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LECTURE 9 SULFUR AND SULFURIC ACID

Ch 61 Industrial Chem

SULFUR - USES ¨  a component of black gunpowder, ¨  vulcanization of natural rubber, ¨  used as a fungicide and ¨  insecticide, and ¨  used to make fertilizers. ¨  used to produce sulfuric acid for batteries ¨  fireworks, paper, and some medicines. ¨  Sulfur dioxide is a cheap and excellent refrigerant, and was once widely

used in home refrigerators before it was replaced by the freons.

NATURAL SOURCES ¨  Deposits ¨  Sour Crude Oil ¨  Natural Gas (sour)

RECOVERY Recovery of sulphur is done in three basic ways: ¨  Mined through the use of wells drilled to sulphur

deposits and worked with the “Frasch” method; (See Diagram)

¨  Extracted from the oil or gas stream at a processing plant;

¨  Scraped from the surface of the earth or dug out of open pits.

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About 11 tons of sulfur are extracted daily from volcanic craters and are then used in sugar refineries, cosmetics, medicines and explosives.

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MINING – THE FRASCH METHOD The Frasch method, developed by Herman Frasch of Canada, must be used to recover these deposits. This process consists of circulating superheated hot water and steam through the well melting the sulfur. (MP = 115oC) Then, compressed air is forced into the well, pushing the liquid sulfur up to the surface. This allows the recovery of the sulfur at the surface, where it is metered, cooled, and solidified.

EXTRACTION FROM SOUR CRUDE OIL ¨  Sour crude oil is defined as crude oil containing sulfur. During

the early 1900s, sour crude had been considered almost unusable because of its odor and poor burning characteristics.

¨  Herman Frasch developed a method for using sour crude to make kerosene. Today, since unleaded gasoline for automobiles is mandated for air pollution reasons, oil refiners have been forced to install sulfur extraction units in order to remove the sulfur from the crude oil in order to produce unleaded gasoline.

EXTRACTION FROM SOUR NATURAL GAS Extraction of sulfur from natural gas is another source for sulfur. The natural gas is sent to a processing plant where the hydrogen sulfide (sulfur) is extracted and where sweet (marketable) gas remains. A method called the “Claus Process” is used to convert the sulphur compounds in sour gas to elemental sulphur. This process, invented by English scientist C. F. Claus, is now the most widely used method for extraction of sulfur from sour gas.

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CLAUS PROCESS H2S(g) + 3/2O2(g) D SO2(g) + H2O(g) ∆H = -518.8 kJ

SO2(g) + 2H2S(g) D 3S(l) + H2O(g) ∆H = -142.8 kJ

Catalyst: Fe2O3

SULFURIC ACID - USES

plays some part in the production of nearly all manufactured goods.

production of fertilizers, e.g., superphosphate of lime and ammonium sulfate.

manufacture of chemicals, e.g., in making hydrochloric acid, nitric acid, sulfate salts, synthetic detergents, dyes and pigments, explosives, and drugs.

Production of Sulfuric Acid reached 98.5 million tons in 2011 and is expected to reach 120 million tons by 2017. SULFURIC ACID - USES

It is used in petroleum refining to wash impurities out of gasoline and other refinery products. Sulfuric acid is used in processing metals, e.g., in pickling (cleaning) iron and steel before plating them with tin or zinc. Rayon is made with sulfuric acid. It serves as the electrolyte in the lead-acid storage battery commonly used in motor vehicles (acid for this use, containing about 33% H2SO4

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¨  The contact process is the current method of producing sulfuric acid in the high concentrations needed for industrial processes. Vanadium(V) oxide is the catalyst employed.

¨  This process was patented in 1831 by the British vinegar merchant Peregrine Phillips, as a far more economical process for producing sulphur trioxide and concentrated sulfuric acid than the previous lead chamber process method used for producing sulfuric acid.

THE CONTACT PROCESS

The process can be divided into three stages: ¨  Preparation and purification of sulfur dioxide ¨  Catalytic oxidation (using vanadium oxide catalyst) of sulfur

dioxide to sulfur trioxide ¨  Conversion of sulfur trioxide to sulfuric acid

¨  Purification of air and SO2 is necessary to avoid catalyst poisoning (i.e. removing catalytic activities). The gas is then washed with water and dried by H2SO4.

¨  To conserve energy, the mixture is heated by exhaust gases from the catalytic converter by heat exchangers.

¨  Sulfur dioxide and oxygen then react in the manner as follows:

2SO2(g) + O2(g) ↔ 2 SO3(g) : ΔH = −197 kJ mol−1

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To increase the reaction rate, high temperatures (450 °C), high pressures (200 kPa or 2 atm), and vanadium(V) oxide (V2O5) are used to ensure a 99.5% conversion. Platinum would be a more suitable catalyst, but it is very costly and easily poisoned.

MAKING SULFUR DIOXIDE ¨  This can either be made by burning sulfur in an excess of

air: . . . or by heating sulfide ores like pyrite in an excess of air:

¨  In either case, an excess of air is used so that the sulfur dioxide produced is already mixed with oxygen for the next stage.

Converting the sulfur dioxide into sulfur trioxide ¨  This is a reversible reaction, and the formation of

the sulphur trioxide is exothermic. A flow scheme for this part of the process looks like this:

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¨  Hot sulfur trioxide passes through the heat exchanger and is dissolved in concentrated H2SO4 in the absorption tower to form oleum: H2SO4(l) + SO3 (g) → H2S2O7(l)

¨  Note that directly dissolving SO3 in water is impractical due to the highly exothermic nature of the reaction. Acidic vapor or mists are formed instead of a liquid.

¨  Oleum is reacted with water to form concentrated H2SO4. H2S2O7(l) + H2O(l) → 2 H2SO4(l)

Converting the sulphur trioxide into sulphuric acid ¨  This can't be done by simply adding water to the

sulfur trioxide - the reaction is so uncontrollable that it creates a fog of sulfuric acid.

¨  Instead, the sulfur trioxide is first dissolved in concentrated sulfuric acid:

¨  The product is known as fuming sulfuric acid or oleum.

This can then be reacted safely with water to produce concentrated sulphuric acid - twice as much as you originally used to make the fuming sulphuric acid.

¨  The mixture of sulphur dioxide and oxygen going into the reactor is in equal proportions by volume.

¨  Avogadro's Law says that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. That means that the gases are going into the reactor in the ratio of 1 molecule of sulphur dioxide to 1 of oxygen.

¨  That is an excess of oxygen relative to the proportions demanded by the equation.

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¨  According to Le Chatelier's Principle, Increasing the concentration of oxygen in the mixture causes the position of equilibrium to shift towards the right. Since the oxygen comes from the air, this is a very cheap way of increasing the conversion of sulphur dioxide into sulphur trioxide.

¨  Why not use an even higher proportion of oxygen? This is easy to see if you take an extreme case. Suppose you have a million molecules of oxygen to every molecule of sulphur dioxide.

¨  The equilibrium is going to be tipped very strongly towards sulphur trioxide - virtually every molecule of sulphur dioxide will be converted into sulphur trioxide. Great! But you aren't going to produce much sulphur trioxide every day. The vast majority of what you are passing over the catalyst is oxygen which has nothing to react with.

¨  By increasing the proportion of oxygen you can increase the percentage of the sulfur dioxide converted, but at the same time decrease the total amount of sulfur trioxide made each day. The 1 : 1 mixture turns out to give you the best possible overall yield of sulfur trioxide.

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Equilibrium considerations - Temperature

¨  You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of sulfur trioxide in the equilibrium mixture.

¨  The forward reaction (the production of sulfur trioxide) is exothermic.

¨  According to Le Chatelier's Principle, this will be favoured if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat.

¨  In order to get as much sulfur trioxide as possible in the equilibrium mixture, you need as low a temperature as possible. However, 400 - 450°C isn't a low temperature!

Rate considerations ¨  The lower the temperature you use, the slower the reaction

becomes. A manufacturer is trying to produce as much sulfur trioxide as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of sulfur trioxide if it takes several years for the reaction to reach that equilibrium.

¨  You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor.

The compromise ¨  400 - 450°C is a compromise temperature

producing a fairly high proportion of sulfur trioxide in the equilibrium mixture, but in a very short time.

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Equilibrium considerations - Pressure ¨  Notice that there are 3 molecules on the left-hand side of the equation,

but only 2 on the right. ¨  According to Le Chatelier's Principle, if you increase the pressure the

system will respond by favoring the reaction which produces fewer molecules. That will cause the pressure to fall again.

¨  In order to get as much sulfur trioxide as possible in the equilibrium mixture, you need as high a pressure as possible. High pressures also increase the rate of the reaction. However, the reaction is done at pressures close to atmospheric pressure!

Economic considerations ¨  Even at these relatively low pressures, there is

a 99.5% conversion of sulphur dioxide into sulphur trioxide. The very small improvement that you could achieve by increasing the pressure is not worth the expense of producing those high pressures.

Rate considerations – The Catalyst

¨  The catalyst has no effect whatsoever on the position

of the equilibrium. Adding a catalyst doesn't produce any greater percentage of sulphur trioxide in the equilibrium mixture. Its only function is to speed up the reaction.

Rate considerations ¨  In the absence of a catalyst the reaction is so

slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor.

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Ch 16 / 17 P and K industries Ch 9 Cement / Ceramic Industries Ch 22 Explosives, Propellants Ch 23 Food and Food By-Products Ch 37/38 – Petroleum Processing and Petrochemicals Ch 29 Soap and Detergents Ch 34 Plastic Industries Ch 40 Pharmaceutical Environmental Protection, R&D, Safety and Regulatory Intellectual Property