The chemistry of Nitrogen

1. The chemistry of Nitrogen Chapter 16

2. Nitrogen • 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)

3. Nitrogen • 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)

4. Inversion of Nitrogen

5. Nitrogen • 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.

6. 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.

7. Nitrogen • 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.

8. Nitrogen • 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 :

9. Nitrides • 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

10. Nitrides • 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.

11. 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.

12. 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:

13. 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.

14. Nitrogen hydrides • The sequence in industrial utilisation of atmospheric nitrogen is • Ammonia is extremely soluble in water.

15. 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.

16. Thermal decomposition of ammonium salts

17. Hydrazine • 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.

18. Hydrazine • 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

19. Hydrazine • 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.

20. Hydroxylamine • 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.

21. Azides • 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.

22. 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.

23. Nitrogen oxides • Nitrogen monoxide

24. Nitrogen oxides • Dinitrogen trioxide • The anhydride of nitrous acid

25. Summary of reactions

26. Summary of reactions

27. Summary of reactions

28. 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.

29. Phosphorus • 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).

30. Elements • 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.

31. Elements • 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.

32. Hydrides • 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..

33. Halides • 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.

34. Important reactions of P and Halides

35. Oxides • Phosphorus pentoxide (P4O10) • It is used as one of the most effective drying agents. Reacts with water to form phosphoric acid

36. Oxo acids • Phosphoric acid: • PCl3 or P4O6 are hydrolised in water • Phosphorus acid Hypophosphorus acid

37. 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.

38. Summary of Group trends

41. Summary of reactions of P4

42. Summary of reactions of PCl5