The physical and chemical properties of the group 5/15 elements, in particular, nitrogen and phosphorus are described and explained in detail. Data table, symbol equations, oxidation states, formulae of oxides & chlorides etc.

(1) Group 5/15 Position in the periodic table - introduction, data, trends and electron configurations

• Generally speaking down a p block group the element becomes more metallic in chemical character.

• Nitrogen and phosphorus are non–metals, arsenic and antimony are semi–metals, bismuth is a true metal.

• na = not applicable

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(2) NITROGEN – summary of a few important points about its chemistry

• The structure of the element:

  • Non–metal existing as diatomic molecule N₂, with a triple covalent bond, NN.

• Physical properties: 

  • Colourless gas; mpt –210^oC; bpt –196^oC; poor conductor of heat/electricity.

• Group, electron configuration (and oxidation states): 

  • Gp5; e.c. 2.5  or 1s²2s²2p³; Variety of oxidation states from –3 to +5 e.g.

  • NH₃ (–3), N₂O (+1), NO (+2), NCl₃ (+3), NO₂ (+4) and N₂O₅ and HNO₃ (+5).

• Reaction of nitrogen with oxygen: 

  • At high temperatures e.g. in car engines, nitrogen(II) oxide (nitrogen monoxide) is formed.

    • N_2(g) + O_2(g) ==> 2NO_(g) 

  • and the nitrogen(II) oxide rapidly reacts in air to form nitrogen(IV) oxide (nitrogen dioxide).

    • 2NO_(g) + O_2(g) ==> 2NO_2(g) 

  • The theoretical highest oxide is N₂O₅ nitrogen(V) oxide (nitrogen pentoxide) and does exist.

• Reaction of nitrogen oxides with water:

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

    • 2NO_2(g) + H₂O_(l) ==> HNO_2(aq) + HNO_3(aq) 

      • or 2NO_2(g) + 2H₂O_(l) ==> HNO_2(aq) + H₃O⁺_(aq) + NO₃^–_(aq) 

  • NO and N₂O are neutral oxides but nitrogen(V) oxide is strongly acidic and dissolves to form nitric acid.

    • N₂O_5(s) + H₂O_(l) ==> 2HNO_3(aq) 

      • or N₂O_5(s) + 3H₂O_(l) ==> 2H₃O⁺_(aq) + 2NO₃^–_(aq) 

• Reaction of nitrogen oxides with acids:

  • None, only acidic (N₂O_{3 (very unstable)}, NO₂ and N₂O₅) or neutral (N₂O and NO), in nature.

• Reaction of nitrogen oxides with bases/alkalis:

  • Nitrogen(IV) oxide or nitrogen dioxide forms sodium nitrite and sodium nitrate with sodium hydroxide solution.

  • 2NO_2(g) + 2NaOH_(aq) ==> NaNO_2(aq) + NaNO_3(aq) + H₂O_(l)

  • ionic equation: 2NO_2(g) + 2OH^–_(aq) ==> NO₂^–_(aq) + NO₃^–_(aq) + H₂O_(l)

  • 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: 

  • None, but the unstable yellow oily liquid chloride can be made indirectly.

• Reaction of chloride with water:

  • Slowly hydrolyses to form weak nitrous acid and strong hydrochloric acid.

    • NCl_3(l) + 2H₂O_(l) ==> HNO_2(aq) + 3HCl_(aq) 

      • or NCl_3(l) + 2H₂O_(l) ==> HNO_2(aq) + 3H⁺_(aq) + 3Cl^–_(aq) 

      • or NCl_3(l) + 5H₂O_(l) ==> HNO_2(aq) + 3H₃O⁺_(aq) + 3Cl^–_(aq) 

• Reaction of nitrogen with water:

  • Slightly soluble but NO reaction.

• Other comments:

  • An essential element for plants, hence need for nitrogen compounds in compost and artificial fertilisers (NPK bags!).

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(3) PHOSPHORUS  – summary of a few important points about its chemistry

• The structure phosphorus:

  • Two solid allotropes (red and white) consisting of P₄ molecules, also a polymer form.

• Physical properties of phosphorus: 

  • Colourless gas; mpt 44^oC; bpt 280^oC; poor conductor of heat/electricity.

• Group, electron configuration (and oxidation states): 

  • Gp5; e.c. 2,8,5  or 1s²2s²2p⁶3s²3p³; Variety of oxidation states from –3 to +5 e.g.

  • PH₃ (–3), P₄O₆ (+3), P₄O₁₀, PCl₅ and H₃PO₄ (+5).

• Reaction of phosphorus with oxygen: 

  • With limited air/oxygen, on heating the phosphorus, the covalent white solid phosphorus(III) oxide is formed.

    • P_4(s) + 3O_2(g) ==> P₄O_6(s) 

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

    • P_4(s) + 5O_2(g) ==> P₄O_10(s) 

• Reaction of the oxides with water: Both oxides are acidic, typical non–metallic element behaviour, and both phosphorus oxides dissolve in water to form acidic solutions.

  • Phosphorus(III) oxide forms (i) phosphonic acid and (II) a little phosphoric(III) acid, which is an isomer.

    • (i) major reaction: P₄O_6(s) + 6H₂O_(l) ==> 4HPO(OH)_2(aq)

      • [a dibasic/diprotic acid, often written as H₃PO₃, but behaves as O=PH(OH)₂]

      • Note that phosphorus still has an oxidation state equivalent to +3 in phosphonic acid.

    • (ii) minor reaction: P₄O_6(s) + 6H₂O_(l) ==> 4H₃PO_3(aq)   (which could be written as P(OH)₃)

  • Phosphorus(V) oxide forms phosphoric(V) acid, no complications here!

    • P₄O_10(s) + 6H₂O_(l) ==> 4H₃PO_4(aq)   [a tribasic/triprotic acid, O=P(OH)₃]

• Reaction of phosphorus with chlorine: 

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

    • P_4(s) + 3Cl_2(g) ==> 4PCl_3(l) 

  • With excess chlorine, on heating the phosphorus, the ionic^* solid phosphorus(V) chloride is formed.

    • P_4(s) + 5Cl_2(g) ==> 4PCl_5(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:

  • None, only acidic in nature.

• Reaction of phosphorus oxides with strong bases/alkalis:

  • Both oxides dissolve in alkalis to form a whole series of phosphate(III) and phosphate(V) salts.

  • So, with strong bases like sodium hydroxide, the simplified equations are:

    • initially: P₄O_6(s) + 4NaOH_(aq) + 2H₂O_(l) ==> 4NaH₂PO_3(aq)

      • forming the mono sodium salt of phosphonic acid,

      • then, with excess sodium hydroxide, you get the disodium salt of phosphonic acid

        • NaH₂PO_3(aq) + NaOH_(aq) ==> Na₂HPO_3(aq) + H₂O_(l)

        • and there is NO trisodium salt because H₃PO₃ behaves as dibasic/diprotic O=PH(OH)₂

        • So, overall with excess sodium hydroxide the reaction is ...

        • P₄O_6(s) + 8NaOH_(aq) ==> 4Na₂HPO_3(aq) + 2H₂O_(l)

    • P₄O_10(s) + 12NaOH_(aq) ==> 4Na₃PO_4(aq) + 6H₂O_(l) sodium phosphate(V) formed from phosphorus(V) oxide

      • ionic equation: P₄O_10(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 solution, other salts can be formed.

    • e.g. P₄O_10(s) + 4NaOH_(aq) + 2H₂O_(l) ==> 4NaH₂PO_4(aq)      sodium dihydrogen phosphate(V)

    • or  P₄O_10(s) + 8NaOH_(aq) ==> 4Na₂HPO_4(aq) + 2H₂O_(l)        disodium hydrogen phosphate(V)

• Reaction of phosphorus chlorides with water:

  • Phosphorus(III) chloride hydrolyses rapidly and exothermically to form phosphoric(III) acid or phosphonic acid (see reaction of oxides with water).

    • PCl_3(l) + 3H₂O_(l) ==> H₃PO_3(aq) + 3HCl_(aq)

      • or    PCl_3(l) + 3H₂O_(l) ==> HPO(OH)_2(aq) + 3HCl_(aq)

  • Phosphorus(V) chloride initially hydrolyses to form phosphoryl chloride (phosphorus oxychloride) and hydrochloric acid.

    • PCl_5(s) + H₂O_(l) ==> POCl_3(aq) + 2HCl_(aq) 

    • Then on boiling the aqueous solution, phosphoric(V) acid is formed and more hydrochloric acid.

    • POCl_3(aq) + 3H₂O_(l) ==> H₃PO_4(aq) + 3HCl_(aq) 

    • overall the hydrolysis reaction is: PCl_5(s) + 4H₂O_(l) ==> H₃PO_4(aq) + 5HCl_(aq) 

• Reaction of element with water:

  • None.

• Reaction of phosphorus acids with strong bases/alkalis e.g. sodium hydroxide.

  • The simplified formulae are used in the neutralisation equations below, so for ...

  • Phosphonic acid H₃PO₃ aka O=PH(OH)₂

    • initially (i)  H₃PO_3(aq) + NaOH_(aq) ==> NaH₂PO_3(aq) + H₂O_(l)

    • with excess alkali (ii) NaH₂PO_3(aq) + NaOH_(aq) ==> Na₂PHO_3(aq) + H₂O_(l)

    • overall with excess alkali: H₃PO_3(aq) + 2NaOH_(aq) ==> Na₂PHO_3(aq) + 2H₂O_(l)

    • ... and for ...

  • Phosphoric(V) acid, H₃PO₄ aka O=P(OH)₃

    • Here the neutralisation can occur in three stages as each labile (acidic) proton is replaced by the sodium ion ...

    • (i) H₃PO_4(aq) + NaOH(aq) ==> NaH₂PO_4(aq) + H₂O_(l)

    • (ii) NaH₂PO_4(aq) + NaOH(aq) ==> Na₂HPO_4(aq) + H₂O_(l)

    • (iii) Na₂HPO_4(aq) + NaOH(aq) ==> Na₃PO_4(aq) + H₂O_(l)

    • overall with excess alkali: H₃PO_4(aq) + 3NaOH(aq) ==> Na₃PO_4(aq) + 3H₂O_(l)

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(4) The shapes and bond angles 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 109^o. Note: the exact H–N–H angle is 107^o 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 109^o 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 90^o and 180^o based on the vertical Q–X–Q bond and 120^o 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 ~109^o.

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) + 6NH_3(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) + 4NH_3(aq) [Cu(NH₃)₄(H₂O)₂]²⁺_(aq) + 4H₂O_(l)

or [Cu(H₂O)₄]²⁺_(aq) + 4NH_3(aq) [Cu(NH₃)₄]²⁺_(aq) + 4H₂O_(l)

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• 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).
  • CH₄ + H₂O ==> 3H₂ + CO
  • e.g. C₈H₁₈ ==> C₈H₁₆ + H₂
• The synthesis equation for this reversible reaction is ...

  • N_2(g) + 3H_2(g) 2NH_3(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's equilibrium rules, the formation of ammonia is favoured by  ...
  • (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.
  • (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,
  • 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.
  • 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 ...
  • Nitrogen is very stable molecule and not very reactive i.e. chemically inert, so the rate of reaction is too slow at low temperatures.
  • To speed up the reaction an iron catalyst is used as well as a higher temperature (e.g. 400–450^oC).
  • The higher temperature is an economic compromise, i.e. it is more economic to get a low yield fast, than a high yield slowly!
  • 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 (–196^oC) and hydrogen (–252^oC) have much lower boiling points than ammonia (–33^oC). 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.

• The Haber Synthesis of ammonia - nitrogen fixation

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(6) The uses of ammonia and derived compounds e.g. manufacture nitric acid and fertilisers

• Ammonia is oxidised with oxygen from air using a hot platinum catalyst to form nitrogen monoxide and water.
• 4NH_3(g) + 5O_2(g) ==> 4NO_(g) + 6H₂O_(g)
• The gas is cooled and reacted with more oxygen to form nitrogen dioxide.
• 2NO_(g) + O_2(g) ==> 2NO_2(g)
• This is reacted with more oxygen and water to form nitric acid.
• 4NO_2(g)+ O_2(g) + 2H₂O_(l) ==> 4HNO_3(aq)
• Nitric acid is used to make nitro–aromatic compounds from which dyes are made.
• It is also used in the manufacture of artificial nitrogenous fertilisers (like ammonium nitrate, see below).

Ammonia is used to manufacture 'artificial' nitrogenous fertilisers

• Ammonia is a pungent smelling alkaline gas that is very soluble in water.

• The gas or solution turns litmus or universal indicator blue because it is a soluble weak base or weak alkali and is neutralised by acids to form salts.

• Ammonium salts are used as 'artificial' or 'synthesised' fertilisers i.e. nitrogenous fertilisers 'man–made' in a chemical works, and used as an alternative to natural manure or compost etc.

• The fertiliser salts are made by neutralising ammonia solution with the appropriate acid.

• The resulting solution is heated, evaporating the water to crystallise the salt e.g.

  • ammonia + sulphuric acid ==> ammonium sulphate

  • 2NH_3(aq) + H₂SO_4(aq) ==> (NH₄)₂SO_4(aq)

  • AND

  • ammonia + nitric acid ==> ammonium nitrate

  • NH_3(aq) + HNO_3(aq) ==> NH₄NO_3(aq)

• The salt Ammonium chloride is used in zinc–carbon dry cell batteries. The slightly acid paste, made from the salt, slowly reacts with the zinc to provide the electrical energy from the chemical reaction.

  • ammonia + hydrochloric acid ==> ammonium chloride

  • NH_3(aq) + HCl_(aq) ==> NH₄Cl_(aq) 

    • or NH₄OH_(aq) + HCl_(aq) ==> NH₄Cl_(aq) + H₂O_(l)

• If ammonium salts are mixed with sodium hydroxide solution, free ammonia is formed (detected by smell and damp red litmus turning blue).

  • e.g. ammonium chloride + sodium hydroxide ==> sodium chloride + water + ammonia

  • NH₄Cl + NaOH ==> NaCl + H₂O + NH₃

• Ammonium sulphate or nitrate salts are widely used as 'artificial or synthetic fertilisers (preparation reactions above). There are several advantages to using artificial fertilisers in the absence of sufficient manure–silage etc. e.g. relatively cheap mass production, easily used to make poor soils fertile or quickly enrich multi–cropped fields.

• Artificial fertilisers are important to agriculture and used on fields to increase crop yields but they should be applied in a balanced manner.

  • Fertilisers usually contain compounds of three essential elements for healthy and productive plant growth to increase crop yield. They replace nutrient minerals used by a previous crop or enriches poor soil and more nitrogen gets converted into plant protein. 

    • Nitrogen (N) e.g. from ammonium or nitrate salts like ammonium sulphate, ammonium sulphate or ammonium phosphate or urea (e.g. look for the N in the formula of ammonium salts)

    • Phosphorus (P) e.g. from potassium phosphate or ammonium phosphate

    • Potassium (K) e.g. from potassium phosphate, potassium sulphate.

    • The fertiliser is marked with an 'NPK' value, i.e. the nitrogen : phosphorus : potassium ratio

  • Fertilisers must be soluble in water to be taken in by plant roots.

• Overuse of ammonia fertilisers on fields can cause major environmental problems as well as being uneconomic.
• Ammonium salts are water soluble and get washed into the groundwater, rivers and streams by rain contaminating them with ammonium ions and nitrate ions.
• This contamination causes several problems.
• Excess fertilisers in streams and rivers cause eutrophication.
  • Overuse of fertilisers results in appreciable amounts of them dissolving in rain water.
  • This increases levels of nitrate or phosphate in rivers and lakes.
  • This causes 'algal bloom' i.e. too much rapid growth of water plants on the surface where the sunlight is the strongest.
  • This prevents light from reaching plants lower in the water.
  • These lower plants decay and the active aerobic bacteria use up any dissolved oxygen.
  • This means any microorganisms or higher life forms relying on oxygen cannot respire.
  • All the eco–cycles are affected and fish and other respiring aquatic animals die.
  • The river or stream becomes 'dead' below the surface as all the food webs are disrupted.
• Nitrates are potentially carcinogenic (cancer or tumour forming).
  • The presence in drinking water is a health hazard.
  • Rivers and lakes can be used as initial sources for domestic water supply.
  • You cannot easily remove the nitrate from the water, it costs too much!
  • So levels of nitrate are carefully monitored in our water supply.
• Cost – The hydrogen for the Haber Process for manufacturing ammonia is usually obtained from hydrocarbon sources e.g. methane gas. Therefore, as oil becomes more scarce, the cost of producing 'artificial' fertilisers will increase.

The manufacture and uses of ammonia-nitric acid-fertilisers (and preparation of ammonium salts)

Use of NPK fertilisers and environmental problems

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(7) The Nitrogen Cycle for the gaseous element N₂(g)

• Nitrogen is an extremely important element for all plant or animal life! It is found in important molecules such as amino acids, which are combined to form proteins. Protein is used everywhere in living organisms from muscle structure in animals to enzymes in plants/animals.
• Nitrogen from the atmosphere:
  • Nitrifying bacteria, e.g. in the root nodules of certain plants like peas/beans (the legumes), can directly convert atmospheric nitrogen into nitrogen compounds in plants e.g. nitrogen => ammonia => nitrates which plants can absorb.
    • However, most plants can't do this conversion from nitrogen => ammonia, though they can all absorb nitrates, so the 'conversion' or 'fixing' ability might be introduced into other plant species by genetic engineering.
  • The nitrogen from air is converted into ammonia in the chemical industry, and from this artificial fertilisers are manufactured to add to nutrient deficient soils. However, some of the fertiliser is washed out of the soil and can cause pollution.
  • The energy of lightning causes nitrogen and oxygen to combine and form nitrogen oxides which dissolve in rain that falls on the soil adding to its nitrogen content.
    1. N_2(g) + O_2(g) ==> 2NO_(g), then 
    2. then 2NO_(g) + O_2(g) ==> 2NO_2(g) 
    3. NO_2(g) + water ==> nitrates_(aq) in rain/soil
    4. Incidentally, reactions 1. and 2. can also happen in a car engine, and NO₂ is acidic and adds to the polluting acidity of rain as well as providing nutrients for plants!
• Nitrogen recycling apart from the atmosphere:
  • Nitrogen compounds, e.g. protein formed in plants or animals, are consumed by animals higher up the food chain and then bacterial and fungal decomposers break down animal waste and dead plants/animals to release nitrogen nutrient compounds into the soil (e.g. in manure/compost) which can then be re–taken up by plants. 
• Nitrogen returned to the atmosphere:
  • However, denitrifying bacteria will break down proteins completely and release nitrogen gas into the atmosphere.

    Use of NPK fertilisers and environmental problems

    The nitrogen cycle

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(8) The chemistry of nitrogen oxides

• The equilibria between oxygen O₂, nitrogen (II) oxide NO, nitrogen(IV) oxide NO₂ and its dimer N₂O₄.

• NO₂ can be made from the irreversible  thermal decomposition of lead(II) nitrate in a pyrex boiling tube connected to a 100 cm³ gas syringe in a fume cupboard.

  • lead(II) nitrate ==> lead(II) oxide + nitrogen(IV) oxide + oxygen

  • 2Pb(NO₃)_2(s) ==> 2PbO_(s) + 4NO_2(g) + O_2(g)

• (a) 2NO_{2(g, brown)} 2NO_{(g, colourless)} + O_{2(g, colourless)}  (ΔH = +113 kJ mol^–1)

  • (a) The temperature effect can be observed by strongly heating the gases in the pyrex tube above 400^oC.

• (b) 2NO_{2(g, brown)} N₂O_{4(g, colourless)}  (ΔH = –58 kJ mol^–1)

  • (b) The temperature effect can be observed by cooling and warming below 100^oC.

  • (b) The pressure effect can be observed by sealing the cool gases in the gas syringe and compressing and decompressing it.

• Temperature and energy change (ΔH)

  • (a) Increases in temperature favours the endothermic decomposition of NO₂ to NO and O₂, so at high temperatures the brown colour fades.

  • (b) Decrease in temperature favours the exothermic formation of the dimer N₂O₄ from NO₂, so the brown colour fades on cooling the gas mixture.

• Gas pressure change (ΔV)

  • (a) Increase in pressure favours the LHS, more NO₂, because 2 mol gas <== 3 mol gas, so theoretically the mixture would get darker.

  • (b) Increase in pressure favours N₂O₄ formation from NO₂, 2 mol gas ==> 1 mol gas, so the mixture would get lighter in colour.

    • (b) This can be demonstrated by compressing/decompressing the gas mixture in the syringe to see the brown colour intensity increase/decrease.

    • In fact you can even see the dynamic equilibrium 'kinetics' in operation here. There is a time lag of about 1–2 seconds before the new equilibrium position is established as the 'imposed' colour intensity change becomes constant.

• Concentration change

  • (a) Theoretically an increase in O₂ would lead to decrease in NO and increase in NO₂, so the mixture would get darker.

  • (b) Increase in NO₂ would increase N₂O₄, but overall the colour would still be darker because not all of the 'extra' NO₂ can be converted to maintain the equilibrium.

• The formation of nitrogen(II) oxide at high temperature e.g. in a car engine
  • N_2(g) + O_2(g) 2NO_(g) (ΔH = +181 kJ mol^–1)
  • Temperature and energy change (ΔH)
    • Increase in temperature favours the endothermic formation of NO.
    • This reaction does not happen at room temperature but is formed at the high temperatures in car engines.
    • Unfortunately when released through the car exhaust, it cools to normal temperatures when NO irreversibly reacts with oxygen in air to form nitrogen(IV) oxide, NO₂,
    • 2NO(g) + O₂(g) ==> 2NO₂(g)
    • Nitrogen dioxide is acidic, a lung irritant and a reactive free radical molecule involved in the chemistry of photochemical smog not good!
  • Its concentration in car exhaust gases can be reduced, along with that of carbon monoxide, by using a catalytic converter.
  • Using platinum, and other transition metal, based catalysts, the following reaction can be made to take place producing harmless nitrogen and carbon dioxide.
    • 2NO(g) + 2CO(g) ==> N₂(g) + 2CO₂(g)

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(9) Some acid–base chemistry of ammonia

• Definition and examples of WEAK BASES

  • A weak base is only weakly or partially ionised in water e.g.

  • A good example is ammonia solution, which is only about 2% ionised :

    • NH_3(aq) + H₂O_(l) NH₄⁺_(aq) + OH^–_(aq)

      • Ammonia is the base and the ammonium ion NH₄⁺ is its conjugate acid.

      • Water is the acid and the hydroxide ion is its conjugate base.

      • This equilibrium is sometimes referred to as a base hydrolysis.

    • The low % of ionisation gives a less alkaline solution of lower pH than for strong soluble bases (alkalis), but pH is still > 7.

    • The concentration of water is considered constant and to solve simple problems, the base ionisation equilibrium expression is written as:

    • Kb =	[NH4+(aq)] [OH–(aq)]
      	––––––––––––––––––––––
      	[NH3(aq)]

    • K_b is the base ionisation/dissociation constant (mol dm^–3) for any base.

• Salts of weak bases and strong acids give acidic solutions.

  • e.g. ammonium chloride. The chloride ion is such a weak base that there is no acid–base reaction with water, but the ammonium ion is an effective proton donor. As a general rule, the conjugate acid of a weak base is quite strong. The result here is that ammonium salt solutions have a pH of 3–4.

  • NH₄⁺_(aq) + H₂O_(l) NH_3(aq) + H₃O⁺_(aq)

  • In zinc–carbon batteries an acidic ammonium chloride paste dissolves the zinc in the cell reaction, though an oxidising agent must be added (MnO₂) to oxidise the hydrogen formed into water, or batteries might regularly explode!

  • If you place a piece of magnesium ribbon or a zinc granule in ammonium chloride or ammonium sulphate solution you will see fizzing as hydrogen gas is formed.

    • 2H₃O⁺_(aq) + M_(s) ==> M²⁺_(aq) + H₂O_(l) + H_2(g)

    • M = zinc or magnesium

• Buffering action ammonium salts

  • A mixture of a weak base and the salt of the weak base with a strong acid.

  • e.g. ammonia NH₃ and ammonium chloride NH₄⁺Cl^–

  • NH₄⁺ and NH₃ constitute a conjugate acid–base pair.

  • In solution most of the ammonia is NOT ionised (and even suppressed by the ammonium ions from the salt).

    • It is the weak base that 'removes' most of any added hydrogen ions.

      • NH_3(aq) + H⁺_(aq) NH₄⁺_(aq)

  • The salt is fully ionised in solution giving relatively high concentrations of the ammonium ion.

    • It is the ammonium ion that removes most of any added hydroxide ions.

      • NH₄⁺_(aq) + OH^–_(aq) NH_3(aq) + H₂O_(l)

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 (10) The thermal decomposition of ammonium salts

• On heating strongly above 340^oC, the white solid ammonium chloride, thermally decomposes into a mixture of two colourless gases ammonia and hydrogen chloride.
• On cooling the reaction is reversed and solid ammonium chloride reforms.
  • This is an example of sublimation but here it involves both physical and chemical changes.
  • When a substance sublimes it changes directly from a solid into a gas without melting and on cooling reforms the solid without condensing to form a liquid.

  • Ammonium chloride + heat ammonia + hydrogen chloride

  • NH₄Cl_(s) NH_3(g) + HCl_(g)

  • so the thermal decomposition of ammonium chloride is the forward reaction, and the formation of ammonium chloride is the backward reaction.

• Note:

  • Reversing the reaction conditions reverses the direction of chemical change, typical of a reversible reaction.
  • Thermal decomposition means using 'heat' to 'break down' a molecule into smaller ones.
  • The decomposition is endothermic (heat absorbed or heat taken in) and the formation of ammonium chloride is exothermic (heat released or heat given out).
  • This means if the direction of chemical change is reversed, the energy change is also reversed.
  • Ammonium fluoride (>?^oC), ammonium bromide (>450^oC) and ammonium iodide (>550^oC), with a similar formula, all sublime in a similar physical–chemical way when heated, so the equations will be similar i.e. just swap F, Br or I for the Cl.
    • Similarly, ammonium sulphate also sublimes when heated above 235^oC and thermally decomposes into ammonia gas and sulphuric acid vapour.
      • (NH₄)₂SO_4(s) NH_3(g) + H₂SO_4(g)
    • Ammonium nitrate does not undergo a reversible sublimation reaction, it melts and then decomposes into nitrogen(I) oxide gas (dinitrogen oxide) and water vapour.
    • NH₄NO_3(s) N₂O_(g) + 2H₂O_(g)
    • This is very different reaction, in fact its an irreversible redox reaction. The nitrate ion, NO₃^–, or any nitric acid formed, HNO₃, act as an oxidising agent and oxidise the ammonium ion. If the products are cooled, ammonium nitrate is NOT reformed.
  • For more on sublimation, see the States of Matter webpage.

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(11) REDOX analysis of selected reactions - oxidation state changes

The oxidation of ammonia with molecular oxygen

• The concept of oxidation state can now be fully applied to reactions which do not involve ions e.g.

• The oxidation of ammonia via a Pt catalyst at high temperature which is part of the chemistry of nitric acid manufacture.

• 4NH_3(g) + 5O_2(g) ===> 4NO_(g) + 6H₂O_(g)  

• The oxidation number analysis is:

  • 4N at (–3) each in NH₃ and 10O all at (0) in O₂ change to ...

  • 4N at (+2) each in NH₃, 4O at (–2) each and 6 O at (–2) each in H₂O.

  • H is +1 throughout i.e. does not undergo an ox. state change.

  • Oxygen is reduced from ox. state (0) to (–2).

  • Nitrogen is oxidised from ox. state (–3) to (+2).

  • The total increase in ox. state change of 4 x (–3 to +2) for nitrogen is balanced by the total decrease in ox. state change of 10 x (0 to –2) for oxygen i.e. 20 e^– or ox. state units change in each case.

  • Oxygen is the oxidising agent (gain/accept e^–s, lowered ox. state) and ammonia is the reducing agent (loses e^–s, inc. ox. state of N).

The reaction between ammonium and nitrate(III) (nitrite) ions

• NH₄⁺_(aq) +  NO₂^–_(aq) ===> 2H₂O_(l) + N_2(g)

• Here its the opposite of disproportionation where two species of an element in different oxidation states react to produce one species of a single intermediate oxidation state.

• Ox. state changes: Nitrogen in a (–3) and a (+3) state both end up in the (0) state.

• Oxygen at (–2) and hydrogen (+1) remain unchanged in oxidation state.

• The nitrite ion acts as the oxidising agent and gets reduced (N +3 to 0, 3e's gained, decrease of 3 ox. state units)

• and the ammonium ion acts as the reducing agent and gets oxidised (N –3 to 0, 3 e's lost, inc. ox. state 3 units).

• The nitrite ion acts as the oxidising agent (gains/accepts e^–s, lowered ox. state of N) and the ammonium ion acts as the reducing agent (loses/donates e^–s, inc. ox. state of N).