Can plants directly use sulfuric acid?

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In the explanation of the sulfur cycle, it is often said that sulfur moves from the atmosphere to the ground by acid rain in the form of sulfuric acid. Can plants directly use sulfuric acid to assimilate sulfur?

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Plants have sulfate transporters which they use to assimilate sulfur. Since sulfate is the conjugate base of sulfuric acid, this could be construed as a "yes" for your question. Sulfuric acid would be present in its conjugate base form at physiological pH values.

However, it would be unwise to water your plants with sulfuric acid, as many plants have an optimal pH range which will quickly cause death when sulfuric acid lowers the pH of the soil below that range. As a result, areas afflicted by acid rain may require treatment with alkaline limestone to resolve this issue.

Sulfuric Acid

Sulfuric acid (H₂S0₄) is a corrosive substance, destructive to the skin, eyes, teeth, and lungs. Severe exposure can result in death. Workers may be harmed from exposure to sulfuric acid. The level of exposure depends on dose, duration, and type of work being done.

Sulfuric acid is used in many industries. It&rsquos used to produce other chemicals, explosives and glue to refine petroleum to cure metal and in lead-based car batteries. Some examples of workers at risk of being exposed to sulfuric acid include the following:

• Outdoor workers who work in areas where coal, oil, or gas are burned
• Mechanics who handle dirty batteries
• Plumbers and contractors who come in contact with toilet bowl cleaners mixed with water
• Workers in publishing, printing or photography shops
• Fire fighters and steelworkers who are exposed to acid mists

NIOSH recommends that employers use the Hierarchy of Controls to prevent or reduce worker exposures. If you work in an industry that uses sulfuric acid, read chemical labels and the accompanying Safety Data Sheet for hazard information. Visit NIOSH&rsquos page on Managing Chemical Safety in the Workplace to learn more about controlling chemical workplace exposures.

The following resources provide information about occupational exposure to sulfuric acid. Useful search terms for sulfuric acid include &ldquobattery acid,&rdquo &ldquohydrogen sulfate,&rdquo &ldquooil of vitriol,&rdquo and &ldquosulfuric acid (aqueous).&rdquo

NIOSH Chemical Resources

The NIOSH Pocket Guide to Chemical Hazards (NPG) helps workers, employers, and occupational health professionals recognize and control workplace chemical hazards.

The NIOSH Manual of Analytical Methods (NMAM) is a collection of methods for sampling and analysis of contaminants in workplace air, and in the blood and urine of workers who are occupationally exposed.

The Health Hazard Evaluation Program (HHE) conducts onsite investigations of possible worker exposure to chemicals. Search the HHE database for more information on chemical topics.

This is a straightforward question with a not very straightforward answer. Acid rain and plant damage go hand in hand in areas prone to this type of pollution, but the changes to a plant’s environment and tissues are gradual. Eventually, a plant exposed to acid rain will die, but unless your plants are incredibly sensitive, the acid rain unusually potent and frequent or you’re a very bad gardener, the damage is not fatal.

The way that acid rain damages plants is very subtle. Over time, the acidic water alters the pH of the soil where your plants are growing, binding and dissolving vital minerals and carrying them away. As the soil pH falls, your plants will suffer increasingly obvious symptoms, including yellowing between the veins on their leaves.

Rain that falls on leaves can eat away the outer waxy layer of tissue that protects the plant from drying out, leading to the destruction of the chloroplasts that drive photosynthesis. When a lot of leaves are damaged at once, your plant may become very stressed and attract a host of pests and diseases organisms.

Effects of Different Treatments on Seed Germination Improvement of Calotropis persica

The purpose of this study was to investigate the effects of different treatments on seed germination in the desert plant species Calotropis persica (Gand.). This species is known to have long time for seed germination considering arid region condition and short time of access moist. An experiment was performed with 13 treatments and 4 replications in a completely randomized design. Treatments included KNO₃ with concentrations of 0.1, 0.2, and 0.3 percent, immersion in hot water for five min, acetylsalicylic acid 100, 200, and 300 mg L −1 , ethereal sulfuric acid (60%) for 5 and 10 min, thiourea with concentrations of 0.1% and 0.3%, and prechilling for 10 days. Tap water was used as the control. Our findings indicate that KNO₃ 0.1% and 100 mg L −1 acetylsalicylic acid were the most effective treatments for improvement of seed germination properties in this species. In a comparison of the two mentioned treatment, KNO₃ 0.1% treatments is the best.

1. Introduction

Germination is a critical stage in the life cycle of weeds and crop plants and often controls population dynamics, with major practical implications. Seed germination is the critical stage for species survival [1, 2]. In recent 20 years, desertification has been recognized as a major environmental problem and is a major focus of United Nations Environment Programme [3]. Vegetation is a protector of the soil against water and wind erosion as well as a casualty of soil erosion [4, 5]. Each desert-inhabiting plant has its own complex of strategies that enables it to persist in desert habitats [6]. Strategies for improving the growth and development of arid region plant species have been investigated for many years. Treated seeds with chemical compound usually would exhibit rapid germination when absorbing water under field conditions [7].

Calotropis is a genus of flowering plants in the dogbane family, Apocynaceae. They are commonly known as milkweeds because of the latex they produce. Calotropis species are considered common weeds in some parts of the world. The flowers are fragrant and are often used in making floral tassels in some mainland Southeast Asian cultures. Calotropis persica is growing in tropical region only. Iran is a country in the mid-latitude belt of arid and semiarid regions of the Earth. Approximately 60% of Iran is classified as arid and semiarid [8]. Based on results, the seed of full ripening fruits with scarification had the highest germination percent. [9] investigated the effects of salt stress and prime on germination improvement and seedling growth of Calotropis procera L. seeds and the results showed that priming improved the seedling characteristics in all samples, especially in −0.05 MPa, but a decrease with decrease in osmotic potential. The work in [10] studied the effect of temperature, light, pretreatment, and storage on seed germination of Rhodomyrtus tomentosa and their result showed that light significantly improved germination of fresh seeds but storage decreased the light-sensitivity of germination. Soaking for 24 hours in 250–600 mg L −1 gibberellic acid, 5–20% potassium nitrate, or 10% hydrogen peroxide solution increased seed germination. Calotropis sp. is an important economic plant used for drug and other purposes. The purpose of this study was to develop methods to increase germination percentage, shorten germination time, provide more rate germination, and result in more efficient seed propagation techniques for C. persica seeds.

2. Material and Methods

Seeds of C. persica were collected from Jiroft arid regions in southern Iran in 2013. A preliminary germination test was performed and low germination percentage was obtained. To solve this problem, we implemented an experiment with a randomized complete design. Before the start of experiment, seeds were surface sterilized in 1% sodium hypochlorite solution for 5 min, then rinsed with sterilized water, and air-dried for 28 h before putting in petri dishes. Treatments included pretreatment with KNO₃ (0.1 and 0.3 percent) for 48 hours, acetylsalicylic acid to the moisture in the petri dish (100, 200 and 100 mg L −1 ), prechilling (4 degrees centigrade for 10 days), hot water (70°C) for 5 min, ethereal sulfuric acid (60%) for 5 and 10 min, thiourea with concentrations of 0.1% and 0.3%, and control treatment (irrigation with distilled water). The seeds were placed on top of Whatman paper number 1 within 10 cm petri dishes containing 10 mL distilled water. Counting number of germinating seeds began from the first day and was done till the end of the experiment (19 days). Germination percentage was recorded daily during the study period. Rate of germination was estimated using modified Timpson’s index of germination velocity [11]. Mean germination time (MGT) was calculated to assess the rate of germination [12]:

where is the number of seeds which in day grow, the total number of seeds grown, and the number of days from the date of germination and the germination rate index was obtained by reversing MGT at the end of this period final germination percentage was recorded. There are no outliers normality of data was checked and nonnormal data transformed by arc sin to verification of this hypothesis arc sin transformation was used for germination percentage before analysis [13]. Experimental data was analyzed by SPSS 17.0 to analyze the data and Duncan’s test at 5% level was used to compare the means.

3. Results

The results of ANOVA (Table 1) showed that there are significant differences (at 1% level) between effective treatments on germination characteristics and the different treatments resulted in significant differences among germination properties (Table 1).

The results of this research showed that germination percentage of C. persica increases due to application of KNO₃ in different concentrations and acetylsalicylic acid 100 and 200 mg L −1 and decreased germination percentage due to application of hot water for 5 min, prechilling for 10 days, sulfuric acid 5 and 10 min, and thiourea 0.3%. Acetylsalicylic acid 300 mg L −1 and thiourea 0.1% have the same effect on germination percentage in comparison to control treatment. The increased germination percentage by KNO₃ 0.1, 0.2, and 0.3% and acetylsalicylic acid 100 mg L −1 was significant (Figure 1).

The seed germination rates of C. persica increased significantly when KNO₃ 0.1% was used. Acetylsalicylic acid 200 and 300 mg L −1 and thiourea 0.3% increased seed germination rate, but this increase was not significant. However, the germination rate was decreased when hot water for 5 min, prechilling for 10 days, sulfuric acid for 5 and 10 min, KNO₃ 0.2 and 0.3%, acetylsalicylic acid 100 mg L −1 , and thiourea 0.3% were used (Figure 2).

Mean germination time of C. persica decreased by using KNO₃ 0.1% but this difference was not significant. In seeds of C. persica, all treatments, except for KNO₃ 0.1%, caused increase in mean germination time (Figure 3).

4. Discussion and Conclusion

According to the obtained results, KNO₃ 0.1% and acetylsalicylic acid 100 mg L −1 were the most effective treatments for improvement of seed germination properties in C. persica plant species. In a comparison of the two mentioned treatments, KNO₃ 0.1% treatment is the best. This technique has become a common seed treatment that can increase rate, percentage, and uniformity of germination or seedling emergence, mainly under unfavorable environmental conditions. Rapid seed germination and stand establishment are critical factors for crop production under stress conditions. Hot water for 5 min and prechilling for 10 days did not show positive effect on germination improvement. The study result of scarification of seeds of Acacia angustissima showed that seeds soaking in hot water cause seed germination induction but increasing duration of seed contact with hot water leads to decline of seed germination percentage [14]. In a research it is shown that prechilling for 10 days had a positive effect on germination rate and mean germination time of both medicinal species of Foeniculum vulgare and Cuscuta epithymum but germination percentage decreased due to application of prechilling [15]. Sulfuric acid for 5 and 10 min did not have positive effect on seed germination of C. persica as a result, seed treatment with sulfuric acid cannot improve seed germination. This result demonstrated that above mentioned treatment had the destructive effect on embryo. It is notified that increasing in doses of sulfuric acid caused germination improvement and suggested chemical scarification in concentrated H₂SO₄ for 2 hours [16]. In this research thiourea did not have a positive effect on germination improvement of C. persica. Acetylsalicylic acid 100 mg L −1 improved mean germination time and germination percentage in comparison with amount of 200 and 300 mg L −1 .

In this research, KNO₃ 0.1% is recognized as the best treatment for improvement seed germination properties of C. persica. Similar results were reported in previous studies for the species of Citrullus colocynthis [17], Foeniculum vulgare and Cuscuta epithymum [15], Hypericum aviculariifolium [18], and Avena fatua [19]. According these results, KNO₃ 0.1% treatment is suggested for improvement of C. persica germination and this treatment is proper for propagation of studied species. Positive effect of KNO₃ could be due to its role in balancing hormonal portion within seed which in turn results in germination inhibitors ratio like ABA. (abscisic acid). Virtually all of the cellular and metabolic events that are known to occur before the completion of germination of nondormant seeds also occur in imbibed dormant seeds indeed, the metabolic activities of the latter are frequently only subtly different from those of the former [20]. The seeds of most Mediterranean and desert species have dormancy characteristics or structural properties that prevent immediate germination of at least a proportion of the seeds [21–24]. The results obtained will be useful in carrying out tree improvement and plantings of C. persica trees for fuel wood, local medicine, and industrial production. Rapid seedling growth is also essential for reclamation of desert. This information could ultimately help in the sustainable development of the arid zones.

Conflict of Interests

The authors declare that there is no conflict of interests regarding the publication of this paper.

Acknowledgments

This research was supported by Iranian Revolutionary Guards Navy. The authors would like to thank Hamid Reza Ahmadinia Ph.D. student of Fishery in Gorgan University of Agriculture and Natural Resources for his help with the creation of this work.

References

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18. C. Çirak, K. Kevseroǧlu, and A. K. Ayan, “Breaking of seed dormancy in a Turkish endemic Hypericum species: Hypericum aviculariifolium subsp. depilatum var. depilatum by light and some pre-soaking treatments,” Journal of Arid Environments, vol. 68, no. 1, pp. 159–164, 2007. View at: Publisher Site | Google Scholar
19. J. R. Hilton, “The influence of light and potassium nitrate on the dormancy and germination of Avena fatua L. (wild oat) seed and its ecological significance,” New Phytologist, vol. 96, no. 1, pp. 31–34, 1984. View at: Publisher Site | Google Scholar
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24. D. T. Bell, D. P. Rokich, C. J. McChesney, and J. A. Plummer, “Effects of temperature, light and gibberellic acid on the germination of seeds of 43 species native to Western Australia,” Journal of Vegetation Science, vol. 6, no. 6, pp. 797–806, 1995. View at: Publisher Site | Google Scholar

Copyright

Copyright © 2014 Asghar Farajollahi et al. This is an open access article distributed under the Creative Commons Attribution License, which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

Sulfur bacterium

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Sulfur bacterium, plural Sulfur Bacteria, any of a diverse group of microorganisms capable of metabolizing sulfur and its compounds and important in the sulfur cycle (q.v.) in nature. Some of the common sulfur substances that are used by these bacteria as an energy source are hydrogen sulfide (H₂S), sulfur, and thiosulfate (S₂O₃ 2- ). The final product of sulfur oxidation is sulfate (SO₄ 2- ).

Thiobacillus, widespread in marine and terrestrial habitats, oxidizes sulfur, producing sulfates useful to plants in deep ground deposits it generates sulfuric acid, which dissolves metals in mines but also corrodes concrete and steel. Desulfovibrio desulficans reduces sulfates in waterlogged soils and sewage to hydrogen sulfide, a gas with the rotten egg odour so common to such places. Thiothrix, common in sulfur springs and in sewage, and Sulfolobus, confined to sulfur-rich hot springs, transform hydrogen sulfide to elemental sulfur.

Many species in the families Chromatiaceae (purple sulfur bacteria) and Chlorobiaceae (green sulfur bacteria) utilize energy from light in an oxygen-free environment to transform sulfur and its compounds to sulfates.

Sulfuric acid

Table (1.2.1): physical properties1

Sulfuric acid is a very important chemical commodity, and indeed, a nation’s sulfuric acid production is a good indicator of its industrial strength.

It is used as electrolyte in lead-acid batteries (accumulators) .

It is important in the production of fertilizers such as ammonium sulfate (sulfate of ammonia), (NH4)2SO4, and superphosphate, Ca(H2PO4)2, which is formed when rock phosphate is treated with sulfuric acid.

It is used to remove oxides from iron and steel before galvanising or electroplating .

Concentrated sulfuric acid is used as a dehydrating agent, that is, to remove water, since it has a tendency to form hydrates such as H2SO4.H2O, H2SO4.2H2O.

Sulfuric acid is used in the production of nitroglycerine, an inorganic ester & organic nitrate, which is used as an explosive.

It is used in petroleum refining to wash impurities out of gasoline and other refinery products.

It is used in manufacturing of hydrochloric acid, nitric acid, phosphoric acid, ether, plastics, metal sulfates, cellophane, dyes, drugs, perfumes, disinfectants and even glue.1

This chart shows the distribution of using sulfuric acid

Figure (1.3.1): Sulfuric Acid Distribution.1

Specification of raw materials

Yellow colored lumps, crystals, powder, or formed shape

Lumps 75-115 lbs./ft3 Powder 33-80 lbs./ft3

388.36 K, 115.21 °C, 239.38 °F

717.8 K, 444.6 °C, 832.3 °F

Table (1.3.1): Physical & Chemical Properties of Sulfur.1

2.1 History and Current processes

The discovery of sulfuric acid is credited to the 8th century chemist and alchemist, Jabir ibn Hayyan (Geber). The acid was later studied by 9th century Persian physician and alchemist Ibn Zakariya al-Razi (Rhazes), who obtained the substance by dry distillation of minerals including iron(II) sulfate heptahydrate, FeSO4·7H2O, and copper(II) sulfate pentahydrate, CuSO4·5H2O. When heated, these compounds decompose to iron(II) oxide and copper(II) oxide, respectively, giving off water and sulfur trioxide, which combine to produce a dilute solution of sulfuric acid. 1

This method was popularized in Europe through translations of Arabic and Persian treatises, as well as books by European alchemists, such as the 13th-century German Albertus Magnus.1

There are two major processes (lead chamber and contact) for production of sulfuric acid and it is available commercially in a number of grades and concentrations. The lead chamber process, the older of the two processes, is used to produce much of the acid used to make fertilizers it produces a relatively dilute acid (62%-78% H2SO4). The contact process produces a purer, more concentrated acid but requires purer raw materials and the use of expensive catalysts. n both processes sulfur dioxide is oxidized and dissolved in water. The sulfur dioxide is obtained by burning sulfur, by burning pyrites (iron sulfides), by roasting nonferrous sulfide ores preparatory to smelting, or by burning hydrogen sulfide gas. Some sulfuric acid is also made from ferrous sulfate waste solutions from pickling iron and steel and from waste acid sludge from oil refineries. 1

This table shows the production rates of sulfuric acid (in metric tones) in some countries at different years.

Production of sulfuric acid in metric tones

Table (2.2.1): Production Rates of Sulfuric Acid.3

This table shows the production and sales amounts of sulfuric acid and the consumption rate of sulfur in Jordan from 2000 to 2005, these amounts in (ton/year).

Table (2.2.2)Jordan Production, Sales and Raw Material Consumption.5

2.3 Prices trends of the raw material and product

The global sulfuric acid market experienced an unprecedented rise and fall in pricing between fall 2007 and spring 2009. Consumption of sulfuric acid for fertilizers fell steeply in the second half of 2008 due to the collapse in the global economy. The second half of 2009 is expected to experience almost flat to slightly positive growth, anticipating the improvement in market conditions in 2010. Trade is expected to fall globally, except for Southeast Asia, which would continue to depend on imports. As of early spring 2009, the market is continuing to deteriorate as the supply shortage situation has been replaced by product oversupply in almost all regions.

And the world sulfuric acid supply trends are shown in the following chart.

Figure (2.3.1): World Sulfuric Acid Supply.2

Sulfuric acid is an important raw material used in many industrial processes, such as phosphate fertilizer production and to a much lesser extent for nitrogen and potassium fertilizers, sulfuric acid is produced by catalytic oxidation of sulfur dioxide to sulfur trioxide, which is subsequently absorbed in water to form sulfuric acid.

There are no major variations of commercial interests on this mentioned chemistry. There are alternatives as to source of Sulfur dioxide and method of conversion to sulfur trioxide. The two most common methods for the conversion of sulfur dioxide to sulfuric

This is an old process and was introduced in Europe in near the middle of 18th century, it’s used to produce much of the acid used to make fertilizers it produces a- relatively dilute acid (62%-78% H2SO4).The classic lead chamber process consists of three stages: Glover tower, lead chambers and Guy-Lussac Tower.

In this method hot sulfuric dioxide gas enters the bottom of the reactor called a Glover tower where it is washed with nitrous vitriol (sulfuric acid with nitric oxide, NO, and nitrogen dioxide, NO2, dissolved in it) and mixed with nitric oxide and nitrogen dioxide gases.

The Glover tower serves two functions: concentration of the chamber acid and stripping of nitrogen oxides from the liquid to the gas. Concentration of the chamber acid (62% to 68% H2SO4) is achieved by the hot gases entering the tower which evaporate water from the acid.

Some of the sulfur dioxide is oxidized to sulfur trioxide and dissolved in the acid wash to form tower acid or Glover acid (about 78% H2SO4). The dissolved nitrogen oxides are stripped from the acid and carried with the gas out of the Glover tower into the lead chambers.

From the acid tower a mixture of gases (including sulfur dioxide and trioxide, nitrogen oxides, nitrogen, oxygen, and steam) is transferred to a lead-lined chamber where it is reacted with more water.

Sulfuric acid is formed by a complex series of reactions it condenses on the walls and collects on the floor of the chamber. There may be from three to twelve chambers in a series. The acid produced in the chambers, often called chamber acid or fertilizer acid, contains 62% to 68% H2SO4.

After the gases have passed through the chambers they are passed into a reactor called the Gay-Lussac tower where they are washed with cooled concentrated acid (from the acid tower) the nitrogen oxides and unreacted sulfur dioxide dissolve in the acid to form the nitrous vitriol used in the acid tower. Remaining waste gases are usually discharged into the atmosphere.

Product acid at a concentration of 78% H2SO4 is drawn from the cooled acid stream that is circulated from the Glover tower to the Guy-Lussac tower. Nitrogen losses are made up with nitric acid which is added to the Glover tower.

The major disadvantage includes the limitations in throughput, quality and concentration of the acid produced, also the environmental pollution.

Figure (3.1.1): Typical process flow sheet for the lead Chamber.

Because of economic reasons Contact plants are widely used compared to the lead plants, they are classified according to the raw materials charged to them: elemental Sulfur burning, spent sulfuric acid and hydrogen sulfide burning, and metal sulfide ores and smelter gas burning. The contributions from these plants to the total acid production are 81, 8, and 11 percent, respectively.

The contact process incorporates three basic operations (stages), each of which corresponds to a distinct chemical reaction.

First, elemental sulfur is received in a solid form containing various impurities. The sulfur is melted in the sulfur melter in the presence of hydrated lime which neutralizes any acidity present in the sulfur. This neutralization prevents problems of acid corrosion which would otherwise be encountered.

Heat for the melting of the sulfur is supplied from steam coils. The molten sulfur is kept agitated to improve heat transfer, to prevent solids settling on the bottom of the sulfur pits and to prevent a crust forming on top. The “dirty sulfur” is filtered to remove impurities present and after filtering is transferred to the “clean sulfur” pit where it is kept molten until it is pumped to the burner.

Molten sulfur at a temperature of 130°C is sprayed into the burner in the presence of warm, dry air. The sulfur burns, forming sulfur dioxide

S + O2 € € € € SO2 ∆H = -300 kJ mol-1

The resulting sulfur dioxide is fed to a process unit called a converter, where it is catalytically oxidized to sulfur trioxide (SO3):

2SO2 + O2 € € € € 2SO3 ΔH = -100 kJ mol-1

It’s apparent that the equation gives a decrease in volume this reaction would be aided by pressure. High conversions are however, obtainable with catalysts at 400 to 500oC with a small excess of oxygen and the use of pressure.

The available methods to maximize the formation of SO3:

As this is an exothermic process, a decrease in temperature by removal of the heat will favour the formation of SO3.

Increased oxygen concentration.

SO3 removal (as in the case of the double absorption process).

Catalyst selection, to reduce the working temperature (equilibrium).

In the contact processes, the sulfur dioxide is converted to sulfur trioxide by the use of metal oxide catalyst, the characteristics of the used catalyst are:

Porous carrier having large surface area, controlled pore size and resistance to process gases at high temperature in pellet form if used in fixed bed and

powdered form if used for fluidized bed. Ex- Alumina, silica gel, zeolites.

Preparations are generally kept secret for the competitive reasons but they usually

consist of adding water soluble compounds to gels or porous substrates and firing

at temperature below the sintering point.

Alkali and/or metallic compounds added in trace amounts to enhance the activity

A catalyst, vanadium pentoxide (V2O5) is used to increase the reaction rate because it’s relatively immune to poisons, also because of its low initial investment and only 5% replacement per year. It is only effective above its melting point of 400 °C. The greatest conversion of SO2 to SO3 is reached by passing the gas over several catalyst beds, cooling the gas between each pass so that the reaction temperature remains between 400 and 500 °C. As can be seen the figure.

The disadvantages of using the V2O5 catalyst are that it must use dilute SO2 input (7-10%), as a catalyst it is less active and requires high oxygen or sulfur dioxide to give economic conversions also it requires larger converters and thus higher initial investment.

Finally, the sulfur trioxide is absorbed in to very concentrated sulfuric acid (a 98-99 percent solution of H2SO4 in water), This operation takes place in the absorbing tower where the gas travels up through the tower, counter-current to the acid falling from the top of the tower producing a thick fuming liquid called oleum, the oleum is mixed carefully with water to avoid producing fine mist of sulfuric acid that is difficult to condense and could escape to pollute the air, the sulfur trioxide in the oleum reacts with the water as follows:

SO3 + H2O € € € € € € € € H2SO4 ∆H = -200 kJ mol-1

It is clear that the reaction is exothermic and the absorbing sulfuric acid has to be cooled continuously the heat is available at a relatively low temperature and is not worth recovering.

The efficiency of the absorption step is related to :

The H2SO4 concentration of the absorbing liquid. (98.5 – 99.5%).

The temperature range of the liquid (normally 70 -120 0C).

The technique of the acid distribution.

The raw gas humidity (mist passes the absorption equipment).

The temperature of incoming gas.

The co-current or countercurrent character of the gas stream in the absorbing liquid.

Main disadvantages of the contact process are that concentrated acid (98%) of high purity can be produced directly and that compact plants of quite high capacity have now become rather common place.

The contact process can be applied in different techniques three of those techniques are described in the following sections

3.2.1 Single contact / single absorption process

After purification and drying, the SO2 is converted to SO3 using a series of four catalyst beds, containing alkali and V2O5. Afterwards, the SO3 is absorbed in concentrated sulfuric acid and, if necessary, an oleum absorber is installed upstream. SO3 reacts with the water contained in the absorber acid to yield H2SO4. The absorber acid is kept at the desired concentration of approximately 99% w/w by addition of water or dilute H2SO4.

The single contact/single absorption process is generally used for gases with an SO2 Content from 3 – 6 %. New single contact plants are built only for inlet gases with substantial fluctuation of the SO2 content.

The investment cost of this technique is low compared to the investment cost of double contact plants.

Figure (3.2.1.1): Typical process flow sheet for a single catalysis plant.

3.2.2 Double Contact/ Double Absorption Process

The double contact process was implemented to develop the single contact/single absorption process. In this process a primary SO2 conversion of 85 – 95 % is achieved in the first catalysis stage of the converter before entry into an intermediate absorber, depending on the arrangement of the converter beds and the contact time.

What makes the double contact/double absorption process more advantageous is that its ability to feed gases with higher SO2 concentrations than would be possible with the single catalysis process. Which leads to smaller gas volumes and therefore smaller equipment with comparable production capacities.

This results in a considerably higher conversion rate, if the residual gas is passed through the following converter beds (usually one or two). The SO3 which is formed in the second catalysis stage is absorbed in the final absorber.

In general the process uses gases with an SO2 content of 10 t o11 %. The inlet gas temperature is about 4000C. Gases with lower temperatures require reheating from 50 to 4000C. This is usually carried out with recovered heats from the conversion process.

Operating the double contact process at an elevated pressure of 5 bar increases the conversion rate by shifting the conversion equilibrium and favouring the formation of SO3.

The disadvantages are higher electricity consumption and, at the same time, less steam production. Higher NOx emissions are caused by higher sulfur combustion temperatures (18000C), but savings of 10 -17 % on investment costs are gained.

Figure 3.2.2.1: Typical process flow sheet for a sulfur burning double catalysis

3.2.3 Wet catalysis process

The wet catalysis process is applicable to wet SO2 gases. The potential for the formation of sulfuric acid mist might require tail gas treatment.

Wet SO2 gases (eg. from the burning of H2S gases or from the catalytic conversion of H2S gases) are directly supplied into the contact tower without previous drying. SO3 formed by the catalytic conversion immediately reacts with the moisture of the gases, thereby forming the acetic acid. The sulfuric acid is condensed in a condenser installed after the contact tower.

Sulfuric Acid Production By Lead Chamber process

Sulfuric Acid Production By single contact/single absorption process

Sulfuric Acid Production By double contact/double absorption process

Sulfuric Acid Production By Wet Catalysis process

Health and safety hazards involved

Less safe, waste gases are discharged to the atmosphere

Less amount of SO3 is absorbed so the rest is discharged to the atmosphere

A larger amount of SO3 is absorbed

A larger amount of SO3 is absorbed

Melted sulfur, H2O, O2, SO2, SO3.

Waste products and by products

Exhaust gases are discharged to the atmosphere

Large amounts of SO2 gas are discharged to the atmosphere

Less amounts of SO2 gas are discharged to the atmosphere, less heat released after each successive catalyst bed.

A larger amount of SO3 is absorbed

Acid Tower (Glover Tower), Lead Chambers, Reactor (Gay-lussac Tower)

Air dryer, burner, waste heat boiler, converter, single absorption column.

Air dryer, burner, waste heat boiler, converter, intermediate and external absorption column.

Contents

Grades of sulfuric acid Edit

Although nearly 100% sulfuric acid solutions can be made, the subsequent loss of SO
₃ at the boiling point brings the concentration to 98.3% acid. The 98.3% grade is more stable in storage, and is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes. Some common concentrations are: [12] [13]

Mass fraction H2SO4	Density (kg/L)	Concentration (mol/L)	Common name
<29%	1.00-1.25	<4.2	diluted sulfuric acid
29–32%	1.25–1.28	4.2–5.0	battery acid (used in lead–acid batteries)
62–70%	1.52–1.60	9.6–11.5	chamber acid fertilizer acid
78–80%	1.70–1.73	13.5–14.0	tower acid Glover acid
93.2%	1.83	17.4	66 °Bé ("66-degree Baumé") acid
98.3%	1.84	18.4	concentrated sulfuric acid

"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in the lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid) and tower acid being the acid recovered from the bottom of the Glover tower. [12] [13] They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid (the modern equivalent of chamber acid, used in many titrations), is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher. [13]

Sulfuric acid reacts with its anhydride, SO
₃ , to form H
₂ S
₂ O
₇ , called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less commonly, Nordhausen acid. Concentrations of oleum are either expressed in terms of % SO
₃ (called % oleum) or as % H
₂ SO
₄ (the amount made if H
₂ O were added) common concentrations are 40% oleum (109% H
₂ SO
₄ ) and 65% oleum (114.6% H
₂ SO
₄ ). Pure H
₂ S
₂ O
₇ is a solid with melting point of 36 °C.

Pure sulfuric acid has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, [14] and 98% sulfuric acid has a <1 mmHg vapor pressure at 40 °C. [15]

Pure sulfuric acid is a viscous clear liquid, like oil, and this explains the old name of the acid ('oil of vitriol').

Commercial sulfuric acid is sold in several different purity grades. Technical grade H
₂ SO
₄ is impure and often colored, but is suitable for making fertilizer. Pure grades, such as USP grade, are used for making pharmaceuticals and dyestuffs. Analytical grades are also available.

Nine hydrates are known, but four of them were confirmed to be tetrahydrate (H₂SO₄·4H₂O), hemihexahydrate (H₂SO₄· 6 + 1 ⁄ 2 H₂O) and octahydrate (H₂SO₄·8H₂O).

Polarity and conductivity Edit

Equilibrium of anhydrous sulfuric acid [16]

Species	mMol/kg
HSO − 4	15.0
H 3 SO + 4	11.3
H 3 O +	8.0
HS 2 O − 7	4.4
H 2 S 2 O 7	3.6
H 2 O	0.1

Anhydrous H
₂ SO
₄ is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis. [16]

The equilibrium constant for the autoprotolysis is [16]

The comparable equilibrium constant for water, K_w is 10 −14 , a factor of 10 10 (10 billion) smaller.

In spite of the viscosity of the acid, the effective conductivities of the H
₃ SO +
₄ and HSO −
₄ ions are high due to an intramolecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions.

Reaction with water and dehydrating property Edit

Because the hydration reaction of sulfuric acid is highly exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid. [17] Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent. This reaction is best thought of as the formation of hydronium ions:

Because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent. Concentrated sulfuric acid has a very powerful dehydrating property, removing water (H₂O) from other chemical compounds including sugar and other carbohydrates and producing carbon, heat, and steam.

In the laboratory, this is often demonstrated by mixing table sugar (sucrose) into sulfuric acid. The sugar changes from white to dark brown and then to black as carbon is formed. A rigid column of black, porous carbon will emerge as well. The carbon will smell strongly of caramel due to the heat generated. [19]

C 12 H 22 O 11 ⏞ sucrose → H 2 SO 4 12 C (black graphitic foam) + 11 H 2 O (g,l) >> ^< ext> >]>> >>>>+>_>> Similarly, mixing starch into concentrated sulfuric acid will give elemental carbon and water as absorbed by the sulfuric acid (which becomes slightly diluted). The effect of this can be seen when concentrated sulfuric acid is spilled on paper which is composed of cellulose the cellulose reacts to give a burnt appearance, the carbon appears much as soot would in a fire. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric. ( C 6 H 10 O 5 ) n ⏞ polysaccharide → H 2 SO 4 6 n C + 5 n H 2 O >>> ^> >]>> 6n>+5n>> The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystal is changed into white powder as water is removed. CuSO 4 ⋅ 5 H 2 O (blue crystal) ⏞ copper(II) sulfate hydrate → H 2 SO 4 CuSO 4 (white powder) ⏞ Anhydrous copper(II) sulfate + 5 H 2 O >>>> ^> >]>> overbrace >>>> ^>+>> Acid-base properties Edit As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, the blue copper salt copper(II) sulfate, commonly used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid: Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH 3 COOH , and forms sodium bisulfate: Similarly, reacting sulfuric acid with potassium nitrate can be used to produce nitric acid and a precipitate of potassium bisulfate. When combined with nitric acid, sulfuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO + 2 , which is important in nitration reactions involving electrophilic aromatic substitution. This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols. When allowed to react with superacids, sulfuric acid can act as a base and be protonated, forming the [H3SO4] + ion. Salt of [H3SO4] + have been prepared using the following reaction in liquid HF: The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply HF/SbF5, however, have met with failure, as pure sulfuric acid undergoes self-ionization to give [H3O] + ions, which prevents the conversion of H2SO4 to [H3SO4] + by the HF/SbF5 system: [20] Reactions with metals Edit Even dilute sulfuric acid reacts with many metals via a single displacement reaction as with other typical acids, producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series) such as iron, aluminium, zinc, manganese, magnesium, and nickel. Concentrated sulfuric acid can serve as an oxidizing agent, releasing sulfur dioxide: [6] Lead and tungsten, however, are resistant to sulfuric acid. Reactions with carbon Edit Hot concentrated sulfuric acid oxidizes carbon [21] (as bituminous coal) and sulfur. Reaction with sodium chloride Edit Electrophilic aromatic substitution Edit Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids: [22] Pure sulfuric acid is not encountered naturally on Earth in anhydrous form, due to its great affinity for water. Dilute sulfuric acid is a constituent of acid rain, which is formed by atmospheric oxidation of sulfur dioxide in the presence of water – i.e., oxidation of sulfurous acid. When sulfur-containing fuels such as coal or oil are burned, sulfur dioxide is the main byproduct (besides the chief products carbon oxides and water). Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as iron sulfide. The resulting water can be highly acidic and is called acid mine drainage (AMD) or acid rock drainage (ARD). This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly colored, toxic solutions. The oxidation of pyrite (iron sulfide) by molecular oxygen produces iron(II), or Fe 2+ : The Fe 2+ can be further oxidized to Fe 3+ : The Fe 3+ produced can be precipitated as the hydroxide or hydrous iron oxide: The iron(III) ion ("ferric iron") can also oxidize pyrite: When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process. ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals. Sulfuric acid is used as a defense by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales) concentrates sulfuric acid in cell vacuoles. [23] Stratospheric aerosol Edit In the stratosphere, the atmosphere's second layer that is generally between 10 and 50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical: [24] Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in the stratospheric aerosol layer. [24] Extraterrestrial sulfuric acid Edit The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain. [25] Jupiter's moon Europa is also thought to have an atmosphere containing sulfuric acid hydrates. [26] Sulfuric acid is produced from sulfur, oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA). Contact process Edit In the first step, sulfur is burned to produce sulfur dioxide. The sulfur dioxide is oxidized to sulfur trioxide by oxygen in the presence of a vanadium(V) oxide catalyst. This reaction is reversible and the formation of the sulfur trioxide is exothermic. The sulfur trioxide is absorbed into 97–98% H 2 SO 4 to form oleum ( H 2 S 2 O 7 ), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid. Directly dissolving SO 3 in water is not practiced. Wet sulfuric acid process Edit In the first step, sulfur is burned to produce sulfur dioxide: or, alternatively, hydrogen sulfide ( H 2 S ) gas is incinerated to SO 2 gas: The sulfur dioxide then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst. 2 SO 2 + O 2 ⇌ 2 SO 3 (−99 kJ/mol) (reaction is reversible) The sulfur trioxide is hydrated into sulfuric acid H 2 SO 4 : The last step is the condensation of the sulfuric acid to liquid 97–98% H 2 SO 4 : Other methods Edit A method that is the less well-known is the metabisulfite method, in which metabisulfite is placed at the bottom of a beaker and 12.6 molar concentration hydrochloric acid is added. The resulting gas is bubbled through nitric acid, which will release brown/red vapors of nitrogen dioxide as the reaction proceeds. The completion of the reaction is indicated by the ceasing of the fumes. This method does not produce an inseparable mist, which is quite convenient. In principle, sulfuric acid can be produced in the laboratory by burning sulfur in air followed by dissolving the resulting sulfur dioxide in a hydrogen peroxide solution. [ citation needed ] Alternatively, dissolving sulfur dioxide in an aqueous solution of an oxidizing metal salt such as copper (II) or iron (III) chloride: Two less well-known laboratory methods of producing sulfuric acid, albeit in dilute form and requiring some extra effort in purification. A solution of copper (II) sulfate can be electrolyzed with a copper cathode and platinum/graphite anode to give spongy copper at cathode and evolution of oxygen gas at the anode, the solution of dilute sulfuric acid indicates completion of the reaction when it turns from blue to clear (production of hydrogen at cathode is another sign): More costly, dangerous, and troublesome yet novel is the electrobromine method, which employs a mixture of sulfur, water, and hydrobromic acid as the electrolytic solution. The sulfur is pushed to bottom of container under the acid solution, then the copper cathode and platinum/graphite anode are used with the cathode near the surface and the anode is positioned at bottom of the electrolyte to apply the current. This may take longer and emits toxic bromine/sulfur bromide vapors, but the reactant acid is recyclable, overall only the sulfur and water are converted to sulfuric acid (omitting losses of acid as vapors): 2HBr → H2 + Br2 (electrolysis of aqueous hydrogen bromide) Br2 + Br − ↔ Br3 − (initial tribromide production, eventually reverses as Br − depletes) 2S + Br2 → S2Br2 (bromine reacts with sulfur to form disulfur dibromide) S2Br2 + 8H2O + 5Br2 → 2H2SO4 + 12HBr (oxidation and hydration of disulfur dibromide) Prior to 1900, most sulfuric acid was manufactured by the lead chamber process. [27] As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants. In the early to mid nineteenth century "vitriol" plants existed, among other places, in Prestonpans in Scotland, Shropshire and the Lagan Valley in County Antrim Ireland where it was used as a bleach for linen. Early bleaching of linen was done using lactic acid from sour milk but this was a slow process and the use of vitriol sped up the bleaching process. [28] Sulfuric acid is a very important commodity chemical, and indeed, a nation's sulfuric acid production is a good indicator of its industrial strength. [8] World production in the year 2004 was about 180 million tonnes, with the following geographic distribution: Asia 35%, North America (including Mexico) 24%, Africa 11%, Western Europe 10%, Eastern Europe and Russia 10%, Australia and Oceania 7%, South America 7%. [29] Most of this amount (≈60%) is consumed for fertilizers, particularly superphosphates, ammonium phosphate and ammonium sulfates. About 20% is used in chemical industry for production of detergents, synthetic resins, dyestuffs, pharmaceuticals, petroleum catalysts, insecticides and antifreeze, as well as in various processes such as oil well acidicizing, aluminium reduction, paper sizing, and water treatment. About 6% of uses are related to pigments and include paints, enamels, printing inks, coated fabrics and paper, while the rest is dispersed into a multitude of applications such as production of explosives, cellophane, acetate and viscose textiles, lubricants, non-ferrous metals, and batteries. [30] Industrial production of chemicals Edit The major use for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilizers. In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite, though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate, hydrogen fluoride (HF) and phosphoric acid. The HF is removed as hydrofluoric acid. The overall process can be represented as: Ca 5 F ( PO 4 ) 3 fluorapatite + 5 H 2 SO 4 + 10 H 2 O ⟶ 5 CaSO 4 ⋅ 2 H 2 O calcium sulfate + HF + 3 H 3 PO 4 >++10H2O-> Ammonium sulfate, an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry. Another important use for sulfuric acid is for the manufacture of aluminium sulfate, also known as paper maker's alum. This can react with small amounts of soap on paper pulp fibers to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making aluminium hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminium sulfate is made by reacting bauxite with sulfuric acid: Sulfuric acid is also important in the manufacture of dyestuffs solutions. Sulfur–iodine cycle Edit The sulfur–iodine cycle is a series of thermo-chemical processes possibly usable to produce hydrogen from water. It consists of three chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen. 2 I 2 + 2 SO 2 + 4 H 2 O → 4 HI + 2 H 2 SO 4 (120 °C, Bunsen reaction)2 H 2 SO 4 → 2 SO 2 + 2 H 2 O + O 2 (830 °C)4 HI → 2 I 2 + 2 H 2 (320 °C) The compounds of sulfur and iodine are recovered and reused, hence the consideration of the process as a cycle. This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied. The sulfur–iodine cycle has been proposed as a way to supply hydrogen for a hydrogen-based economy. It is an alternative to electrolysis, and does not require hydrocarbons like current methods of steam reforming. But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it. The sulfur–iodine cycle is currently being researched as a feasible method of obtaining hydrogen, but the concentrated, corrosive acid at high temperatures poses currently insurmountable safety hazards if the process were built on a large scale. [31] [32] Industrial cleaning agent Edit Sulfuric acid is used in large quantities by the iron and steelmaking industry to remove oxidation, rust, and scaling from rolled sheet and billets prior to sale to the automobile and major appliances industry. [ citation needed ] Used acid is often recycled using a spent acid regeneration (SAR) plant. These plants combust spent acid [ clarification needed ] with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide ( SO 2 ) and sulfur trioxide ( SO 3 ) which are then used to manufacture "new" sulfuric acid. SAR plants are common additions to metal smelting plants, oil refineries, and other industries where sulfuric acid is consumed in bulk, as operating a SAR plant is much cheaper than the recurring costs of spent acid disposal and new acid purchases. Hydrogen peroxide ( H 2 O 2 ) can be added to sulfuric acid to produce piranha solution, a powerful but very toxic cleaning solution with which substrate surfaces can be cleaned. Piranha solution is typically used in the microelectronics industry, and also in laboratory settings to clean glassware. Catalyst Edit Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanone oxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H 2 SO 4 is used in petroleum refining, for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane, a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also often used as a dehydrating or oxidizing agent in industrial reactions, such as the dehydration of various sugars to form solid carbon. Electrolyte Edit Sulfuric acid acts as the electrolyte in lead–acid batteries (lead-acid accumulator): Domestic uses Edit Sulfuric acid at high concentrations is frequently the major ingredient in acidic drain cleaners [11] which are used to remove grease, hair, tissue paper, etc. Similar to their alkaline versions, such drain openers can dissolve fats and proteins via hydrolysis. Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned. The study of vitriol, a category of glassy minerals from which the acid can be derived, began in ancient times. Sumerians had a list of types of vitriol that they classified according to the substances' color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides (first century AD) and the Roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis, in the treatise Phisica et Mystica, and the Leyden papyrus X. [33] Medieval Islamic chemists like Jābir ibn Ḥayyān (died c. 806 – c. 816 AD, known in Latin as Geber), Abū Bakr al-Rāzī (865 – 925 AD, known in Latin as Rhazes), Ibn Sina (980 – 1037 AD, known in Latin as Avicenna), and Muḥammad ibn Ibrāhīm al-Watwat (1234 – 1318 AD) included vitriol in their mineral classification lists. [34] Sulfuric acid was called "oil of vitriol" by medieval European alchemists because it was prepared by roasting "green vitriol" (iron(II) sulfate) in an iron retort. The first vague allusions to it appear in the works of Vincent of Beauvais, in the Compositum de Compositis ascribed to Saint Albertus Magnus, and in pseudo-Geber's Summa perfectionis (all thirteenth century AD). [35] In the seventeenth century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO 3 ), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to SO 3 , which combines with water to produce sulfuric acid. In 1736, Joshua Ward, a London pharmacist, used this method to begin the first large-scale production of sulfuric acid. In 1746 in Birmingham, John Roebuck adapted this method to produce sulfuric acid in lead-lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the lead chamber process or "chamber process", remained the standard for sulfuric acid production for almost two centuries. [3] Sulfuric acid created by John Roebuck's process approached a 65% concentration. Later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS 2 ) was heated in air to yield iron(II) sulfate, FeSO 4 , which was oxidized by further heating in air to form iron(III) sulfate, Fe2(SO4)3, which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid. [3] In 1831, British vinegar merchant Peregrine Phillips patented the contact process, which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method. [36] Laboratory hazards Edit Sulfuric acid is capable of causing very severe burns, especially when it is at high concentrations. In common with other corrosive acids and alkali, it readily decomposes proteins and lipids through amide and ester hydrolysis upon contact with living tissues, such as skin and flesh. In addition, it exhibits a strong dehydrating property on carbohydrates, liberating extra heat and causing secondary thermal burns. [6] [7] Accordingly, it rapidly attacks the cornea and can induce permanent blindness if splashed onto eyes. If ingested, it damages internal organs irreversibly and may even be fatal. [5] Protective equipment should hence always be used when handling it. Moreover, its strong oxidizing property makes it highly corrosive to many metals and may extend its destruction on other materials. [6] Because of such reasons, damage posed by sulfuric acid is potentially more severe than that by other comparable strong acids, such as hydrochloric acid and nitric acid. Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). Solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 M are labeled "IRRITANT". However, even the normal laboratory "dilute" grade (approximately 1 M, 10%) will char paper if left in contact for a sufficient time. The standard first aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly. Dilution hazards Edit Preparation of the diluted acid can be dangerous due to the heat released in the dilution process. To avoid splattering, the concentrated acid is usually added to water and not the other way around. Water has a higher heat capacity than the acid, and so a vessel of cold water will absorb heat as acid is added. Comparison of sulfuric acid and water Physical property H2SO4 Water UnitsDensity 1.84 1.0 kg/LVolumetric heat capacity 2.54 4.18 kJ/LBoiling point 337 100 °C Also, because the acid is denser than water, it sinks to the bottom. Heat is generated at the interface between acid and water, which is at the bottom of the vessel. Acid will not boil, because of its higher boiling point. Warm water near the interface rises due to convection, which cools the interface, and prevents boiling of either acid or water. In contrast, addition of water to concentrated sulfuric acid results in a thin layer of water on top of the acid. Heat generated in this thin layer of water can boil, leading to the dispersal of a sulfuric acid aerosol or worse, an explosion. Preparation of solutions greater than 6 M (35%) in concentration is most dangerous, because the heat produced may be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (such as an ice bath) are essential. Reaction rates double for about every 10-degree Celsius increase in temperature. [37] Therefore, the reaction will become more violent as dilution proceeds, unless the mixture is given time to cool. Adding acid to warm water will cause a violent reaction. On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water. Industrial hazards Edit Sulfuric acid is non-flammable. The main occupational risks posed by this acid are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. Repeated occupational exposure to sulfuric acid mists may increase the chance of lung cancer by up to 64 percent. [38] In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m 3 : limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to vitamin B12 deficiency with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show demyelination, loss of axons and gliosis. International commerce of sulfuric acid is controlled under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988, which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances. [39] A few months ago I called my local nursery to ask if they carried ammonium sulfate. He said "Oh, you want aluminum sulfate to acidify the soil for blueberries." I cringed in horror that this advice is being dispensed so regularly. Aluminum is not known to be a nutrient for plant growth in any quantity and is actually more well-known for being a toxin. Aluminum (Al) is the most abundant metal in the earth's crust, comprising about 7% of its mass. Since many plant species are sensitive to micromolar concentrations of Al, the potential for soils to be Al toxic is considerable. Fortunately, most of the Al is bound by ligands or occurs in other nonphytotoxic forms such as aluminosilicates and precipitates. However, solubilization of this Al is enhanced by low pH and Al toxicity is a major factor limiting plant production on acid soils. Why would you want to add more of this to your soil? Aluminum doesn't acidify the soil, but aluminum becomes more readily available to plants as the soil becomes more acid. Ph is a measure of H+ ions in the soil. It has nothing to do with aluminum. In technical terms, pH is the negative logarithm of the activity of the (solvated) hydronium ion, more often expressed as the measure of the hydronium ion concentration. The exact meaning of the "p" in "pH" is disputed, but according to the Carlsberg Foundation pH stands for "power of hydrogen". Actually, what blueberries like is Magnesium. They grow well in acid soils because calcium is low, which lowers the Ca:Mg ratio and provides the plant with more access to Mg with less competition from Ca. pH is a measure of free Hydrogen ions in water. It measures Hydrogen ion concentration, H+ and OH-, and that’s all it does. One can change the soil pH with any acid or alkali. You can raise the pH with sodium hydroxide, which is lye, drain cleaner, or lower it with hydrochloric acid, for instance, but they aren’t going to give you much growth stimulus. They will probably kill the plant. A slightly acid pH of about 6 or 6.5 is ideal, because it gives just the right amount of electrical conductivity in the soil, but plants aren’t nearly as finicky about pH as they are about having the right balance of soil minerals. Rhododendrons, for instance, are supposed to require an acid soil. What they really prefer is a high Magnesium soil. Experimenters in Scotland raised the pH of soil from 5.0 to nearly 8.0 with Magnesium Carbonate, and the rhodies grew better and better as the soil pH went up because the Magnesium level was going up. pH had little to do with it. So, this is a good thing to know if you are trying to grow rhododendrons in New Mexico, for instance, where the soil is frequently alkaline to start with, although there you would want to use an acid form of Magnesium like Magnesium sulfate, Epsom salts. But your garden, your farm crops and your fruits and berries wouldn’t necessarily like it (except the blueberries). High levels of Magnesium in relation to Calcium are common in Organic gardening and farming, though, because people are told to lime their soils with dolomite lime, which is high in Magnesium. So, its not ph that matters as much as its the right nutrient combination for the plant. It just so happens that the right nutrient combination is most often found in a naturally acid soil. That observation doesn't necessarily mean that the soil must be acid and in no way infers anything about aluminum. Aluminum sulfate acidifies the soil because of the sulfate, not the aluminum. Actually, its possible aluminum, being Al+++, could take the place of 3 hydrogen ions (H+), though more likely is it would take the place of a Ca++ and 1 H+. Anyway, its the sulfate which causes acidity. Sulfate + water = sulfuric acid + oxygen Ammonium sulfate is NH4 + SO4, so not only will the sulfate make sulfuric acid, but the NH4 (ammonium) will break down to NO3 (nitrate) and release extra H+ into the soil, which makes the soil more acid. Same thing with epsom salts. Mg + SO4. Sulfate when added to water takes the hydrogen out of water and makes an acid. Blueberries have sensitive roots which lack root hairs found on most plants. The application of nitrates is known to "burn" blueberry roots and could kill the plant. See this link and take note on page 9 of the pic of blueberry roots with NH4 and NO3. Blueberries, and their relatives’ cranberries, lingonberries, and bilberries have somewhat unique N requirements. They are not able to use nitrate forms of N (NO3-N) effectively. These plants have evolved in soil conditions that do not naturally contain a significant amount of NO3-N and they depend more on ammonium-N (NH4-N). Blueberries take up both forms of N, but they have limited nitrate reductase activity. Nitrate reductase is an enzyme that is needed to convert nitrate to amino acids and proteins. The limited nitrate reductase system in blueberries means that they cannot efficiently utilize nitrate forms of N. Some reports also state that excessive nitrate fertilization can lead to leaf burn. (I'm pretty sure "leaf burn" is a misprint, it should probably be "root burn") Also consider from the 1st page: Blueberries have fine, fibrous roots that do not develop root-hairs. Going back to the first link I posted: The most easily recognized symptom of Al toxicity is the inhibition of root growth, and this has become a widely accepted measure of Al stress in plants. In simple nutrient solutions micromolar concentrations of Al can begin to inhibit root growth within 60 min. So if Al harms roots and blueberries have fine, sensitive roots, why would you want to add aluminum to your soil for the benefit of blueberries? Furthermore, aluminum sulfate dissolves in water and becomes immediately available to plants (as a toxin). As opposed to the aluminum in soil that is locked-up. Lastly, from page 45 of spectrum analytics site which contained the root pics: Aluminum deserves special attention with blueberries because of the very acid soil pH which the crop requires. At these acid pH’s, there is often a considerable amount of soluble Al in the soil solution. This can cause several negative results. First, soluble Al has a strong affinity for soluble P. This is the same form of P that the blueberries require. The result of the excess soluble Al can be the requirement for a higher soil P test than suggested by us and other authorities. Another potential problem is the cation competition caused by excessive soluble Al. The likely result of this cation competition is reduced uptake of one or more of the cation micronutrients (Cu, Mn, Fe, and Zn). About the only way to identify either of these potential problems is with leaf analysis. Do not use aluminum sulfate to lower the soil pH because aluminum is toxic to blueberries and is already present in many soils in the region in quantities that can negatively impact blueberry plants once the pH is lowered. Do NOT use aluminum sulfate, as this material is toxic to blueberries. Hopefully one day nurseries will stop recommending aluminum sulfate for anything. Sulfuric Acid and Water Safety

If you spill some sulfuric acid on your skin, you want to wash it off with copious amounts of running, cold water as soon as possible. Water is less dense than sulfuric acid, so if you pour water on the acid, the reaction occurs on top of the liquid. If you add the acid to the water, it sinks. Any wild and crazy reactions have to get through the water or beaker to get to you. How do you remember this? Here are some mnemonics:

• AA: Add Acid
• Acid to Water, like A&W Root Beer
• Drop acid, not water
• If you think your life's too placid, add the water to the acid
• First the water, then the acid, otherwise it won't be placid

Personally, I don't find any of those mnemonics easy to remember. I get it right because I figure if I get it wrong, I'd rather have a whole container of water splash on me than a whole container of sulfuric acid, so I take my chances with the small volume of acid and the large volume of water.

Alkali Soils: Factors, Effects and Reclamation

Saline soils represent a group of soils in which percentages of soluble salts, usually chlorides and sulphates of the alkali bases are very high.

The pH of saline soils are always high. In India saline soils occur in many provinces, as U.P., West Bengal, Punjab, Bihar, Orissa, Maharashtra Tamil Nadu, M.P., Andhra Pradesh, Gujarat, Delhi and Rajasthan covering an area of about 7 million hectares.

There are 4 major tracts in India where salinity problem is acute.

(a) The arid tract of Rajasthan and Gujarat,

(b) Semi-arid alluvial tracts of Punjab, Haryana and Uttar Pradesh,

(c) The arid and semi-arid tracts of Southern States, and

In U.P, the saline (usar) soils are distributed in Kanpur, Lucknow, Hardoi, Unnao, Allahabad, Rai Bareh, Azamgarh and many other districts covering about 1.29 million hectares. In Punjab alone saline soil covers about 2 million hectare area. In U.P. and Punjab the saline soils are gradually increasing in area. These lands are known by a variety of names in local agricultural parlance. By far the most common of them is mar derived from the Sanskrit word Ushtra meaning sterile or barren. Other terms like reh, char, lone, thur or shora are also popular. The word alkali is of Arabic origin meaning ash-like and is used to designate hard and intractable soils generally known by the names rakkar, kallar, bara and bari.

The salty soils are of three types:

(i) Saline or solonchak or white alkali soils:

In these, salinity is caused by soluble salts other than alkali salts. They have high soluble salts and low exchangeable sodium.

(ii) Alkali or sodic or solonetz or black alkali soils:

These are formed by accumulation of alkalies, such as Na, K etc. in excess. Such soils have low salt content but high exchangeable sodium.

(iii) Saline-alkali soils:

In these, alkali and other soluble salts have combined effects. They are also called saline sodic as they have high salt content and high exchangeable sodium. The United States Salinity Laboratory recently designated these soils scientifically on the basis of soil analysis following the ideas of Sigmond and Gedroiz (1954) (Table 23.4).

Three classes of salty soil:

Soils in which salinity is mainly due to accumulation of alkali salts are called alkali soils or usar soils. High alkalinity in the soil adversely affects the plant growth, thereby reduces the crop yield. Such sterile or unproductive soils are called barren soils. The main salts present in the alkali soils are Na₂SO₄, K₂SO₄, NaCl, and KCl.

Types of alkali soils:

These soils are of two types:

(i) Black alkali soils—In these soils, Na₂CO₃ is found in excess.

(ii) White alkali soils—In this group, NaCl is present in excess.

Russians call such soils as solonchack. Bertholet suggested that Na₂CO₃ was formed in black alkali soil by interaction of NaCl and CaCO₃.

Factors Which Make the Soils Alkaline:

1. Poor drainage in arid region,

2. Rapid evaporation of alkaline soil solution, and

3. Excess uptake of alkaline salts and little percolation.

In arid and semi-arid regions, the rainfall is too low to leach or remove the saline matter from the top soils. Besides this, water along with dissolved alkali salts moves upward by capillary action which on reaching to the soil surface evaporates and the salts accumulate in the form of a hard layer or pan in the subsoil. This hard layer is responsible for impermeability of such soils. Miller is of the opinion that many plants absorb excess acidic ions, e.g., NO – ₃, than the basic ions. This excessive removal of acidic ions results in the accumulation of basic ions which make the soil alkaline.

According to a chemical hypothesis, alkali soils may result in the following steps:

(a) Reaction between NaCl or KCl and soil (S):

NaCl + S (Soil) → Na (S) + CI – ion

(b) Then the soluble products are leached away from the soil surface by drainage water, and

(c) Finally, reaction between insoluble Na (S) complex and carbonates.

2 Na (S) + CaCO₃ → Ca (S) (alkaline) + Na₂CO₃ (alkaline)

Effects of alkali salts on vegetation:

The alkali salts show the following effects on plants:

(i) Due to excessive accumulation of salts, concentration of soil solution becomes high. This decreases absorption of nutrients by plants and causes plasmolysis of cell cytoplasm in the plants which may be fatal sometimes. All these effects are responsible for stunted growth of plants.

(ii) If sodium is absorbed by the plants in excess, it shows toxic effects. Chloride salt of alkaline elements causes the death of trees. BaCO₃ and BaCl are toxic to all plants.

(iii) Presence of excess salts in the soil retards the germination of seeds and growth of seedlings. Plants die before bearing fruits.

(iv) Alkali salts in the soil also affect the plant growth by reducing the size of leaves in alkali soil, plant roots remain superficial, bark of stem turns brown or black, green tissues are less developed.

Alkali tolerance in plant:

Some plants are resistant to alkali salts. Barley, wheat oats sorghum, sugar-beet, berseem are best suited to grow in alkaline soil. Cotton and grapes are also alkah tolerants. Uppal et al (1961) prepared a list of crops that can be grown at various stages of reclamation. They can be categorized as high, medium and low salt tolerant crops. High salt tolerant crops are Dhaincha, paddy, sugarcane in kharif and oat, berseem, lucerne, sanji (Trigonella spp), and barley in rabi. Medium salt tolerant crops to be tried during second stage of reclamation are castor, cotton, jowar, bajra, and maize for kharif and mustard and wheat for rabi.

Low salt tolerant crops are sesamum, moong, urd, arhar, and sannhemp in kharif and gram peas, linseed during rabi. Uppal et al also listed babul, dhak, jhand, khair, chokra, neem, Lasora, Sisham, siris, bahera and reonjha as trees that can be planted on saline alkali soils.

The alkali tolerance of plants depends upon:

(i) Physiological constitution of cell cytoplasm of the plant.

(ii) Length of roots. Shallow roots are more affected by alkalinity than the deeper roots.

(iii) Alkali salts reduce and correct the soil acidity and improve the physical conditions of the soils.

(iv) Calcium salts provide calcium to the plants.

(v) Many alkali salts change toxic elements, such as aluminum and Mn. into their harmless compounds.

Reclamation of Alkali, Saline and Saline-Alkali Soils:

The excessive accumulation of alkali salts in the soils is injurious for plants growth It is necessary, therefore, to reduce the percentage of salts to optimum or normal level so that plants may grow luxuriantly in such soils.

There are several methods of reclamation which can be grouped as follows:

(A) Chemical method in which some chemicals are added to the soil in order to brine the alkalinity to desired level.

(B) Mechanical practices such as improving drainage and leaching, mechanical shattering of clay pans, and scrapping.

(C) Cultural method (growing salt tolerant plants).

Since fundamental causes in various groups of salty soils are different, their reclaiming techniques are different. Hence, these are discussed separately.

1. Reclamation of Alkali Soils:

Alkali soils are best reclaimed by the following methods:

(A) Chemical method:

(1) By cationic exchange (replacement of alkali from soil colloids by calcium ions). Application of calcium sulphate (gypsum) in the soil reduces alkalinity to a great extent and makes the soil fertile.

The reaction proceeds in the following way:

Good drainage leaches away Na₂SO₄.

(2) Alkali salt percentage can also be reduced in the soil by the use of acid forming chemical amendments such as sulphur, ferrous sulphate and limestone. Sulphur, when applied to the soil, oxidises and forms sulphuric acid which converts carbonates of sodium and potassium to Na₂SO₄ and K₂SO₄ respectively that may be removed from top soil by drainage water. The amount of gypsum and sulphur required to reclaim the alkali soils will be different depending upon the degree of alkalinity, drainage and buffering capacity of soils.

The types of reaction which occur when an amendment is applied to an alkali soil are given below:

In the next step, if soil is calcareous—

But if the soil is non-calcareous—

(v) 2Na-Clay + H2S04= 2H-clay + Na₂S0₄

(2) With lime-sulphur:

Now if the soil is calcareous—

(iii) 2Na-Clay + CaSO4 = Ca-Clay + Na₂SO₄

But if the soil is non-calcareous—

(3) With ferrous sulphate:

Now if the soil is calcareous—

But if the soil is non-calcareous—

(4) With limestone on non-calcareous soils:

(i) Na-Clay + H₂O = H-Clay + NaOH

(5) With any H-Clay:

The use of pyrite (FeS₂) as an amendment is a recent development in the chemical amelioration and reclamation of alkali soils. In presence of moisture and air, pyrite is converted into sulphuric acid which then replaces exchangeable sodium by hydrogen or calcium released from insoluble calcium present in the soil. In addition it is said to correct iron deficiency and lime induced iron chlorosis in alkali soils. It is important to mention that the formation of H2SO4 in the soil by the application of pyrite may take place through chemical and microbiological actions. Pyrite is oxidised according to the following equation suggested by Bloomfield (1973).

FeS₂ + 2Fe +3 = 3Fe +2 + 2s (Chemical)

Sulphur thus formed could be the substrate for thioxidants which convert it into H₂SO₄.

Temple and Kochler (1954) explained the action of ferroxidans on the formation of H₂SO₄ as follows:

FeSO₄ formed in the above reaction may be converted into H2SO4 by hydrolysis.

In brief, the pyrite is oxidized in soils to ferrous sulphate and sulphuric acid as depicted in the following equation:

Both sulphuric acid and ferrous sulphate help in reclamation of calcareous as well as non- calcareous salt affected soils by lowering the pH and solubilising free calcium from calcium carbonate present.

The reactions are given below:

In salt affected calcareous soils:

H₂SO₄ formed in reaction II reacts as per equations la and lb

III. H₂SO₄ also neutralizes NaHCO₃ and Na₂CO₃ present in these soils.

But if the soil is non-calcareous:

H₂SO₄ formed in reaction II acts in similar manner as in reaction I.

(3) Dhar’s method. In India, Dr. Neel Ratan Dhar (1935) succeeded in reducing the alkalinity and salinity of the soil by the use of molasses and press-mud.

For one acre land he recommended the mixture of the following substances:

(i) 2 tons of molasses, (ii) 1-2 tons of press-mud (a waste product of sugar industry) and (iii) 50-100 pounds of P₂O₅ in the form of basic slag.

The molasses is fermented by soil microbes and as a result of fermentation organic acids are produced which lower the alkalinity and increase the availability of phosphates. The press- mud contains Ca which forms calcium salts that reduce the content of exchangeable sodium. Phosphate helps in the microbial fixation of nitrogen into nitrogenous compounds in the soil.

(B) Mechanical methods:

The alkali salts are removed by:

(1) Scraper or by rapidly moving streams of water,

(2) Deep ploughing of the land which reduces the alkalinity and makes the soil more permeable.

(3) Application of green manures of Dhaincha, guar, jantar (Sesbania aculeata) has been found most successful in reclamation of alkali and saline soils.

(4) Spreading of straw and dried grasses and leaves on the alkaline soil.

(C) Cultural method:

Growing of alkali tolerant crops and plants, such as sugar-beet, rice, patsann (Hibiscus cannabinus), wild indigo and babul in such soils successfully reduces alkalinity. Rice is commonly the first crop grown on salty lands to be reclaimed. In Punjab the usual practice of reclamation of salty lands involves growing of paddy after first initial leaching followed by berseem or senji which has higher water requirement than Dhaincha as green manure which IS followed by sugarcane and then wheat or cotton.

Introduction of leguminous crops helps in building up of nitrogen supply and opens the soils. Dhaincha-paddy-berseem rotation has been found to be the best cropping pattern on mild type of alkali soils in Punjab region. In U.P. also, paddy or dhaincha-paddy are the usual crops taken during first stage of reclamation of salty soils. This is followed by berseem or barley in winter. Pulse crops like gram or peas show poor performance.

II. Reclamation of saline soil:

Saline soil can be reclaimed by the following methods:

(1) By lowering the water table 5-6 feet below the surface. In sloppy area, it can be done by making network of 5-6 feet deep trenches at right angles to the slopes. In course of 2 or 3 Successive leaching, harmful salts are removed. A deep ploughing is also helpful in reclamation of saline soil. This also makes the soil loose and thus facilitates the downward movement of salty water in the soil.

(2) Salt tolerant crops, e.g., rice, sugar cane, barley and castor gradually remove salts from the soil.

(3) In case of saline soils which do not contain calcium salts, the addition of CaSO₄ (gypsum) is beneficial. Supply of calcium in the soil can indirectly be maintained by of organic matters which on decomposition produce CO₂. The CO₂ gas, so produced, combines with insoluble calcium carbonate in moist condition to form soluble calcium bicarbonate. This also reduces alkalinity.

(4) Application of green manure, organic manures, organic residues, acids or acid formers is yet another good way to reduce salinity.

III. Reclamation of saline-alkaline soil:

Here the problem of reclamation is two-fold because of:

(a) Heavy accumulation of different types of salts,

(b) Poor percolation due to the presence of hard clay pan and highly dispersed sodium clay.

Such soil can be reclaimed by:

(i) Mechanical shattering of clay pans. This helps in downward movement of water.

(ii) Application of gypsum in the soil. This is followed by flushing with plenty of water.

(iii) Green maturing with Dhaincha (Sesbenia aculeata).

(iv) Growing of salt tolerant plants, e.g., paddy in kharif and oat and barley in rabi seasons are recommended for such soils. ,

Schoonover (1959) worked on the soils of India and enlisted the following technical requirements for reclamation of saline and alkaline soils:

(1) Necessity of good drainage.

(2) Availability of sufficient water to wash the excess salts from the top soils.

(3) Good soil management including land leveling, good bonding for irrigation and recent and advanced agronomic practices.

Watch the video: Sulfuric Acid vs Plants (January 2023).