Sulphuric acid




Sulphuric acid

Sulphuric acid is probably the most important of all chemicals,
because of its extensive use in a very large number of manufacturing operations. Of the immense quantities made yearly, the greater part does not come upon the market; for, being expensive and difficult to ship, consumers of large amounts generally make their own acid.
The commercial grades of acid have special names. A moderately strong acid (50°_55° He.), such as condenses in the lead chambers, is known as "chamber acid." It contains from 62 to 70 per cent of H2S04, and is strong enough for use in the manufacture of fertilizer, and for other purposes requiring a dilute acid. By concentrating this chamber acid, an acid of 60° Be. is obtained, containing about 78 per cent of H2S04> which is sufficiently strong for most technical uses. Further evaporation in platinum or iron pans yields an acid of 66° Be., containing 93.5 per cent of H2S04, and known as oil of vitriol. Faming or Nordhausen acid, which is still more concentrated, is prepared by special means. It is essentially a solution of sulphuric anhydride (S03) in sulphuric acid. This is the acid which was prepared by the alchemists in the Middle Ages.
In about the year 1740, Ward, an Englishman, began to make sulphuric acid on a moderately large scale. He burned sulphur and nitre (KNO3) together, and condensed the vapors in glass vessels containing a little water. The dilute acid so formed was then concentrated in glass alembics or retorts. In this way all acid was produced at a lower price than the fuming acid could be made. and the industry was soon established on [t commercial scale. The reactions involved ill Ward's process are the bases of the method now in use; this consists in bringing together, under suitable conditions, sulphur dioxide, oxygen, and water as steam, in the presence of certain oxides of nitrogen. The latter probably act as carriers of the oxygen, causing it to unite with the sulphur dioxide and water to form the acid. The apparent reaction is;-

S02 + H20 + 0 = H2S04-

Hut this does not represent the actual process, which is more
complicated than it at first appears. Several theories have been advanced to explain the reactions occurring in the lead chambers, and the part taken by the nitrogen oxides, but the most generally accepted one, that of Lunge, regards nitrous anhydride (N2O3) * as the essential factor. t According to this view, the principal reactions involved are as follows:

1) 2S02 + N2O3 + O2 + H20 = 2S02.(OH).(ONO) (Nitrosylsulphuric acid);
2) 2 S02.(OH).(ONO) + H20 = 2 S02(OH)2 + N2O3; or,
3) 2 SO2(OH)(ONO) + S02 +0 +H20 = 3 S02(OH)2 +N203.

First there is a union of sulphur dioxide, nitrous anhydride,
oxygen, and water, to form nitrosylsulphuric acid, which probably separates as part of the mist or fog seen in the lead chambers. But in the presence of water vapor or of dilute sulphuric acid, this nitrosylsulphuric acid is at once decomposed, according to reaction (2), sulphuric acid being formed, and nitrous anhydride regenerated; or if sulphur dioxide and oxygen are concerned in the process, then reaction (3) occurs. This cycle of reactions repeats an indefinite number of times. But in the first lead chamber, where the temperature is rather high and an excess of water vapor is usually present, the following secondary reactions probably occur
to a greater or less extent:-

4) 2 S02.(OH).(ONO) + S02 + 2 H20 = 3 H2S04 + 2 NO,

this reaction being only momentary.
Since there is usually an excess of oxygen present, however, the nitric oxide here formed is at once brought into action again, thus:-

5) 2 S02 + 2 NO + 30 + H20 = 2 S02.(OH).(0N0).

If there is a deficiency of oxygen, the nitric oxide is not returned to the process, but passes through the several chambers and, since it is not absorbed by the concentrated acid in the Gay-Lussac tower, it escapes into the atmosphere and is lost.
The nitrogen oxides are derived from nitric acid, or by the action of sulphuric acid on sodium nitrate in the nitre pots. When nitric acid is used, it must be introduced in the form of vapor, or at least as a very fine spray, whereupon it reacts as follows:-

6) 2 S02 + 2 HNO2 + H20 = 2 H2S04 + N203,

Perhaps this reaction really occurs in two stages, thus:-

(a) S02 + HNO3 = S02.(0H). (ONO);
(b) 2 S02.(OH).(ONO) + H20 =2 H2S04 + N203.

The formula assigned to the nitrosylsulphuric acid may perhaps be written:

and the compound would then be called nitrosulphonic acid. But in either ease the existence of the substance is only transitory, it being broken up at once by the steam and sulphur dioxide present when the process is working properly. In case there is a deficiency of water vapor in the chambers, and especially if the temperature falls too low, the nitrosylsulphuric acid may separate as crystals, which deposit at various points on the walls, forming "chamber crystals" This is an undesirable accident, for when steam or water come in contact with them they decompose into sulphuric acid, nitric acid, and nitrogen peroxide (N204) :-

4 S02 . (OH) . (ONO) + 2 H20 = 4 H2S04 + N204 + 2 NO.

Then the nitrogen peroxide unites with some of the water,

N204 + H20 = HN02 + HNO3,

forming nitrous and nitric acids directly on the walls, corroding the lead at the point where the cluster of crystals was attached. To prevent this separation of "chamber crystals" and the retention of nitrogen oxides in the sulphuric acid. an excess of steam in the lead chambers is preferred, although it dilutes the acid somewhat.
The raw materials employed in sulphuric acid making are:-
1. Sulphur.
(a) Crude brimstone.
(b) Metallic sulphides, such as iron pyrites, chalcopyrite
(copper pyrites), sphalerite (zinc blende), etc.
(c) Hydrogen sulphide (seldom used).
2. Sodium nitrate or nitric acid.
3. Water as steam.
4. Oxygen as air.

The acid may be made from brimstone, pyrites, blende, or hydrogen sulphide, but they are not used together.
Since the burners and chambers employed for one source of sulphur cannot be adapted to that of another without extensive alterations, the manufacturer must decide what material he will use, and erect his plant accordingly. Crude sulphur from the calceroni gives a very pure acid free from arsenic, iron, copper, or zinc, and much smaller condensing chambers may be used for a given yield than when pyrites or blende is employed. The sulphur is placed in iron retorts or on trays in brick ovens, and ignited. The combustion is easily controlled by regulating the amount of air admitted to the retort. Usually the hot vapors from the brimstone burners are
passed through a narrow flue or passage, into which a regulated supply of ail' is admitted. This insures complete combustion of any sulphur that may distil owing to too great heat in the retort. To prevent clogging and loss, as much of the sulphur as possible is converted to dioxide.
Pyrites, or natural disulphide of iron (FeS2), is a dense, hard
mineral of crystalline structure and pale yellow color. The largest deposits in the United States are in Virginia, at Mineral City, and at Charlemont in Massachusetts. Of the foreign deposits, those in Spain are the most important. A pure pyrites contains 53.3 per cent of sulphur, but that commonly used for acid making carries from 43 to 48 per cent. It seldom pays to use an ore with less than 35 per cent of sulphur, for it will not support its own combustion.
The first proposal to use pyrites originated with an Englishman named Hill, who took out a patent for the process in 1818. But it was not until 1838, when the Sicilian government sold the monopoly of the sulphur export to a French firm which nearly trebled the price of crude brimstone, that pyrites began to find favor with acid makers, At the present time, because it is cheap and easily obtained, pyrites has almost completely replaced sulphur for acid making. The product from pyrites is usually contaminated with arsenic, and often with zinc, copper, and selenium.
By the oxidation of pyrites in a suitable furnace, the sulphur is converted to dioxide, and iron oxide remains. The reaction may
be written as follows:-

2 FeS2 + 11 0 =4 S02+Fe203

This is not exact, however, as some sulphur remains in the ore
and some sulphur trioxide is formed. The proper regulation of the pyrites burners is one of the problems of the manufacturer. If the ore contains over 35 per cent of sulphur, the burning, once started, generates sufficient heat to maintain the combustion, and no fuel is necessary. But zinc sulphide and the" mattes" from metallurgical processes must be heated by fuel.
The complete burning of pyrites is difficult. With lump ore there is apt to be a kernel in the centre of the lump, from which the sulphur is not burned out. If the temperature rises too high, the charge fuses together, forming clinkers or "scar," and choking the furnace. If too much air is admitted, the furnace cools below the temperature at which fresh pyrites will ignite, and the gases leaving the burner are so diluted that the desired reactions do not take place in the lead chambers. With "smalls" the tendency to fuse is more marked than with lump ore, and the fine ore packs together so densely that the air will not penetrate it, and unless it is constantly stirred only the surface is burned. (The lump ore is that which has been broken to about the size of a goose egg, the "smalls" constituting what will pass through a half-inch screen.)
Pyrites burners are usually built in benches containing from three to thirty furnaces, in order that the supply of gas may not be broken, while charging or cleaning one furnace.
A burner for lump ore (Fig. 21) * consists of a brick furnace, containing a grate formed of single loose iron bars (B, B) having a square section, and resting in grooves at each end. These bars may be turned parallel with their longitudinal axes, but have no lateral motion. They are so adjusted that their sides are at an angle of 45° to the vertical. After a charge is burned, the bars are given several quarter turns by means of a key, to allow the cinders on them to drop through into the ash pit. Ail' is admitted by dampers beneath the grate. When properly working, the cinders resting on the bars are nearly cold, the hottest part of the fire being eighteen inches above the grate. The furnaces are lined with fire-brick, and to prevent any access of air except through the dampers, the doors (0, D) for charging, cleaning, raking, etc., are make to fit closely, and are generally luted with clay.
All the burners in one bench deliver their sulphur dioxide gas
into a common, wide flue, or "dust box" (F), where any fine dust carried along by the gases may settle before they enter the Glover towel'. This dust consists of unburned pyrites, arsenic, antimony or zinc oxides, iron oxide, etc.
In one or more of the burners a cast-iron "nitre pot" (P) may be set, in which nitrous gases are generated by the action of sulphuric

chamber built into the flue (F), and heated by the waste heat of
the burners. Sometimes, however, the pots are placed in separate
furnaces.
A lump burnersof average size has a grate area of 15 to 25 square
feet. The furnace is sometimes made slightly hopper-shapedinside,
so that it is larger at the level of the charging doors than at the grate bars.' About 40 pounds of pyrites, containing 48 per cent of sulphur, are burned per square foot of grate area in 24 hours, a
larger quantity of such high-grade ore being liable to cause fusion,
unless great care is exercised. A larger quantity of poorer ore may
be burned daily, without danger of fusion.


Organic Chemistry for the industry

Inorganic Chemistry for the industry

  • Lixiviation
  • Levigation
  • Evaporation
  • Distillation
  • Sublimation
  • Filtration
  • Crystallization
  • Calcination
  • Refrigeration
  • Density
  • Fuels
  • Liquid fuels
  • Gaseous fuels
  • Water
  • Sulphur
  • Sulphur Derivatives
  • Sulphuric Acid
  • Sulphuric acid burners
  • Fuming Sulphuric acid
  • Salt
  • Hydrochloric Acid
  • Soda Industry
  • Caustic Soda
  • Treatment of tank
  • Ammonia Soda
  • Cryolite Soda process
  • Chlorine Industry
  • Electrolytic Chlorine
  • Hypochlorites
  • Chlorates
  • Nitric Acid
  • Nitrates
  • Ammonia
  • Potash Industry
  • Fertilizers
  • Lime, Cement
  • Cement
  • Glass
  • Ceramic Industries
  • Pigments
  • Bromine
  • Iodine
  • Phosphorus
  • Boric Acid
  • Arsenic Compounds
  • Peroxides
  • Oxygen
  • Sulphates
  • Alum








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