). About 60 percent of the sulfur dioxide produced throughout
the world comes from burning sulfur, and approximately 40 percent
is derived from roasting sulfide minerals. (Roasting is the
process by which ores are oxidized by heating in air.) Sulfur
dioxide is then oxidized to
sulfur trioxide, SO
. This oxidation reaction is exothermic (i.e., releases
energy in the form of heat) and reversible. Accordingly, a vanadium
oxide catalyst is used on an inert support to increase the rate
of the oxidation without decreasing the yield. Under optimum
conditions, the feed gas consists of equimolar quantities of
oxygen and sulfur dioxide (i.e., a 5:1 ratio of air to
sulfur dioxide) that passes through a four-stage catalytic converter
operating at various temperatures (see Figure
12). After the gas mixture has passed over three of the
catalyst beds and approximately 93 percent conversion to sulfur
trioxide has occurred, it is cooled and absorbed into sulfuric
acid in ceramic-packed towers. A final conversion of greater
than 99 percent is achieved after passage through the final
reaction bed (see Figure 12).
All three reactions used to produce sulfuric acid, as shown
below, are exothermic. Efficient utilization of this energy
to generate electricity, for example, is a key component
in maintaining the inexpensive price of this heavily used acid.
S + O 
SO
2SO + O
2SO
SO + H
O (in 98% H
SO
)
HSO
Pure sulfuric acid is a colourless, oily, dense (1.83 grams
per cubic centimetre) liquid that freezes at 10.5
C. It fumes when heated because of its decomposition to water
and sulfur trioxide. Because SO
has a lower boiling point than water, more SO
is lost during heating. When a concentration of 98.33 percent
acid is reached, the solution boils at 338
C without any further change in concentration. This is called
a constant boiling solution, and it is this concentration
that is sold as concentrated sulfuric acid. Anhydrous sulfuric
acid mixes with water in all proportions in a very exothermic
reaction. Adding water to concentrated acid can cause explosive
spattering. Because it reacts with organic compounds in the
skin, concentrated sulfuric acid can cause severe
burns. Thus, to decrease
the risk of injury in the laboratory, sulfuric acid should always
be added to water, slowly and with stirring to distribute the
heat.
{sup -} is nonetheless considered a moderately strong acid.
Because it is a diprotic acid, H
SO forms
two series of salts: hydrogen sulfates, HSO
{sup -}, and sulfates, SO
{sup 2-}. The sulfates of the alkaline earth metals--calcium
(Ca), strontium (Sr), and barium (Ba)--as well as that of lead
(Pb) are virtually insoluble, and these salts are found as naturally
occurring minerals. These important minerals include
gypsum (CaSO
2H
O), celestite (SrSO
), barite (BaSO
), and anglesite (PbSO
). These insoluble salts can be prepared in the laboratory by
metathesis reactions. A metathesis reaction
is one in which compounds exchange anion-cation partners. For
example, if a solution of barium nitrate, Ba(NO
), is added
to a solution of sodium sulfate, Na
SO, a precipitation
of barium sulfate, BaSO
, occurs. This is an important reaction because it can be used
as both a qualitative and quantitative test for the sulfate
ion and the barium ion. (Qualitative tests are used to determine
the presence or absence of a substance, while quantitative tests
are used to measure the amount of a constituent.) In addition
to metathesis reactions, sulfate salts can generally be prepared
by dissolution of metals in aqueous H
SO, neutralization
of aqueous H
SO with metal
oxides or hydroxides, oxidation of metal sulfides (a sulfide
contains S{sup 2-}) or sulfites (SO
{sup 2-}), or decomposition of salts of volatile acids, such
as carbonates, with aqueous H
SO. Some
important soluble sulfate salts are Glauber's salt, Na
SO
10HO; Epsom
salt, MgSO
7H
O; blue vitriol, CuSO
5H
O; green vitriol, FeSO
7H
O; and white vitriol, ZnSO
7H
O.SO
undergoes extensive self-ionization (sometimes called
autoprotolysis).
2HSO
H
SO{sup +}
+ HSO{sup
-}
This autoprotolysis reaction is, however, only one of the equilibrium
reactions that occur in pure H
SO to give
it an extremely high electrical conductivity. There are three
additional equilibrium reactions that take place because of
the ionic self-dehydration of sulfuric acid.
2HSO 
HO{sup +}
+ HSO
{sup -}
HO + H
SO
HO{sup +}
+ HSO{sup
-}
HS
O + H
SO
HSO
{sup +} + HS
O{sup -}
Thus, there are at least seven well-defined species that exist
in "pure" H
SO. The value
of the dielectric constant of the acid is also quite high (
= 100).
Concentrated sulfuric acid is not a very strong oxidizing agent
unless it is hot. When it acts as an oxidizing agent, however,
it can be reduced to several different sulfur species, including
SO, HSO
{sup -}, SO{sup
2-}, elemental sulfur (S
), hydrogen sulfide (H
S), and the sulfide anion, (S{sup 2-}). Concentrated sulfuric
acid is a good dehydrating agent, as it reacts with many organic
materials to remove the elements of water.
The amount of sulfuric acid used in industry exceeds that of any other manufactured compound. In the United States approximately 67 percent of the acid is utilized to convert phosphate rock to phosphoric acid. The phosphoric acid is then converted to phosphate fertilizers. Other major uses include the refining of petroleum, the removal of impurities from gasoline and kerosene, the pickling of steel (the cleaning of its surface), and the manufacture of other chemicals, such as nitric and hydrochloric acids. It also is utilized in lead storage batteries and in the production of paints, plastics, explosives, and textiles.
