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The chemistry of Nitrogen. Chapter 16. Nitrogen. Nitrogen can complete its valence valence shell by: 1.) Electron gain: N 3- ion This is found in saltlike nitrides. 2.) formation of electron pair bonds: A) single bonds NH3 B) multiple bonds :N≡N: ; -N=N-, or NO 2

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  • Nitrogen can complete its valence valence shell by:
  • 1.) Electron gain: N3- ion
      • This is found in saltlike nitrides.
  • 2.) formation of electron pair bonds:
      • A) single bonds NH3
      • B) multiple bonds :N≡N: ; -N=N-, or NO2
  • 3.) formation of electron pair bonds with electron gain, NH2- or NH2-
  • 4.) Formation of electron pair bonds with electron loss (substituted ammonium ions)
  • Three-Covalent Nitrogen
  • NR3 molecules are sp3 hybridised, the lone pair occupies the fourth position.
  • 1.) all NR3 compounds behave as Lewis bases, give donor-acceptor complexes with lewis-acids, act as ligands towards transition metal ions [Co(NH3)6]3+
  • 2.) Pyramidal molecules (NRR’R’’) should be chiral. Optical isomers can not be isolated, because N oscillates through the plane of the R-groups. The energy barrier is only 24kJ/mol. (Inversion)
  • 3.) in few cases 3-covalent nitrogen is planar;
  • N-N single bond energy
  • The difference between C and N in bonding energies is attributable to the effects of repsulsion between nonbonding lone pairs. Nitrogen has little tendency to catenation.
multiple bonds
Multiple bonds
  • Nitrogens propensity to form pπ- pπ multiple bonds is a feature that distinguishes it from phosphorus and the other GroupVB elements.
  • N2 has a high bond strength and a short internuclear distance (1.094Å). P forms infinite layer structures with only single bonds or P4 molecules.
  • The oxo anions NO2- and NO3- , multiple bonds may be formulated in either resonance or MO terms.
  • Nitrogen occurs as dinitrogen. N2 (bp 77.3 K).
  • 78% of the atmosphere is N2
  • N14/N15 has a ratio of 272.
  • N15 compounds are used in tracer studies.
  • The NN triple bond is responsible for the inert behaviour of N2.
  • N2 is prepared by liquefaction and fractionation of air.
  • N2 only reacts with Li to give Li3N.
  • With certain transition metal complexes oand with nitrogen fixing bacteria.
  • Typical reactions of N2 at elevated temperatures :
  • Nitrides of eletropositives metals have structures with discrete nitrogen atoms and can be regarded as ionic (Ca2+)3(N3-)2
  • The Nitrides hydrolyse to ammonia and metal hydroxides.
  • Preparation:
  • Direct interaction
  • Loss of ammonia from amides on heating
  • Transition metal nitrides are often nonstoichiometric and have nitrogen atoms in the interstices of close-packed arrays of metal atoms.
  • They are like the carbides or borides hard, chemically inert, high melting and electrically conducting.
  • Numerous covalent nitrides (BN,S4N4,P3N5)
  • These nitrides have very differing properties, depending on the element.
nitrogen hydrides
Nitrogen Hydrides
  • Ammonia is formed by the action of a base on an ammonium salt:
  • Industrially Ammonia is made by the haber-Bosch process at 400-500 deg C and 100-1000atm.
nitrogen hydrides1
Nitrogen hydrides
  • Ammonia is a colorless gas.
  • In liquid form it has a high heat of evaporation .
  • Liuid ammonia resembles water in its physical behaviour. It forms strong nydrogen bonds.
  • Its dielectric constant is around 22 at -34degC.
  • Liquid ammonia has lower reactivity towards electropositive metals and dissolves many of them.
  • AgI is insoluble in water but soluble in ammonia.
  • Ammonia burns in air:
nitrogen hydrides2
Nitrogen hydrides
  • At 750-900 deg C in the presence of a catalyst (platinum, platinum-rhodium) :
  • NO reacts on with O2 to form the mixed oxides which can be absorbed in water to form nitric acid.
nitrogen hydrides3
Nitrogen hydrides
  • The sequence in industrial utilisation of atmospheric nitrogen is
  • Ammonia is extremely soluble in water.
ammonium salts
Ammonium salts
  • Ammonium salts
  • Crystalline salts of ammonium are mostly water soluble.
  • Ammonium salts generally resemble those of potassium and rubidium in solubility and structure. The three ions have comparable radii.
  • Hydrazine can be described as a reaction of ammonia with one ammonia as the substituent.
  • 2 series of hydrazinium salts can be obtained:
  • N2H5+ are stable in water
  • N2H62+ are hydrolysed in water.
  • Anhydrous hydrazine is a fuming colorless liquid. It is considerable stable and burns in air
  • Aqueous hydrazine is a powerful reducing agent in basic solution.
  • Hydrazine is synthesized by the inateraction of aqueous ammonia with sodium hypochlorite
  • But there is a competing reaction when hydrazine first is formed:
  • To prevent this reaction one needs to add gelatine. It complexes Cu2+ ions better than EDTA.
  • Hydroxylamine is a weaker base than NH3:
  • It is prepared by reduction of nitrates or nitrites either electrolytically or with SO2 under controlled conditions.
  • Hydroxylamine is a white unstable solid.
  • It is used as a reducing agent.
  • Heavy metal azides are explosive and lead or mercury azides have been used in detonation caps.
  • The pure acid is a dangerously explosive liquid.
  • It can act as a ligand in metal complexes, it is linear molecule.
nitrogen oxides
Nitrogen oxides
  • Dinitrogen monoxide
  • It has a linear structure is realtively unreactive , is inert towards:
  • Halogens,
  • Alkali metals
  • Ozone at RT.
  • It is used as an anaesthetic.
nitrogen oxides1
Nitrogen oxides
  • Nitrogen monoxide
nitrogen oxides2
Nitrogen oxides
  • Dinitrogen trioxide
  • The anhydride of nitrous acid
phosphorous arsen anitmony bismuth
Phosphorous, Arsen, Anitmony, Bismuth
  • Phosphorous occurs in minerals of the apatite family.
  • As, Sb,Bi occur mainly as sulfide minerals.
  • The electron configuration is ns2np3.
  • P and N are very different in their chemistry.
  • P is a true non metal, down the period the metallic trend is increasing.
  • Differences between N and P:
  • Diminished ability to form pπ- pπ multiple bonds
  • The possibility to use the lower 3d orbitals
  • Nitrogen forms esters, phosphorus gives P(OR)3. Nitrogen oxides and oxoacids involve multiple bonds, whereas the phosphorus oxides have single bonds. Phosphoric acid PO(OH)3 in contrast NO2(OH).
  • Phosphorus is obtained by reduction of phosphate rocks.
  • Phosphorus distills and is condensed in water.
  • White P is stored under water to protect from air.
  • Red and black P are stable in air, burn on heating.
  • P is soluble in organic solvents.
  • As,Sb,Bi are obyained by reduction of the oxides with carbon or Hydrogen.
  • All elements react readily with halogens.
  • Nitric acid  Phosphoric acid, arsenic acid, Sb trioxide and Bi nitrate.
  • Interactions with metals gives phosphides, arsenides, ....
  • GaAs has semiconductor properties.
  • The stability of the hydrides decreases down the period.
  • Sb and Bi hydrides are very unstable.
  • Phosphine is made from the reaction of acids with zinc phosphide.
  • Phosphine is a nerve toxin..
  • Trihalides are obtained by direct reaction with halogens.
  • They rapidly hydrolize in water
  • Gaseous molecules have pyramidal structure.
  • Iodides of As,SB,and Bi have layer structures based on hexagonal closed packing of iodine atoms with the group VB elements.
  • Phosphorus trifluoride is a colorless toxic gas.
  • It is slowly attacked by water and rapidly by alkali.
  • Phosphorus pentoxide (P4O10)
  • It is used as one of the most effective drying agents. Reacts with water to form phosphoric acid
oxo acids
Oxo acids
  • Phosphoric acid:
  • PCl3 or P4O6 are hydrolised in water
  • Phosphorus acid Hypophosphorus acid
oxo acids1
Oxo acids
  • Orthophosphoric acid
  • Is the oldest known phosphorus compounds. It is a syrupy liquid made by direct reaction of ground phhosphate rock with sulfuric acid.
  • The pure acid is a colorless cyrstalline solid.
  • Stable and has no oxidising properties below 350-400 degC.
  • It will attack quartz.
  • Hydrogen bonding persists in the concentrated solution and is respåonsible ofr the syrupy behaviour.