Nitrogen cycleFrom Wikipedia, the free encyclopediaJump to: navigation, search

Schematic representation of the flow of nitrogen through the environment. The importance of bacteria in the cycle is immediately recognized as being a key element in the cycle, providing different forms of nitrogen compounds assimilable by higher organisms.

The nitrogen cycle is the process by which nitrogen is converted between into ammonia acids. They make proteins used by producers. This transformation can be carried out via both biological and non-biological processes. Important processes in the nitrogen cycle include fixation, mineralization, nitrification, and denitrification. The majority of Earth's atmosphere (approximately 79%) is nitrogen,[1] making it the largest pool of nitrogen. However, atmospheric nitrogen is unavailable for biological use, leading to a scarcity of usable nitrogen in many types of ecosystems. The nitrogen cycle is of particular interest to ecologists because nitrogen availability can affect the rate of key ecosystem processes, including primary production and decomposition. Human activities such as fossil fuel combustion, use of artificial nitrogen fertilizers, and release of nitrogen in wastewater have dramatically altered the global nitrogen cycle.

Contents

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1 Human influences on the nitrogen cycle o 1.1 Wastewater treatment o 1.2 Environmental damage

2 References

[edit] Human influences on the nitrogen cycle

Main article: Human impacts on the nitrogen cycle

As a result of extensive cultivation of legumes (bean plants), growing use of the Haber-Bosch process in the creation of chemical fertilizers, and pollution emitted by vehicles and industrial plants, human beings have more than doubled the annual transfer of nitrogen into biologically available forms.[2] In addition, humans have significantly contributed to the transfer of nitrogen trace gases from Earth to the atmosphere, and from the land to aquatic systems. Human alterations to the global nitrogen cycle are most intense in developed countries and in Asia, where vehicle emissions and industrial agriculture are highest.[3]

N2O (nitrous oxide) has risen in the atmosphere as a result of agricultural fertilization, biomass burning, cattle and feedlots, and other industrial sources.[4] N2O has deleterious effects in the stratosphere, where it breaks down and acts as a catalyst in the destruction of atmospheric ozone. N2O in the atmosphere is a greenhouse gas, currently the third largest contributor to global warming, after carbon dioxide and methane. While not as abundant in the atmosphere as carbon dioxide, for an equivalent mass, nitrous oxide is nearly 300 times more potent in its ability to warm the planet.[5]

Ammonia (NH3) in the atmosphere has tripled as the result of human activities. It is a reactant in the atmosphere, where it acts as an aerosol, decreasing air quality and clinging on to water droplets, eventually resulting in acid rain. Fossil fuel combustion has contributed to a 6 or 7 fold increase in NOx flux to the atmosphere. NO2 actively alters atmospheric chemistry, and is a precursor of tropospheric (lower atmosphere) ozone production, which contributes to smog, acid rain, damages plants and increases nitrogen inputs to ecosystems.[6] Ecosystem processes can increase with nitrogen fertilization, but anthropogenic input can also result in nitrogen saturation, which weakens productivity and can kill plants.[2] Decreases in biodiversity can also result if higher nitrogen availability increases nitrogen-demanding grasses, causing a degradation of nitrogen-poor, species diverse heathlands.[7]

[edit] Wastewater treatment

Onsite sewage facilities such as septic tanks and holding tanks release large amounts of nitrogen into the environment by discharging through a drainfield into the ground. Microbial activity consumes the nitrogen and other contaminants in the wastewater.

However, in certain areas, the soil is unsuitable to handle some or all of the wastewater, and, as a result, the wastewater with the contaminants enters the aquifers. These contaminants accumulate and eventually end up in drinking water. One of the contaminants concerned about the most is nitrogen in the form of nitrates. A nitrate concentration of 10 ppm (parts per million) or 10 milligrams per liter is the current EPA limit for drinking water and typical household wastewater can produce a range of 20–85 ppm.

The health risk associated with drinking water (with >10 ppm nitrate) is the development of methemoglobinemia and has been found to cause blue baby syndrome. Several states have now

started programs to introduce advanced wastewater treatment systems to the typical onsite sewage facilities. The result of these systems is an overall reduction of nitrogen, as well as other contaminants in the wastewater.

[edit] Environmental damage

Additional risks posed by increases in fixed nitrogen in aquatic systems include spurring the creation and growth of eutrophic lakes and oceanic dead zones through algal bloom-induced hypoxia.[8][9]

The extent and effects of the anthropogenically-induced doubling of biologically available nitrogen in the soils, waters, and air of the earth during the past century are still poorly understood.[

The Nitrogen Cycle

Of Microbes and Men

by John Arthur Harrison, Ph.D.

Key Concepts hide

The nitrogen cycle is the set of biogeochemical processes by which nitrogen undergoes chemical reactions, changes form, and moves through difference reservoirs on earth, including living organisms.

Nitrogen is required for all organisms to live and grow because it is the essential component of DNA, RNA, and protein. However, most organisms cannot use atmospheric nitrogen, the largest reservoir.

The five processes in the nitrogen cycle -- fixation, uptake, mineralization, nitrification, and denitrification -- are all driven by microorganisms.

Humans influence the global nitrogen cycle primarily through the use of nitrogen-based fertilizers.

Nitrogen (N) is an essential component of DNA, RNA, and proteins, the building blocks of life. All organisms require nitrogen to live and grow. Although the majority of the air we breathe is N2, most of the nitrogen in the atmosphere is unavailable for use by organisms. This is because the strong triple bond between the N atoms in N2 molecules makes it relatively inert. In fact, in order for plants and animals to be able to use nitrogen, N2 gas must first be converted to more a chemically available form such as ammonium (NH4

+), nitrate (NO3-), or organic nitrogen (e.g.

urea - (NH2)2CO). The inert nature of N2 means that biologically available nitrogen is often in short supply in natural ecosystems, limiting plant growth and biomass accumulation.

Nitrogen is an incredibly versatile element, existing in both inorganic and organic forms as well as many different oxidation states. The movement of nitrogen between the atmosphere, biosphere, and geosphere in different forms is described by the nitrogen cycle (Figure 1), one of the major biogeochemical cycles. Similar to the carbon cycle, the nitrogen cycle consists of various storage pools of nitrogen and processes by which the pools exchange nitrogen (arrows) (see our The Carbon Cycle module for more information).

Figure 1: The nitrogen cycle. Yellow arrows indicate human sources of nitrogen to the environment. Red arrows indicate microbial transformations of nitrogen. Blue arrows indicate physical forces acting on nitrogen. And green arrows indicate natural, non-microbial processes affecting the form and fate of nitrogen.

Five main processes cycle nitrogen through the biosphere, atmosphere, and geosphere: nitrogen fixation, nitrogen uptake (organismal growth), nitrogen mineralization (decay), nitrification, and denitrification. Microorganisms, particularly bacteria, play major roles in all of the principal nitrogen transformations. As microbially mediated processes, these nitrogen transformations tend to occur faster than geological processes like plate motion, a very slow, purely physical process

that is a part of the carbon cycle. Instead, rates are affected by environmental factors that influence microbial activity, such as temperature, moisture, and resource availability.

©The Microbial WorldFigure 2: Part of a clover root system bearing naturally occurring nodules of Rhizobium, bacteria that can fix atmospheric nitrogen. Each nodule is about 2-3 mm long. Image courtesy of http://helios.bto.ed.ac.uk/bto/microbes/nitrogen.htm.

Nitrogen fixation

N2 NH4+ Nitrogen fixation is the process wherein N2 is converted to ammonium, essential

because it is the only way that organisms can attain nitrogen directly from the atmosphere. Certain bacteria, for example those among the genus Rhizobium, are the only organisms that fix nitrogen through metabolic processes. Nitrogen fixing bacteria often form symbiotic relationships with host plants. This symbiosis is well-known to occur in the legume family of plants (e.g. beans, peas, and clover). In this relationship, nitrogen fixing bacteria inhabit legume root nodules (Figure 2) and receive carbohydrates and a favorable environment from their host plant in exchange for some of the nitrogen they fix. There are also nitrogen fixing bacteria that exist without plant hosts, known as free-living nitrogen fixers. In aquatic environments, blue-green algae (really a bacteria called cyanobacteria) is an important free-living nitrogen fixer.

In addition to nitrogen fixing bacteria, high-energy natural events such as lightning, forest fires, and even hot lava flows can cause the fixation of smaller, but significant amounts of nitrogen (Figure 3). The high energy of these natural phenomena can break the triple bonds of N2 molecules, thereby making individual N atoms available for chemical transformation.

Within the last century, humans have become as important a source of fixed nitrogen as all natural sources combined. Burning fossil fuels, using synthetic nitrogen fertilizers, and cultivation of legumes all fix nitrogen. Through these activities, humans have more than doubled the amount of fixed nitrogen that is pumped into the biosphere every year (Figure 3), the consequences of which are discussed below.

Figure 3: Recent increases in anthropogenic N fixation in relation to “natural” N fixation. Modified from Vitousek, P. M. and P. A. Matson (1993). Agriculture, the global nitrogen cycle, and trace gas flux. The Biogeochemistry of Global Change: Radiative Trace Gases. R. S. Oremland. New York, Chapman and Hall: 193-208.

Nitrogen uptake

NH4+ Organic N The ammonia produced by nitrogen fixing bacteria is usually quickly

incorporated into protein and other organic nitrogen compounds, either by a host plant, the bacteria itself, or another soil organism. When organisms nearer the top of the food chain (like us!) eat, we are using nitrogen that has been fixed initially by nitrogen fixing bacteria.

Nitrogen mineralization

Organic N NH4+ After nitrogen is incorporated into organic matter, it is often converted back

into inorganic nitrogen by a process called nitrogen mineralization, otherwise known as decay. When organisms die, decomposers (such as bacteria and fungi) consume the organic matter and lead to the process of decomposition. During this process, a significant amount of the nitrogen contained within the dead organism is converted to ammonium. Once in the form of ammonium, nitrogen is available for use by plants or for further transformation into nitrate (NO3

-) through the process called nitrification.

Nitrification

NH4+ NO3

- Some of the ammonium produced by decomposition is converted to nitrate via a process called nitrification. The bacteria that carry out this reaction gain energy from it. Nitrification requires the presence of oxygen, so nitrification can happen only in oxygen-rich environments like circulating or flowing waters and the very surface layers of soils and sediments. The process of nitrification has some important consequences. Ammonium ions are positively charged and therefore stick (are sorbed) to negatively charged clay particles and soil organic matter. The positive charge prevents ammonium nitrogen from being washed out of the soil (or leached) by rainfall. In contrast, the negatively charged nitrate ion is not held by soil particles and so can be washed down the soil profile, leading to decreased soil fertility and nitrate enrichment of downstream surface and groundwaters.

Denitrification

NO3- N2+ N2O Through denitrification, oxidized forms of nitrogen such as nitrate and nitrite

(NO2-) are converted to dinitrogen (N2) and, to a lesser extent, nitrous oxide gas. Denitrification

is an anaerobic process that is carried out by denitrifying bacteria, which convert nitrate to dinitrogen in the following sequence:NO3

- NO2- NO N2O N2.

Nitric oxide and nitrous oxide are both environmentally important gases. Nitric oxide (NO) contributes to smog, and nitrous oxide (N2O) is an important greenhouse gas, thereby contributing to global climate change.

Once converted to dinitrogen, nitrogen is unlikely to be reconverted to a biologically available form because it is a gas and is rapidly lost to the atmosphere. Denitrification is the only nitrogen transformation that removes nitrogen from ecosystems (essentially irreversibly), and it roughly balances the amount of nitrogen fixed by the nitrogen fixers described above.

Human alteration of the N cycle and its environmental consequences

Early in the 20th century, a German scientist named Fritz Haber figured out how to short circuit the nitrogen cycle by fixing nitrogen chemically at high temperatures and pressures, creating fertilizers that could be added directly to soil. This technology has spread rapidly over the past century, and, along with the advent of new crop varieties, the use of synthetic nitrogen fertilizers has led to an enormous boom in agricultural productivity. This agricultural productivity has helped us to feed a rapidly growing world population, but the increase in nitrogen fixation has had some negative consequences as well. While the consequences are perhaps not as obvious as an increase in global temperatures or a hole in the ozone layer, they are just as serious and potentially harmful for humans and other organisms.

Not all of the nitrogen fertilizer applied to agricultural fields stays to nourish crops. Some is washed off of agricultural fields by rain or irrigation water, where it leaches into surface or ground water and can accumulate. In groundwater that is used as a drinking water source, excess nitrogen can lead to cancer in humans and respiratory distress in infants. The U.S. Environmental Protection Agency has established a standard for nitrogen in drinking water of 10 mg per liter nitrate-N. Unfortunately, many systems (particularly in agricultural areas) already exceed this level. By comparison, nitrate levels in waters that have not been altered by human activity are

rarely greater than 1 mg/L. In surface waters, added nitrogen can lead to nutrient over-enrichment, particularly in coastal waters receiving the inflow from polluted rivers. This nutrient over-enrichment, also called eutrophication, has been blamed for increased frequencies of coastal fish-kill events, increased frequencies of harmful algal blooms, and species shifts within coastal ecosystems.

Reactive nitrogen (like NO3- and NH4

+) present in surface waters and soils, can also enter the atmosphere as the smog-component nitric oxide (NO) and the greenhouse gas nitrous oxide (N2O). Eventually, this atmospheric nitrogen can be blown into nitrogen-sensitive terrestrial environments, causing long-term changes. For example, nitrogen oxides comprise a significant portion of the acidity in acid rain which has been blamed for forest death and decline in parts of Europe and the Northeast United States. Increases in atmospheric nitrogen deposition have also been blamed for more subtle shifts in dominant species and ecosystem function in some forest and grassland ecosystems. For example, on nitrogen-poor serpentine soils of northern Californian grasslands, plant assemblages have historically been limited to native species that can survive without a lot of nitrogen. There is now some evidence that elevated levels of atmospheric N input from nearby industrial and agricultural development have paved the way for invasion by non-native plants. As noted earlier, NO is also a major factor in the formation of smog, which is known to cause respiratory illnesses like asthma in both children and adults.

Currently, much research is devoted to understanding the effects of nitrogen enrichment in the air, groundwater, and surface water. Scientists are also exploring alternative agricultural practices that will sustain high productivity while decreasing the negative impacts caused by fertilizer use. These studies not only help us quantify how humans have altered the natural world, but increase our understanding of the processes involved in the nitrogen cycle as a whole

Making and testing ammoniaIn this experiment pupils make ammonia, investigate its solubility in water and its alkaline nature. The experiment provides a useful precursor to the ammonia fountain experiment.

Read our standard health & safety guidance

Lesson organisation

Because ammonia has a strong smell and is poisonous in quantity, pupils should, ideally, carry out this experiment in fume cupboards. Alternatively, provided that the quantities of reactants are no greater than as stated here, the experiment could be carried out in a well-ventilated laboratory.

Depending on arrangements for setting up and clearing away, the experiments should take around 30 minutes.

Apparatus and Chemicals

Eye protection

Each working group will require:

Retort stand, boss and clampBoiling tube with stopper and delivery tube – see diagram. This apparatus must be dry (see note 1) Boiling tubes, 2 – must be thoroughly dry (see note 1)Beaker (100 cm3) Beaker (250 cm3) or largerBunsen burnerHeat resistant matSpatula

Students will need access to:

Ammonium chloride (Harmful) (see note 3)Calcium hydroxide (Irritant) (see note 3)Calcium oxide (Irritant) - (optional) (see note 3)Concentrated hydrochloric acid (Corrosive) (see note 2)

Red Litmus paperBlue Litmus paperUniversal indicator paper

Technical notes

Ammonia gas (Toxic, Dangerous for the environment) Refer to CLEAPSS Hazcard 5Ammonium chloride (Harmful) Refer to CLEAPSS Hazcard 9ACalcium hydroxide (Irritant) Refer to CLEAPSS Hazcard 18Calcium oxide (Irritant) Refer to CLEAPSS Hazcard 18Concentrated hydrochloric acid (Corrosive) Refer to CLEAPSS Hazcard 47A

1 Ammonia (Toxic, Dangerous for the environment) is a very soluble gas. If the preparation and collection apparatus is not dry, the solubility experiment will not work.

2 Two drops of the acid in the bottom of a stoppered, labelled test-tube are sufficient. A small number of prepared tubes could be arranged around the laboratory so that students can access them only under supervision.

3 Small quantities of the solid chemicals should be provided in stoppered bottles. Use lumps or granules of calcium oxide (optional).

Procedure

HEALTH & SAFETY: Eye protection must be worn at all times. Ammonia gas is Toxic and Dangerous for the environment and pungent-smelling and must not be inhaled. The experiments must only be carried out in a fume cupboard or in a well-ventilated laboratory.

a In a small beaker, mix 2 spatulas of the ammonium chloride with 2 spatulas of the calcium hydroxide together. The two solids begin to react immediately on mixing.

b Hold a piece of each colour of Litmus paper over the mixture and observe the colour change. Test also with a piece of Universal indicator paper.

c Transfer the mixture of ammonium chloride and calcium hydroxide into a boiling tube and set up the apparatus as shown in the diagram.

d Optional step - put the lump of calcium oxide into the boiling tube containing the ammonium chloride/calcium hydroxide mixture. The calcium oxide will absorb the water produced in the reaction and ensure that the ammonia gas is dry.

e Gently warm the reaction mixture.

f Collect a test-tube which contains a few drops of concentrated hydrochloric acid. Remove the stopper from this test-tube and hold the open end near the end of the ammonia gas delivery tube. Observe what happens. Replace the stopper on the test-tube of hydrochloric acid and return the test-tube to its original place.

g Two-thirds fill a large beaker with water. This is needed for step j.

h Continue to gently warm the reaction mixture. Hold one of the dry boiling tubes in position as shown in the second diagram. Notice that the ammonia is collected with the boiling tube upside down. This is because ammonia is less dense than air.

i Test around the open end of the collecting boiling tube with Universal Indicator paper to check that the collecting tube is full of ammonia.

j Hold the tube of ammonia upside down then quickly put it, mouth still downwards, into water in a beaker. The ammonia dissolves in the water and the level of the water should rise up inside the test tube. If you want to try this a second time, use a fresh dry boiling tube.

Teaching notes

Ammonia is a very soluble gas. 1 cm3 of water dissolves about 800 cm3 of ammonia at room temperature. A few drops of water will easily dissolve a test-tube of ammonia. This is why it is essential that the apparatus is dry.

The reaction to produce ammonia also produces water. The purpose of the calcium oxide is to help to prevent this water from coming out of the delivery tube as water vapour.

The equation for the generation of ammonia is:

2NH4Cl(s) + Ca(OH)2(s) → CaCl2(s) + 2NH3(g) + 2H2O(g)

Calcium oxide reacts with water to produce calcium hydroxide:

CaO(s) + H2O(l) → Ca(OH)2(s)

Ammonia is one of the very few common alkaline gases. When it dissolves in water it reacts reversibly according to the equation:

NH3(g) + H2O(l) → NH4+(aq) + OH–(aq)

There is generally enough water on the surface of indicator papers to dissolve ammonia without having to moisten the paper.

The white fumes ('smoke') produced with the hydrogen chloride given off by the concentrated hydrochloric acid consist of fine particles of solid ammonium chloride.

The equation isNH3(g) + HCl(g) → NH4Cl(s)

HABER BOSCH

Dasar teori pembuatan amonia dari nitrogen dan hydrogen ditemukan oleh Fritz Haber (1908), seorang ahli kimia dari Jerman. Sedangkan proses industri pembuatan amonia untuk produksi secara besar-besaran ditemukan oleh Carl Bosch, seorang insinyur kimia juga dari Jerman. Persamaan termokimia reaksi sintesis amonia adalah :

N2(g) + 3H2(g) ⇄ 2NH3(g) ∆H = -92,4Kj Pada 25oC : Kp = 6,2×105

Berdasarkan prinsip kesetimbangan kondisi yang menguntungkan untuk ketuntasan reaksi ke kanan (pembentukanNH3) adalah suhu rendah dan tekanan tinggi. Akan tetapi, reaksi tersebut berlangsung sangat lambat pada suhu rendah, bahkan pada suhu 500oC sekalipun. Dipihak lain, karena reaksi ke kanan eksoterm, penambahan suhu akan mengurangi rendemen. Proses Haber-Bosch semula dilangsungkan pada suhu sekitar 500oC dan tekanan sekitar 150-350 atm dengan katalisator, yaitu serbuk besi dicampur dengan Al2O3, MgO, CaO, dan K2O.

Reaksi kekanan pada pembuatan amonia adalah reaksi eksoterm. Reaksi eksoterm lebih baik jika suhu diturunkan, tetapi jika suhu diturunkan maka reaksi berjalan sangat lambat . Amonia punya berat molekul 17,03. Amonia ditekanan atmosfer fasanya gas. Titik didih Amonia -33,35 oC, titik bekunya -77,7 oC, temperatur & tekanan kritiknya 133 oC & 1657 psi. Entalpi pembentukan (∆H), kkal/mol NH3(g) pada 0oC, -9,368; 25 oC, -11,04. Pada proses sintesis pd suhu 700-1000oF, akan dilepaskan panas sebesar 13 kkal/mol. Kondisi optimum untuk dapat bereaksi dengan suhu 400- 600oC, dengan tekanan 150-300 atm. Kondisi optimum pembuatan amonia (NH3) dapat digambarkan pada tabel berikut :

Tabel : Kondisi Optimum Pembuatan NH3

No Faktor Reaksi : N2(g) + 3H2(g) ⇄ 2NH3(g) ∆H= -924 kJ

Kondisi Optimum

1. Suhu 1. Reaksi bersifat eksoterm

2. Suhu rendah akan menggeser kesetimbangan kekanan.

3. Kendala:Reaksi berjalan lambat

400-600oC

2. Tekanan 1. Jumlah mol pereaksi lebih besar dibanding dengan jumlah mol produk.

2. Memperbesar tekanan akan menggeser kesetimbangan kekanan.

3. Kendala Tekanan sistem dibatasi oleh kemampuan alat dan faktor keselamatan.

150-300 atm

3. Konsentrasi Pengambilan NH3 secara terus menerus akan menggeser kesetimbangan kearah kanan

_

4. Katalis Katalis tidak menggeser kesetimbangan kekanan, tetapi mempercepat laju reaksi secara keseluruhan

Fe dengan campuran Al2O3 KOH dan garam lainnya

Pengaruh katalis pada sistem kesetimbangan adalah dapat mempercepat terjadinya reaksi kekanan atau kekiri, keadaan kesetimbangan akan tercapai lebih cepat tetapi katalis tidak mengubah jumlah kesetimbangan dari spesies-spesies yang bereaksi atau dengan kata lain katalis tidak mengubah nilai numeris dalam tetapan kesetimbangan. Peranan katalis adalah mengubah mekanisme reaksi kimia agar cepat tercapai suatu produk.

Katalis yang dipergunakan untuk mempercepat reaksi memberikan mekanisme suatu reaksi yang lebih rendah dibandingkan reaksi yang tanpa katalis. Dengan energi aktivasi lebih rendah menyebabkan maka lebih banyak partikel yang memiliki energi kinetik yang cukup untuk mengatasi halangan energi aktivasi sehingga jumlah tumbukan efektif akan bertambah sehingga laju meningkat. Perbandingan reaksi dengan katalis dan tanpa katalis dapat dilihat pada gambar dihalaman berikut:

Gambar 1 : Perbandingan mekanisme reaksi menggunakan katalis dan tanpa katalis

Dengan kemajuan teknologi sekarang digunakan tekanan yang jauh lebih besar, bahkan mencapai 700 atm. Untuk mengurangi reaksi balik, maka amonia yang terbentuk segera dipisahkan. Mula-mula campuran gas nitrogen dan hidrogen dikompresi (dimampatkan) hingga mencapai tekanan yang diinginkan. Kemudian campuran gas dipanaskan dalam suatu ruangan yang bersama katalisator sehingga terbentuk amonia. Diagram alur dari proses Haber-bosch untuk sintesis amonia

diberikan pada Gambar 1 berikut ini :

A notional flow-sheet for the cyanamide process.

The first step in the process is to make lime from limestone:CaCO3 + heat → CaO + CO2

this is then heated with Coal in an an anoxic environment to make Calcium Carbide:CaO + 3C + heat → CaC2 + COThe actual nitrogen fixing comes from reacting Calcium Carbide with pure Nitrogen, thus for this process to be industrially practical it required the Linde process of fractionation of liquid air. The reaction takes place at 2atm or ~0.2MPa, heated by through the Ohmic heating of a Carbon rod:CaC2 + N2 → CaCN2 + CFinally in the quest to make Ammonia, the Calcium Cyanamide is mixed with water and NaOH (as a catalyst) for hydrolysis:CaCN2 + H2O → 2NH3 + CaCO3

The Calcium Carbonate can easily be separated as it is a solid, and the Ammonia can be distilled, allowing the NaOH to be recycled back for more hydrolysis.

Contrast this with the Haber-Bosch process for making Ammonia, which at the time required the same costly liquid air seperator as well as an electrolytic seperator for producing hydrogen and higher pressure catalytic reactor:

A notional flow-sheet for an early Haber-Bosh process.

By simply glaring at it one sees that, as a way of making Ammonia, the Haber-Bosh process is by far simpler. Since it doesn’t require multiple furnaces and the intermediary steps of producing Cyanamide its operating costs should be lower (assuming one has an efficient electrolysis system for hydrogen). Of course the Ammonia reactor requires an expensive catalyst and recycle system since a single pass is not particularly efficient.

The nitrate ion, NO3-

The nitrate ion is distinctly awkward and there isn't any simple way of working out its shape. The problem is that it contains a co-ordinate (dative covalent) bond.

Note:  If you are a bit hazy about co-ordinate (dative covalent) bonding you could follow this link before you go on. It probably isn't necessary - all you need to know is that a co-ordinate bond is one in which both electrons come from the same atom.

The nitrate ion comes from nitric acid, so we'll start from the structure of that.

Look carefully at the bonding around the nitrogen. You will see that one of the bonds is formed entirely from the lone pair on the nitrogen. That's the co-ordinate bond.

The nitrate ion is formed by the loss of the hydrogen ion, and so its structure is:

Around the central nitrogen there are 4 pairs of shared electrons, and no remaining lone pair. The original lone pair has now become a bonding pair. Two of those pairs make up a double bond. The double bond unit and the two single bonds arrange themselves as far apart as possible in a trigonal planar arrangement - exactly the same as the carbonate ion.

Note:  Co-ordinate bonds can be shown using an arrow going from the atom donating the lone pair towards the atom receiving it. Once again, this structure is a simplification (OK for this level). In fact, all the nitrogen-oxygen bonds are identical, and the charge is delocalised evenly over all three oxygen atoms.

NITROGEN - brief summary of a few points

The structure of the element: o Non-metal existing as diatomic molecule N2, with a triple covalent bond, N N.

Physical properties:  o Colourless gas; mpt -210oC; bpt -196oC; poor conductor of heat/electricity.

Group, electron configuration (and oxidation states):  o Gp5; e.c. 2.5  or 1s22s22p3; Variety of oxidation states from -3 to +5 e.g. o NH3 (-3), N2O (+1), NO (+2), NCl3 (+3), NO2 (+4) and N2O5 and HNO3 (+5).

Reaction of nitrogen with oxygen:  o At high temperatures e.g. in car engines, nitrogen(II) oxide (nitrogen monoxide) is

formed. N2(g) + O2(g) ==> 2NO(g) 

o and the nitrogen(II) oxide rapidly reacts in air to form nitrogen(IV) oxide (nitrogen

dioxide). 2NO(g) + O2(g) ==> 2NO2(g) 

o The theoretical highest oxide is N2O5 nitrogen(V) oxide (nitrogen pentoxide) and does

exist. Reaction of nitrogen oxides with water:

o Nitrogen(IV) oxide dissolves to form an acidic solution of weak nitrous acid and strong

nitric acid. 2NO2(g) + H2O(l) ==> HNO2(aq) + HNO3(aq) 

or 2NO2(g) + 2H2O(l) ==> HNO2(aq) + H3O+(aq) + NO3

-(aq) 

o NO and N2O are neutral oxides but nitrogen(V) oxide is strongly acidic and dissolves to

form nitric acid. N2O5(s) + H2O(l) ==> 2HNO3(aq) 

or N2O5(s) + 3H2O(l) ==> 2H3O+(aq) + 2NO3

-(aq) 

Reaction of nitrogen oxides with acids:o None, only acidic (N2O3 (very unstable), NO2 and N2O5) or neutral (N2O and NO), in nature.

Reaction of nitrogen oxides with bases/alkalis:o Nitrogen(IV) oxide or nitrogen dioxide forms sodium nitrite and sodium nitrate with

sodium hydroxide solution. o 2NO2(g) + 2NaOH(aq) ==> NaNO2(aq) + NaNO3(aq) + H2O(l) o ionic equation: 2NO2(g) + 2OH-

(aq) ==> NO2-(aq) + NO3

-(aq) + H2O(l)

o As well as being a neutralisation reaction, it is also a redox reaction, the oxidation states  of oxygen (-2) and hydrogen (+1) do not change BUT the oxidation state of nitrogen changes from two at (+4) to one at (+3) and one at (+5). The simultaneous change of an element into an lower and upper oxidation sate is sometimes called disproportionation.

Reaction of nitrogen with chlorine:  o None, but the unstable yellow oily liquid chloride can be made indirectly.

Reaction of chloride with water: o Slowly hydrolyses to form weak nitrous acid and strong hydrochloric acid.

NCl3(l) + 2H2O(l) ==> HNO2(aq) + 3HCl(aq)  or NCl3(l) + 2H2O(l) ==> HNO2(aq) + 3H+

(aq) + 3Cl-(aq) 

or NCl3(l) + 5H2O(l) ==> HNO2(aq) + 3H3O+(aq) + 3Cl-

(aq)  Reaction of nitrogen with water:

o Slightly soluble but NO reaction. Other comments:

o An essential element for plants, hence need for nitrogen compounds in compost and

artificial fertilisers (NPK bags!).

PHOSPHORUS - brief summary of a few points

The structure phosporus: o Two solid allotropes (red and white) consisting of P4 molecules, also a polymer form.

Physical properties of phosphorus:  o Colourless gas; mpt 44oC; bpt 280oC; poor conductor of heat/electricity.

Group, electron configuration (and oxidation states):  o Gp5; e.c. 2,8,5  or 1s22s22p63s23p3; Variety of oxidation states from -3 to +5 e.g. o PH3 (-3), P4O6 (+3), P4O10, PCl5 and H3PO4 (+5).

Reaction of phosphorus with oxygen:  o With limited air/oxygen, on heating the phosphorus, the covalent white solid

phosphorus(III) oxide is formed. P4(s) + 3O2(g) ==> P4O6(s) 

o With excess air/oxygen, on heating the phosphorus, the covalent white solid phosphorus(V) oxide is formed.

P4(s) + 5O2(g) ==> P4O10(s)  Reaction of phosphorus oxides with water: Both oxides dissolve in water to form acidic

solutions. o Phosphoric(III) oxide forms phosphoric(III) acid.

P4O6(s) + 6H2O(l) ==> 4H3PO3(aq)  o Phosphoric(III) oxide forms phosphoric(V) acid.

P4O10(s) + 6H2O(l) ==> 4H3PO4(aq)  Reaction of phosphorus with chlorine: 

o With limited chlorine, on heating the phosphorus, the covalent liquid phosphorus(III) chloride is formed.

P4(s) + 3Cl2(g) ==> 4PCl3(l)  o With excess chlorine, on heating the phosphorus, the ionic* solid phosphorus(V) chloride

is formed. P4(s) + 5Cl2(g) ==> 4PCl5(s)  * PCl5 is a bit unusual for an 'expected covalent' liquid chloride.

It is an ionic solid with the structure [PCl4]+[PCl6]-  Hence its melting point is much greater than the liquid

phosphorus(III) chloride, where the molecules are only held together by the inter-molecular forces.

However, gaseous phosphorus(V) chloride consists of PCl5 covalent molecules.

Reaction of phosphorus oxides with acids:o None, only acidic in nature.

Reaction of phosphorus oxides with bases/alkalis:o Both oxides dissolve in alkalis to form a whole series of phosphate(III) and phosphate(V)

salts. o with excess strong bases like sodium hydroxide, the simplified equations are:

P4O6(s) + 12NaOH(aq) ==> 4Na3PO3(aq) + 6H2O(l) sodium phosphate(III) formed from phosphorus(III) oxide

ionic equation: P4O6(s) + 12OH-(aq) ==> 4PO3

3-(aq) + 6H2O(l)

P4O10(s) + 12NaOH(aq) ==> 4Na3PO4(aq) + 6H2O(l) sodium phosphate(V) formed from phosphorus(V) oxide

ionic equation: P4O10(s) + 12OH-(aq) ==> 4PO4

3-(aq) + 6H2O(l)

If the empirical formulae P2O3 and P2O5 are used, just halve all the balancing numbers.

Other than using excess sodium hydroxide, other salts can be formed. e.g. P4O10(s) + 4NaOH(aq) + 2H2O(l) ==> 4NaH2PO4(aq) sodium dihydrogen

phosphate(V) or  P4O10(s) + 8NaOH(aq) ==> 4Na2HPO4(aq) + 2H2O(l) disodium hydrogen

phosphate(V) o -

Reaction of phosphorus chlorides with water: o Phosphorus(III) chloride hydrolyses rapidly and exothermically to form phosphoric(III)

acid. PCl3(l) + 3H2O(l) ==> H3PO3(aq) + 3HCl(aq)

o Phosphorus(V) chloride initially hydrolyses to form phosphorus oxychloride? and hydrochloric acid.

PCl5(s) + H2O(l) ==> POCl3(aq) + 2HCl(aq)  Then on boiling the aqueous solution, phosphoric(V) acid is formed and more

hydrochloric acid. POCl3(aq) + 3H2O(l) ==> H3PO4(aq) + 3HCl(aq)  overall: PCl5(s) + 4H2O(l) ==> H3PO4(aq) + 5HCl(aq) 

Reaction of element with water: o None.

 

The shapes of some molecules and ions of nitrogen and phosphorus

 

electrons: three bond pairs and one lone pair, PYRAMIDAL or TRIGONAL PYRAMID shape: e.g. ammonia NH3 with bond angle of approximately 109o. Note: the exact H-N-H angle is 107o due to the extra repulsion of one lone pair (for H-X-H angles: NH3 > H2O and < CH4).

 

electrons: three bond pairs and one lone pair, PYRAMIDAL or TRIGONAL PYRAMID shape. e.g. nitrogen trifluoride/trichloride, NCl3, or phosphorus(III) fluoride/chloride (phosphorus trifluoride/trichloride), PF3/PCl3, with bond angles Q-X-Q of approximately 109o and similarly with ions like the oxonium ion H3O+ (Q = F, Cl etc. X = N, P etc.)

 

electrons: 5 bond pairs, TRIGONAL BIPYRAMID shape: e.g. phosphorus(V) fluoride (phosphorus pentafluoride) PF5, gaseous phosphorus(V) chloride, PCl5, with bond angles 90o and 180o based on the vertical Q-X-Q bond and 120o based on the central trigonal planar arrangement. Note that solid PCl5 has an ionic structure and is not a trigonal bipyramid molecule - a tetrahedral [PCl4]+ ion and an octahedral [PCl6]- ion.

 

H3N:=>BF3 Boron trifluoride (3 bonding pairs, 6 outer electrons) acts as a lone pair acceptor (Lewis acid) and ammonia (3 bond pairs) and lone pair which enables it to act as a Lewis base - a an electron pair donor. It donates the lone pair to the 4th 'vacant' boron orbital to form a sort of 'adduct' compound. Its shape is essentially the same as ethane, a sort of double tetrahedral with H-N-H, N-B-F and F-B-F bond angles of ~109o.

 

Nitrogen(IV) oxide, NO2 (nitrogen dioxide) is bent shaped (angular), O-N-O bond angle ~120o because of two bonding groups of bonding electrons and a single lone electron in the same plane as the bonding pairs of electrons.

The nitrate(III) ion, NO2- (nitrite ion)

is bent shaped (angular), O-N-O bond angle ~120o due to two groups of bonding electrons and one lone pair of electrons.

The nitrate(V) ion, NO3- (nitrate ion)

is trigonal planar, O-N-O bond angle 120o due to three bonding groups of electrons and no lone pairs of electrons.

The nitronium ion, NO2+, is linear,

O-N-O bond angle of 180o because there are two groups of bonding electrons and no lone pairs of electrons (you easily see this from the NO2 neutral molecule diagram below).

The shapes are deduced below using dot and cross diagrams and VSEPR theory and illustrated in the valence bond dot and cross diagrams below.

 

Ammonia can act as an electron pair donor ligand in transition metal ion complexes

e.g. tetraamminedichlorochromium(III) complex ion - cis/trans isomers (Z/E isomers)

With excess aqueous ammonia a pale blue hexa-ammine complex is formed with hexaaquanickel(II) ions

[Ni(H2O)6]2+(aq) + 6NH3(aq) [Ni(NH3)6]2+

(aq) + 6H2O(l)

and also excess aqueous ammonia a pale blue hexa-ammine complex is formed with aqueous copper(II) ions

[Cu(H2O)6]2+(aq) + 4NH3(aq) [Cu(NH3)4(H2O)2]2+

(aq) + 4H2O(l)

or [Cu(H2O)4]2+(aq) + 4NH3(aq) [Cu(NH3)4]2+

(aq) + 4H2O(l)

 

The Synthesis of ammonia - The Haber Process

Ammonia gas is synthesised in the chemical industry by reacting nitrogen gas with hydrogen gas in what is known as the Haber-Bosch Process, named after two highly inventive and subsequently famous chemists.

The nitrogen is obtained from liquified air (80% N2). Air is cooled and compressed under high pressure to form liquid air (liquefaction). The liquid air is fractionally distilled at low temperature to separate oxygen (used in welding, hospitals etc.), nitrogen (for making ammonia), Noble Gases e.g. argon for light bulbs, helium for balloons).

The hydrogen is made by reacting methane (natural gas) and water or from cracking hydrocarbons (both reactions are done at high temperature with a catalyst).

o CH4 + H2O ==> 3H2 + COo e.g. C8H18 ==> C8H16 + H2

The synthesis equation for this reversible reaction is ...

o N2(g) + 3H2(g) 2NH3(g)

.. which means an equilibrium will form, so there is no chance of 100% yield even if you use, as you actually do, the theoretical reactant ratio of nitrogen : hydrogen of 1 : 3

In forming ammonia 92kJ of heat energy is given out (i.e. exothermic, 46kJ of heat released per mole of ammonia formed).

Four moles of 'reactant' gas form two moles of 'product' gas, so there is net decrease in gas molecules on forming ammonia.

So applying the Le Chatelier equilibrium rules, the formation of ammonia is favoured by  ...o (a) Using high pressure because you are going from 4 to 2 gas molecules (the high

pressure also speeds up the reaction because it effectively increases the concentration of the gas molecules), but higher pressure means more dangerous and more costly engineering.

o (b) Carrying out the reaction at a low temperature, because it is an exothermic reaction favoured by low temperature, but this may produce too slow a rate of reaction,

o So, the idea is to use a set of optimum conditions to get the most efficient yield of ammonia and this involves getting a low % yield (e.g. 8% conversion) but fast. Described below are the conditions to give the most economic production of ammonia.

o these arguments make the point that the yield* of an equilibrium reaction depends on the conditions used.

* The word 'yield' means how much product you get compared to the theoretical maximum possible if the reaction goes 100%.

For more on chemical economics see Extra Industrial Chemistry page. In industry pressures of 200 - 300 times normal atmospheric pressure are used in line with the

theory. Theoretically a low temperature would give a high yield of ammonia BUT ...

o Nitrogen is very stable molecule and not very reactive i.e. chemically inert, so the rate of reaction is too slow at low temperatures.

o To speed up the reaction an iron catalyst is used as well as a higher temperature (e.g. 400-450oC).

o The higher temperature is an economic compromise, i.e. it is more economic to get a low yield fast, than a high yield slowly!

o Note: a catalyst does NOT affect the yield of a reaction, i.e. the equilibrium position BUT you do get there faster!

At the end of the process, when the gases emerge from the iron catalyst reaction chamber, the gas mixture is cooled under high pressure, when only the ammonia liquefies and is so can be removed and stored in cylinders.

Any unreacted nitrogen or hydrogen (NOT liquified), is recycled back through the reactor chamber, nothing is wasted! [nitrogen (-196oC) and hydrogen (-252oC) have much lower boiling points than ammonia (-33oC). Boiling points increase with pressure, but these normal atmospheric pressure values offer a fair comparison].

To sum up: A low % yield of ammonia is produced quickly at moderate temperatures and pressure, and is more economic than getting a higher % equilibrium yield of ammonia at a more costly high pressure and a slower lower temperature reaction.

AND there are some more general notes on Chemical Economics on the Industrial Chemistry page.