dis­til­la­tion of mer­cury.

In­ci­den­tal­ly, the name of this sil­very ma­te­ri­al came from the Latin word for alum (alu­men), be­cause this el­e­ment is pro­duced from them.

Alums are nat­u­ral min­er­als on the ba­sis of met­als which com­bine a hy­drat­ed dou­ble sul­fate salt in their com­po­si­tion.

Pre­vi­ous­ly, alu­minum was con­sid­ered to be a pre­cious met­al and cost con­sid­er­ably more than gold. This was be­cause the met­al was rather dif­fi­cult to sep­a­rate from im­pu­ri­ties. So only rich and in­flu­en­tial peo­ple could af­ford jew­el­ry made of alu­minum.

But soon, in 1886, a method was de­vel­oped by Charles Mar­tin Hall and Paul Héroult for pro­duc­ing alu­minum in in­dus­tri­al scale, which dras­ti­cal­ly re­duced the cost of this met­al, and al­lowed it to be used in the met­al­lur­gi­cal in­dus­try. The in­dus­tri­al method in­volved the elec­trol­y­sis of an al­loy of cry­o­lite, in which alu­minum ox­ide was dis­solved. It is dif­fi­cult to ap­pre­ci­ate the val­ue of this met­al in its present form, as many items made of alu­minum are used by peo­ple in ev­ery­day life.

Ap­pli­ca­tion of alu­minum

Thanks to its mal­leabil­i­ty and light­ness, and also its cor­ro­sion re­sis­tance, alu­minum is a valu­able met­al of great im­por­tance in mod­ern in­dus­try, from which such house­hold items are made as kitchen­ware, and in in­dus­try: for ex­am­ple, it is wide­ly used in air­craft and au­to­mo­bile con­struc­tion.

Alu­minum is also one of the most in­ex­pen­sive and eco­nom­i­cal ma­te­ri­als, as it can be used end­less­ly, by melt­ing down used alu­minum cans and oth­er ob­jects.

Metal­lic alu­minum is safe, but its com­pounds can have a tox­ic ef­fect on hu­mans and an­i­mals, par­tic­u­lar­ly alu­minum chlo­ride, ac­etate, and sul­fate.

The phys­i­cal prop­er­ties of alu­minum

Alu­minum is a sil­very-white light met­al, which can form al­loys with the ma­jor­i­ty of met­als, es­pe­cial­ly cop­per and mag­ne­sium, and also sil­i­con. It is also very duc­tile, and can eas­i­ly be turned into a thin sheet or foil. The melt­ing point of alu­minum is 660 °C, and its boil­ing tem­per­a­ture is 2470 °C.

The chem­i­cal prop­er­ties of alu­minum

If alu­minum is left at room tem­per­a­ture, a durable pro­tec­tive lay­er of alu­minum ox­ide Al₂O₃ forms on the met­al, which pro­tects it from cor­ro­sion.

Alu­minum prac­ti­cal­ly does not re­act with ox­i­diz­ers be­cause of the ox­ide film that pro­tects it. But this film can be eas­i­ly de­stroyed, so that the met­al dis­plays ac­tive re­duc­tive prop­er­ties. The ox­ide film of alu­minum can be de­stroyed by a so­lu­tion or al­loy of al­ka­lis, or with acids, and also with mer­cury chlo­ride or ox­ide. Thanks to its re­duc­ing prop­er­ties, alu­minum has found ap­pli­ca­tion in in­dus­try for ob­tain­ing oth­er met­als or non-met­als, and this process is called alu­minothermy. This fea­ture of alu­minum in­volves the in­ter­ac­tion with ox­ides of oth­er met­als.

For ex­am­ple, let us ex­am­ine the re­ac­tion with chromi­um ox­ide:

2Al + Cr₂O₃ = Al₂O₃ + 2Cr

Alu­minum eas­i­ly re­acts with sim­ple sub­stances. For ex­am­ple, with halo­gens, flu­o­rine is the ex­cep­tion, alu­minum can form alu­minum io­dide, chlo­ride or bro­mide:

2Al+3Cl₂ = 2Al­Cl₃

With oth­er non-met­als such as flu­o­rine, sul­fur, ni­tro­gen, car­bon etc. alu­minum can only re­act when heat­ed.

The sil­very met­al also re­acts with com­plex chem­i­cal sub­stances. For ex­am­ple, with al­ka­lis it forms alu­mi­nates. i.e. com­plex com­pounds which are ac­tive­ly used in the pa­per and tex­tile in­dus­try. It en­ters into the re­ac­tion as alu­minum hy­drox­ide

Al(ОН)₃ + NaOH = Na[Al(OH)₄],

and as metal­lic alu­minum or alu­minum ox­ide:

2Al + 2NaOH + 6Н₂О = 2Na[Al(OH)₄] + ЗН₂↑

Al₂O₃ + 2NaOH + 3H₂O = 2Na[Al(OH)₄]

With strong acids, for ex­am­ple, sul­fu­ric and hy­drochlo­ric, alu­minum re­acts quite calm­ly, with­out ig­ni­tion.

If a piece of the met­al is put in hy­drochlo­ric acid, a slow re­ac­tion takes place, as ini­tial­ly the ox­ide film will be dis­solved from its sur­face, but then the re­ac­tion speeds up. Alu­minum re­acts with hy­drochlo­ric acid with the re­lease of hy­dro­gen, and the re­sult of the re­ac­tion is alu­minum chlo­ride:

Al₂O₃ + 6HCl = 2Al­Cl₃ + 3H₂O

2Al+6HCl→2Al­Cl₃ +3H₂↑

The re­ac­tion of alu­minum with wa­ter

If you take an alu­minum fil­ing and put it in wa­ter, noth­ing will hap­pen, be­cause alu­minum is pro­tect­ed by an ox­ide film, which does not al­low the met­al to join into a re­ac­tion.

Only af­ter the pro­tec­tive film is re­moved with mer­cury chlo­ride can there be a re­sult. The met­al must be soaked for two min­utes in a so­lu­tion of mer­cury chlo­ride, and then rinsed well. As a re­sult, an amal­gam forms, an al­loy of mer­cury and alu­minum:

3Hg­CI₂ + 2Al = 2Al­CI₃ + 3Hg

The amal­gam does not stay on the sur­face of the met­al. Now, af­ter putting the cleaned met­al in wa­ter, you can ob­serve a slow re­ac­tion, which is ac­com­pa­nied by the re­lease of hy­dro­gen and the for­ma­tion of alu­minum hy­drox­ide:

2Al + 6H₂O = 2Al(OH)₃ + 3H₂↑