chemical of the week
Almost everyone has smelled the sharp, penetrating odor of ammonia, NH3. As the active product of “smelling salts,†the compound can quickly revive the faint of heart and light of head. But more than a sniff of this toxic, reactive, and corrosive gas can make one very ill indeed. It can, in fact, be fatal. Ammonia is pretty nasty stuff. Nevertheless, it is also an extremely important bulk chemical widely used in fertilizers, plastics, and explosives.
             The melting and boiling points of ammonia, –77.7°C and –33.5°C, respectively, are both considerably higher than the corresponding properties of its chemical “cousins,†PH3 and AsH3. This failure of NH3 to follow the usual trend of decreasing melting and boiling points with decreasing molecular weights indicates abnormally strong intermolecular attractions. The forces involved stem from hydrogen bonding, a consequence of the high electronegativity of nitrogen and the small size of the hydrogen atom.
             The NH3 molecule has a large dipole moment, and this is consistent with its geometry, a trigonal pyramid.
             The electronic arrangement in nitrogen obeys the octet rule. The four pairs of electrons (three bonding pairs and one non-bonding lone pair) repel each other, giving the molecule its non-planar geometry. The H–N–H bond angle of 107 degrees is close to the tetrahedral angle of 109.5 degrees. Because of this, the electronic arrangement of the valence electrons in nitrogen is described as sp3 hybridization of atomic orbitals.
             The polarity of NH3 molecules and their ability to form hydrogen bonds explains to some extent the high solubility of ammonia in water. However, a chemical reaction also occurs when ammonia dissolves in water. In aqueous solution, ammonia acts as a base, acquiring hydrogen ions from H2O to yield ammonium and hydroxide ions.
|
NH3(aq) + H2O(l)
 NH4+(aq) + OH-(aq) |
             The production of hydroxide ions when ammonia dissolves in water gives aqueous solutions of ammonia their characteristic alkaline (basic) properties. The double arrow in the equation indicates that an equilibrium is established between dissolved ammonia gas and ammonium ions. Not all of the dissolved ammonia reacts with water to form ammonium ions. A substantial fraction remains in the molecular form in solution. In other words, ammonia is a weak base. A quantitative indication of this strength is given by its base ionization constant:
             In contrast, the ammonium ion acts as a weak acid in aqueous solution because it dissociates to form hydrogen ion and ammonia.
|
NH4+(aq)
NH3(aq) + H+(aq) |
             The ammonium ion is found in many common compounds, such as ammonium chloride, NH4Cl. Typically, ammonium salts have properties similar to the corresponding compounds of the Group IA alkali metals.
             The commercial production of ammonia by the direct combination of nitrogen and hydrogen is an example of equilibrium in the gaseous state. The equation for the reaction and its equilibrium constant expression are
|
N2(g) + 3 H2(g)
2 NH3(g) |
 |
             At 300°C, Kc has a value of 9.6, indicating that at this temperature, an appreciable amount of NH3 forms from N2 and H2. Because the reaction gives off heat (Δ H°= –92.0 kJ for the equation above), increasing the temperature drives the reaction to the left. Thus, Kc decreases with increasing temperature. The equilibrium mixture at 500°C contains less NH3 than at 300°C or at 100°C. If one is in the business of making ammonia (and money), the object is to make as much NH3 as possible as quickly as possible. The temperature dependence of the equilibrium constant suggests that working at low temperatures is better because more ammonia is obtained at equilibrium. Alas, equilibrium isn’t everything! All chemical reactions slow down as the temperature decreases. While a low temperature favors a high equilibrium yield of ammonia, it also dictates that a long time will be required to obtain the yield. The ideal method is a balance between yield and speed.
             A great asset in the production of ammonia is a catalyst which speeds the reaction between nitrogen and hydrogen. Early in this century, a German academic chemist, Fritz Haber, and an industrial colleague, Carl Bosch, found that a mixture of Fe2O3 and Fe3O4 catalyzes this reaction at temperatures in the range of 400°C to 600°C. The yield of ammonia was further enhanced by working at gas pressures between 200 and 400 atmospheres. In the balanced equation for the reaction, the number of moles of product (2 NH3) is less than the total number of moles of reactants (N2 + 3 H2). Therefore, high pressures drive the reaction forward, decreasing the number of moles of gas in the mixture. Although now modified and improved, the Haber-Bosch process continues to be the most common method for making ammonia. The nitrogen is obtained from liquefied air, and the hydrogen is usually from natural gas decomposed by heating.
             The Haber-Bosch process is also an example of the complex impact of chemistry upon life. At the start of World War I, Germany was dependent upon the natural nitrate deposits of Chile for the nitrogen compounds required to manufacture explosives. The Allied blockade of South American ports soon cut off this supply. Had it not been for the alternative source of nitrogen compounds provided by the direct synthesis of ammonia, Germany most likely would have been forced to surrender several years before 1918. By prolonging the war, the Haber-Bosch process indirectly cost thousands of lives. However, over the years, the fertilizer produced by the same process has increased crop yields around the globe and spared millions from starvation.
WHAT EVERY AQUARIST AND POND KEEPER SHOULD KNOW ABOUT AMMONIA TEST KITS…
Interpreting ammonia readings has many users confused, including the experts. It is worthwhile to review the following information to help better understand what ammonia test kits and their readings are all about. First, you should know that what is generally referred to as “ammonia” is in two forms: un-ionized ammonia (NH 3 ) and ionized ammonia (NH 4 +). Ionized ammonia is relatively nontoxic while un-ionized ammonia is toxic to fishes and aquatic invertebrates, even in low concentrations.
Ammonia test kits commonly available in the aquarium pet industry read total ammonia: which include both unionized and ionized ammonia (NH 3 +NH 4 +). The unionized (toxic) form of ammonia is a part of the total reading, and two determinations have to be made to find the amount that is in the toxic form. Consult the paragraph entitled “The effects of pH and temperature on ammonia” at the end of this article for additional information on toxic ammonia levels.
There are two kinds of test kit readings; one for ionic ammonia and the other for ammonia nitrogen (a reading of the nitrogen present in NH 4 +). Some pond and aquarium test kits give readings as ammonia nitrogen, some as total ionic ammonia, and many don’t explain how their readings are calculated. Older Kordon test kits read in units of total ammonia as nitrogen (N) for the salicylate kits and as ammonia ion for the nessler kit (as indicated in the instructions supplied with each kit). The latest generation of Kordon’s ammonia tests kits read both the ion and as nitrogen.
What is the difference between test kit readings of ammonia ion and ammonia/nitrogen?
Many aquarists and pond keepers are perplexed by the use of the terms ammonia ion and ammonia/nitrogen (or ammonia as nitrogen), and do not know what the difference is between the two. The differences between the two terms are in how the chemical composition of the same ammonia molecules are being measured, which can be in two different ways, giving two different readings, each of which is correct. If you have a reading specified as ammonia/nitrogen it can be converted to ammonia ion by multiplying the reading by 1.3. If the reading is expressed as ammonia ion, it can be converted to ammonia/nitrogen by dividing by 1.3. For example, if the ammonia ion reading is 2.6 ppm, you can divide by 1.3 and you will get an ammonia/nitrogen reading of 2.0 ppm (2.6 ÷ 1.3 = 2.0). The conversion factor of 1.3 is based on the atomic weight proportions of nitrogen and hydrogen in ammonia (1.3 weight units of ammonia contain 1.0 weight units of nitrogen).
If the ammonia is measured as the total molecular weight of the molecules of the ammonia compounds in the water, which will be NH 3 + NH 4 +, the total amount of combined nitrogen and hydrogen atoms in the molecules are being measured as ammonia ion. If only the nitrogen atoms contained in the complete ammonia molecules are being measured and the measurement for the hydrogen atom is left out, the reading is for ammonia/nitrogen (N).
Most information in the aquarium and pond keeping hobby refers to ammonia as the total ionic ammonia so when you see the reference to ammonia as 0.8, for example, this usually means in aquarium and pond publications the total molecular weight for all the ammonia molecules. The number 0.8 means that the total weight is in ppm (parts per million) or mg/L (milligrams per liter), which for our purposes are essentially the same measurement. When the information provided does not indicate whether it is as ammonia ion or ammonia as nitrogen, there indeed can be confusion as to which is being measured. However, when you remember that the two readings are different by 1.3x, it may be possible to know to which of the two types are involved; at least for general aquarist and pond keeping purposes the readings are close enough together that it is not of major concern.
Why are there two different ways to refer to measuring the amount of ammonia in the water?
In scientific measurements of water quality in lakes, streams and oceans, the amount of nitrogen in the water, what ever the form, is an important measurement. Therefore, ammonia as nitrogen is widely used in the scientific literature. In biological laboratory research on living aquatic animals, the amount of total ionic ammonia present is important to know, particularly in considering the ammonia ions’ effect on the physiological processes of the animals. Therefore, in this type of research, and in much of aquarium and ornamental pond keeping, it has been more important to measure ammonia as ammonia ion. But keep in mind that both readings are correct, and either reading can be directly acquired by multiplying or dividing by 1.3, respectively.
WHEN THE AMMONIA READINGS ARE MADE, WHAT DO THEY MEAN?
This question is often asked by aquarists and pond keepers wanting to know what the significance of their readings really are. Put simply, it is important to keep the level of ammonia as low as possible in the water, because it is toxic to fishes and aquatic invertebrates (toxic means poisonous; stressful but not necessarily lethal) . Even levels below 0.5 ppm (0.5 mgL) are toxic or enfeebling to many aquatic animals.
To immediately get rid of the ammonia in the water use Kordon’s AmQuel ® or AmQuel® Plus(which quickly and permanently combine with the ammonia molecules to form new, non-toxic molecules, see KPD-51), or make partial or whole water changes. Over time, a properly set up biological filter will control ammonia build up. For specific aquatic animals check the technical aquaristic literature for their tolerance to ammonia. In general: Many of the hardiest fishes cannot survive toxic ammonia levels above 3.0 ppm (= mgL); sick fishes succumb at lower toxic ammonia levels than healthier fishes; most marine animals have a much lower tolerance to ammonia than freshwater animals, although many crustaceans (crabs, shrimp, lobsters) can stand 9.0 ppm and higher levels of toxic ammonia. To keep on the safe side use Kordon’s AmQuel or do water changes as soon as you see 0.2 ppm or more total ammonia in the water.  Aquarium fishes should be kept at 0.25 ppm or lower, and some are adversely affected at as low as 0.5 ppm
WHERE DOES THE AMMONIA COME FROM?
Uneaten food, dead and decaying animals and plants, as well as the fecal, urinary and respiratory waste products of the aquatic animals are broken down into ammonia by heterotrophic bacteria in a process called amonification. A secondary ammonia source is tap water. Municipal water supplies often have high levels of ammonia added to the water combined with chlorine to produce chloramine. Chloramines are used to control bacteria and viruses which are highly toxic to aquatic animals. It is important for fish keepers to know whether or not chloramines are being used in their tap water, contact the local water supplier for information. Chloramines are extremely stable and will not “gass off” like chlorine. For information on dealing with chloramines see the AmQuel Product Data Sheet KPD-51. Whatever the source of ammonia in the water, it exists in a balance between ionized and un-ionized, converting back and forth depending upon the pH and temperature of the water. Ammonia (that is not combined with chlorine) is normally converted to nitrite and ultimately nitrate by naturally occuring nitrifying bacteria. This action prevents the buildup of potentially lethal ammonia concentrations in the water. See KPD-64 on BIOLOGICAL FILTRATION . If anything happens to the water that is detrimental to these bacteria (low alkalinity, low oxygen level, addition of antibiotics, etc.) the bacteria will become dormant or die, and the ammonia level in the water will immediately increase. Therefore, the regular use of Kordon AmQuel Plus, such as dosing once a week to eliminate all toxic nitrogen compounds is recommended, as is monitoring of aquariums and ponds by use of water quality test kits is very important.
WHAT IS THE EFFECT OF pH (ACIDITY/ALKALINITY) AND TEMPERATURE ON AMMONIA?
If you want to be more exact in assessing whether the amount of ammonia in the water is in the toxic form or not, it is necessary to measure the pH as well as the ammonia, because the pH of the water determines how much of the ammonia is in the un-ionic (toxic) form. Dependent upon the pH of the water, the ratio between ionized (nontoxic) and un-ionized (toxic) ammonia changes rapidly. The ammonia molecules are in balance between the nontoxic ionized and the toxic un-ionized; the higher the pH the more ammonia is in the toxic form; conversely, the lower the pH the less ammonia is in the toxic form. Temperature also plays a part in determining how much of the total ammonia molecule is in the toxic form; the higher the temperature, the more toxic the ammonia.
Kordon has created a chart of ammonia tables that show how to determine the amount of toxic ammonia in a given sample of water when the pH and temperature of the water and the amount of total ammonia are known. Once test kit readings are made of the ammonia, pH and temperature, the exact amount of ammonia in the toxic form can be determined. This chart also points out how pH and temperature can have a dramatic effect on the amount of toxic ammonia present in an aquarium or pond. For example, if the ammonia reading is 0.5 ppm (mg/L), if the pH is 7.0 and if the temperature is 73.4° F then the amount of toxic ammonia is .0025; the amount of toxic ammonia in this sample is below a harmful level for most aquatic animals. If only the pH were raised to 8.0, the amount of toxic ammonia would also rise to .0218 ppm (mg/L) which is a harmful level for many aquatic animals. This result would indicate that Kordon’s AmQuel or AmQuel Plus should be added to the water to eliminate the toxic ammonia, or that a water change should be made. Testing should be continued to be sure that harmful levels of toxic ammonia do not recur. Click here for access to the AMMONIA TABLES
Ammonia-free bore cleaner
New from KG Industries is the ammonia-free KG-12 cleaner, specifically designed to clean the copper fouling from large-bore military weapons. KG-12C is for conventional use. The water-based KG-12 requires no neutralizer or waiting time. KG-12 and 12C literally consume the copper without touching any other metal component. Contact KG industries, 16790 S. Hwy 63, Bldg. 63, Hayward, WI 54843, (800) 348-9558, Fax: (715) 934-3570, e-mail: adam@kgcoatings.com,
Ammonia The story behind the gas.
A brief insight into the chemistry of ammonia:
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Ammonia has a triangular pyramidal geometry, and boiling points of 77.7*C and 33.5*C. In its pure form ammonia was prepared in 1774 by Joseph Priestly, and its composition was determined in 1785 by Claude-Louis Berthollet. Ammonia has a chemical formula of NH3, and is sp3 hybridised. Ammonia is highly polarised, due to the electronegativity of nitrogen, and as a result, has a large dipole moment. Ammonias polarisation allows it to dissociate in water forming hydroxide and ammonium ions:
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NH3 (aq) + H2O (l) Û NH4 (aq) + OH (aq)
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Ammonia solutions are basic, due to the hydroxide ions formed in solution.
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Ammonia is commercially produced by the Haber-Bosch process, which is also sometimes referred to as the Haber-Ammonia Process or Synthetic Ammonia Process. Fritz Haber, the German physical chemist, created the process in 1909, and it was further developed by Carl Bosch to make it economically viable. Both chemists won the Nobel prize for their work in this field; Haber in 1918 for its development, and Bosch in 1931 for creating high-pressure conditions which obtained a higher yield, economically.
A Quantum Chemical Study of the Catalysis for Cytidine Deaminase: Contribution of the Extra Water Molecule
Abstract:
Cytidine deaminase is known as an important enzyme responsible for the hydrolytic deamination of cytidine, which is applied as a key step to the conversion of the precursor of the cancer drug to an active form in the living body. Cytidine with water is efficiently converted to uridine with ammonia in the cleft of cytidine deaminase. In this work, the catalysis of cytidine deaminase for the hydrolytic deamination was examined using cytosine as a model of cytidine and the model molecules for the active site of cytidine deaminase by means of the quantum chemical method. We especially investigated the contribution of the water molecule from the solvent to the catalysis, because the X-ray diffraction analysis of a crystal structure has revealed the existence of the water molecule in the vicinity of the substrate bound to the active site inside the cleft. Our computations showed that the extra water molecule from the solvent has a possibility to support the catalysis of cytidine deaminase.
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Aniline Dyes Information
Aniline
Aniline, phenylamine or aminobenzene (C6H5NH2) is an organic chemical compound which is a primary aromatic amine consisting of a benzene ring and an amino group. The chemical structure of aniline is shown at the right.
Synthesis
Aniline can be produced from benzene in two steps. First, benzene is nitrated (reacted with nitric acid, a form of electrophilic substitution reaction) to give nitrobenzene. Second, the nitrobenzene is reduced to give aniline. A variety of reducing agents are effective for the reduction, including H2 (with a catalyst), hydrogen sulfide, iron, zinc, or tin.
Many derivatives of aniline can be prepared similarly.In commerce three brands of aniline are distinguished—aniline oil for blue, which is pure aniline; aniline oil for red, a mixture of equimolecular quantities of aniline and ortho- and para-toluidines; and aniline oil for safranine, which contains aniline and ortho-toluidine, and is obtained from the distillate (échappés) of the fuchsine fusion. Monomethyl and dimethyl aniline are colourless liquids prepared by heating aniline, aniline hydro-chloride and methyl alcohol in an autoclave at 220°C. They are of great importance in the colour industry. Monomethyl aniline boils at 193-195°C; dimethyl aniline at 192°C.
Properties
Aniline is oily and, although colourless, it can be slowly oxidized and resinified in air to form impurities which can give it a red-brown tint. Its boiling point is 184 °C and its melting point is -6 °C. It is a liquid at room temperature.
Like most volatile amines, it possesses a somewhat unpleasant odour of rotten fish, and also has a burning aromatic taste; it is a highly acrid poison. It ignites readily, burning with a large smoky flame.
Chemically, aniline is a weak base. Aromatic amines such as aniline are generally much weaker bases than aliphatic amines. Aniline reacts with strong acids to form salts containing the anilinium (or phenylammonium) ion (C6H5-NH3+), and reacts with acyl halides (such as acetyl chloride (ethanoyl chloride), CH3COCl) to form amides. The amides formed from aniline are sometimes called anilides, for example CH3-CO-NH-C6H5 is acetanilide, for which the modern name is N-phenyl ethanamide.
The sulphate forms beautiful white plates. Although aniline is but feebly basic, it precipitates zinc, aluminium and ferric salts, and on warming expels ammonia from its salts. Aniline combines directly with alkyl iodides to form secondary and tertiary amines; boiled with carbon disulphide it gives sulphocarbanilide (diphenyl thio-urea), CS(NHC6H5)2, which may be decomposed into phenyl mustard-oil, C6H5CNS, and triphenyl guanidine, C6H5N: C(NHC6H5)2. Sulphuric acid at 180° C gives sulphanilic acid, NH2.C6H4.SO3H. Anilides, compounds in which the amino group is substituted by an acid radical, are prepared by heating aniline with certain acids; antifebrin or acetanilide is thus obtained from acetic acid and aniline. The oxidation of aniline has been carefully investigated. In alkaline solution azobenzene results, while arsenic acid produces the violet-colouring matter violaniline. Chromic acid converts it into quinone, while chlorates, in the presence of certain metallic salts (especially of vanadium), give aniline black. Hydrochloric acid and potassium chlorate give chloranil. Potassium permanganate in neutral solution oxidizes it to nitrobenzene, in alkaline solution to azobenzene, ammonia and oxalic acid, in acid solution to aniline black. Hypochlorous acid gives para-amino phenol and para-amino diphenylamine.
Like phenols, aniline derivatives are highly reactive in electrophilic substitution reactions. For example, sulfonation of aniline produces sulfanilic acid, which can be converted to sulfanilamide. Sulfanilamide is one of the sulfa drugs which were widely used as antibacterials in the early 20th century.
Aniline and its ring-substituted derivatives react with nitrous acid to form diazonium salts. Through these, the -NH2 group of aniline can be conveniently converted to -OH, -CN, or a halide.
Uses
Originally the great commercial value of aniline was due to the readiness with which it yields, directly or indirectly, valuable dyestuffs. The discovery of mauve in 1858 by William Perkin was the first of a series of dyestuffs which are now to be numbered by hundreds. Reference should be made to the articles dyeing, fuchsine, safranine, indulines, for more details on this subject. In addition to dyestuffs, it is a starting-product for the manufacture of many drugs such as Acetaminophen/Paracetamol (Tylenol).
Currently the largest market for aniline is preparation of 4,4′-MDI, some 85% of aniline serving this market. Other uses include rubber processing chemicals (9%), herbicides (2%), and dyes and pigments (2%). [1]
History
Aniline was first isolated from the destructive distillation of indigo in 1826 by Otto Unverdorben (Pogg. Ann., 1826, 8, p. 397), who named it crystalline. In 1834, Friedrich Runge (Pogg. Ann., 1834, 31, p. 65; 32, p. 331) isolated from coal tar a substance which produced a beautiful blue colour on treatment with chloride of lime; this he named kyanol or cyanol. In 1841, C. J. Fritzsche showed that by treating indigo with caustic potash it yielded an oil, which he named aniline, from the specific name of one of the indigo-yielding plants, Indigofera anil, anil being derived from the Sanskrit nÄ«la, dark-blue, and nÄ«lÄ, the indigo plant. About the same time N. N. Zinin found that on reducing nitrobenzene, a base was formed which he named benzidam. August Wilhelm von Hofmann investigated these variously prepared substances, and proved them to be identical (1855), and thenceforth they took their place as one body, under the name aniline or phenylamine.
Its first industrial-scale use was in the manufacture of mauveine, a purple dye discovered in 1856 by William Henry Perkin.
p-toluidine, an aniline derivative, can be used in qualitative analysis to prepare carboxylic acid derivitives.
Toxicology
Aniline is toxic by inhalation of the vapour, absorption through the skin or swallowing. It causes headache, drowsiness, cyanosis, mental confusion and in severe cases can cause convulsions. Prolonged exposure to the vapour or slight skin exposure over a period of time affects the nervous system and the blood, causing tiredness, loss of appetite, headache and dizziness.[2]
Oil mixtures containing rapeseed oil denatured with aniline have been clearly linked by epidemiological and analytic chemical studies to the toxic oil syndrome that hit Spain in the spring and summer of 1981, in which 20,000 became acutely ill, 12,000 were hospitalized, and more than 350 died in the first year of the epidemic. The precise etiology though remains unknown.
Some authorities class aniline as a carcinogen, although the IARC lists it in Group 3 (not classifiable as to its carcinogenicity to humans) due to the limited and contradictary data available.
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Catalyst Information
A catalyst (Greek: καταλÏτης, catalytÄ“s) is a substance that accelerates the rate (speed) of a chemical reaction (see also catalysis). Chemical catalysts, the focus of this article, participate in reactions but are neither chemical reactants nor chemical products. More generally, one may sometimes call anything which accelerates a reaction without itself being consumed or transformed a catalyst (for example, a “catalyst for political change”).
Catalysts and reaction energetics
Catalysts enable reactions to occur much faster or at lower temperatures because of changes that they induce in the reactants. Catalysts provide an alternative pathway of lower activation energy, for a reaction to proceed whilst remaining chemically unchanged themselves. This can be observed on a Boltzmann distribution and energy profile diagram. This means that catalysts reduce the amount of energy needed to start a chemical reaction. Molecules that would not have had the energy to react or that have such low energies that they probably would have taken a long time to react are able to react in the presence of a catalyst. Thus, more molecules that need to gain less energy to react will go through the chemical reaction.
Catalysts cannot make energetically unfavorable reactions possible — they have no effect on the chemical equilibrium of a reaction because the rate of both the forward and the reverse reaction are equally affected.
Types of catalysts
Catalysts can be either heterogeneous or homogeneous. Heterogeneous catalysts are present in different phases from the reactants (e.g. a solid catalyst in a liquid reaction mixture), whereas homogeneous catalysts are in the same phase (e.g. a dissolved catalyst in a liquid reaction mixture). A simple model for heterogeneous catalysis involves the catalyst providing a surface on which the reactants (or substrates) temporarily become adsorbed. Bonds in the substrate become weakened sufficiently for new bonds to be created. The bonds between the products and the catalyst are weaker, so the products are released.
For example, in the Haber process to manufacture ammonia, finely divided iron acts as a heterogenous catalyst. The metal uses active sites to allow partial weak bonding to the reactant gases, which are adsorbed onto the metal surface. As a result, the bond within the molecule of a reactant is weakened and the reactant molecules are held in close proximity to each other. In this way the particularly strong triple bond in nitrogen is weakened and the hydrogen and nitrogen molecules are brought closer together than would be the case in the gas phase, so the rate of reaction increases.
Other heterogenous catalysts include vanadium V oxide in the Contact process, nickel in the manufacture of margarine, alumina and silica in the cracking of alkanes and platinum rhodium palladium in catalytic converters.
In car engines, incomplete combustion of the fuel produces carbon monoxide, which is toxic. The electric spark and high temperatures also allow the oxygen and nitrogen to react to form nitrogen monoxide, which is acidic. Catalytic converters reduce such emissions by adsorbing CO and NO onto the catalytic surface, where the gases undergo a redox reaction. Carbon dioxide and nitrogen are desorbed from the surface and emitted as relatively harmless gases:
2CO + 2NO → 2CO(2) + N(2)
Example of homogeneous catalysts are H+(aq) which acts as a catalyst in esterification and chlorine free radicals in the break down of ozone. Chlorine free radicals are formed by the action of ultraviolet radiation on chlorofluorocarbons (CFCs). They react with ozone forming oxygen molecules and regenerating chlorine free radicals:
Cl(.) + O(3) → ClO(.) + O(2)
ClO(.) + O → Cl(.) + O(2)
N.B. Full stops in brackets denote free radicals that should be superscripted. Numbers in brackets should be subscripted
Homogeneous catalysts generally react with one or more reactants to form a chemical intermediate that subsequently reacts to form the final reaction product, in the process regenerating the catalyst. The following is a typical reaction scheme, where C represents the catalyst:
A + C → AC (1)
B + AC → AB + C (2)
Although the catalyst (C) is consumed by reaction 1, it is subsequently produced by reaction 2, so for the overall reaction:
A + B + C → AB + C
the catalyst is neither consumed nor produced. Enzymes are biocatalysts. Use of “catalyst” in a broader cultural sense is in rough analogy to the sense described here. Other biocatalysts are ribozymes and deoxyribozymes.
Poisoning a Catalyst
A catalyst can be poisoned if another compound reacts with it and bonds chemically, but does not release. This effectively destroys the usefulness of the catalyst, as it cannot participate in the reaction with which it was supposed to catalyse, just like Raney nickel catalyst has reduced activity when it is in combination with mild steel. The loss in activity of catalyst can be overcome by having a lining of epoxy or other substances .
Commonly used catalysts
Estimates are that 60% of all commerically produced chemical products involve catalysts at some stage in the process of their manufacture.[1] Some of the most famous catalysts ever developed are the Ziegler-Natta catalysts used to mass produce polyethylene and polypropylene. Probably the best-known catalytic reaction is the Haber process for ammonia synthesis, where ordinary iron is used as a catalyst. Catalytic converters made from platinum and rhodium break down some of the more harmful byproducts of automobile exhaust. The most effective catalysts are usually transition elements.
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Ammonium bicarbonate
Ammonium Bicarbonate also called bicarbonate of ammonia, ammonium hydrogen carbonate, hartshorn, or powdered baking ammonia is the bicarbonate salt of ammonia.
Ammonium bicarbonate is formed as shown above and also by passing carbon dioxide through a solution of the normal compound, when it is deposited as a white powder, which has no smell and is only slightly soluble in water. The aqueous solution of this salt liberates carbon dioxide on exposure to air or on heating, and becomes alkaline in reaction. The aqueous solutions of all the carbonates when boiled undergo decomposition with liberation of ammonia and of carbon dioxide:
NH4HCO3 → NH3 + H2O + CO2
Properties
At room temperature Ammonium bicarbonate is a white, crystalline powder with a slight odour of ammonia that can dissolve in water to give a mildly alkaline solution. It is however insoluble in acetone and alcohols. Ammonium bicarbonate decomposes at 36 to 60 °C into ammonia, carbon dioxide and water vapor in an endothermic process (as it is with many ammonium salts) and so causes a drop in the temperature of the water. When reacted with acids carbon dioxide is produced, while reactions with alkalis give ammonia.
Uses
Ammonium bicarbonate was used in the food industry as a raising agent (e.g. for gingerbread, Chinese Youtiao) before the introduction of baking soda. This compound is used as a component in the production of fire-extinguishing compounds, pharmaceuticals, dyes, pigments and it is also a basic fertilizer being a source of ammonia. Ammonium bicarbonate is still widely used in the plastic and rubber industry, in the manifacture of ceramics, in chrome leather tanning and for the synthesis of catalysts.lll
Safety
Ammonium bicarbonate is irritant to the skin, eyes and respiratory system.
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