JEE Chemistry The P Block Elements Notes Set 22

Group 15: Occurrence, Trends and Anomalous Nitrogen

The group 15 elements are N, P, As, Sb and Bi. Nitrogen forms about 78 percent of air by volume; phosphorus occurs as phosphate rock, \( \ce{Ca3(PO4)2} \) and fluorapatite \( \ce{Ca5(PO4)3F} \), also in bones; arsenic, antimony and bismuth occur as sulphides. The valence shell is \( \ce{ns^2 np^3} \) (half-filled p subshell, so ionisation enthalpies are high), and the common oxidation states are -3, +3 and +5. On going down the group the +3 state becomes more stable than +5 because of the inert pair effect (Bi(V) is a strong oxidant, Sb(V) a weaker one).
Trends: size, metallic character, density and melting/boiling points (except for N) increase down the group; ionisation enthalpy and electronegativity decrease. The covalent radius rises sharply from N to P but only slightly from Sb to Bi because the filled d and f subshells shield poorly.
Nitrogen differs from the rest: it is a diatomic gas with an extremely strong \( \ce{N#N} \) bond (941 kJ/mol), forms p pi-p pi multiple bonds because of its small size and high electronegativity, has no d-orbitals so its maximum covalency is 4 (no \( \ce{NF5} \) or \( \ce{NCl5} \), while \( \ce{PF5} \) and \( \ce{PCl5} \) exist), forms strong hydrogen bonds in \( \ce{NH3} \), and catenates only weakly. Phosphorus exists as a tetrahedral \( \ce{P4} \) molecule.

Allotropes of Phosphorus

Preparation of the Elements

1. Nitrogen: industrially by liquefaction and fractional distillation of air. In the laboratory, \( \ce{NH4Cl + NaNO2 -> N2 + NaCl + 2H2O} \), or \( \ce{(NH4)2Cr2O7 -> N2 + Cr2O3 + 4H2O} \), or \( \ce{2NaN3 -> 2Na + 3N2} \) (very pure). Nitrolim is \( \ce{CaCN2 + C} \), made as \( \ce{CaC2 + N2 -> CaCN2 + C} \).
2. Phosphorus: \( \ce{2Ca3(PO4)2 + 6SiO2 + 10C -> 6CaSiO3 + 10CO + P4} \) in an electric furnace; the vapour is condensed under water.

Compounds of Nitrogen

Ammonia: Haber process \( \ce{N2 + 3H2 <=> 2NH3} \) (about 200 atm, 700 K, iron catalyst with molybdenum promoter); in the laboratory \( \ce{2NH4Cl + Ca(OH)2 -> 2NH3 + CaCl2 + 2H2O} \), dried over quicklime (not \( \ce{H2SO4} \), \( \ce{P2O5} \) or \( \ce{CaCl2} \)). It is a pyramidal Lewis base, very soluble in water, forms deep blue \( \ce{[Cu(NH3)4]^2+} \), reduces hot CuO (\( \ce{3CuO + 2NH3 -> 3Cu + N2 + 3H2O} \)) and burns in oxygen. Hydrazine \( \ce{N2H4} \) and hydrazoic acid \( \ce{HN3} \) are other hydrides.
Ammonium salts: those with non-oxidising anions give ammonia on heating (\( \ce{NH4Cl -> NH3 + HCl} \)), those with oxidising anions give nitrogen or its oxides (\( \ce{NH4NO2 -> N2 + 2H2O} \), \( \ce{NH4NO3 -> N2O + 2H2O} \), dichromate gives \( \ce{N2 + Cr2O3} \)).
Oxidation states and oxides of nitrogen:

Nitric acid by Ostwald's process: \( \ce{4NH3 + 5O2 -> 4NO + 6H2O} \), \( \ce{2NO + O2 -> 2NO2} \), \( \ce{3NO2 + H2O -> 2HNO3 + NO} \). Reactions: with Zn, dilute acid gives \( \ce{N2O} \) and concentrated acid gives \( \ce{NO2} \); Cu gives NO (dilute) or \( \ce{NO2} \) (concentrated); Fe, Al and Cr are passivated by cold concentrated acid; Sn gives hydrated \( \ce{SnO2} \); non-metals are oxidised, e.g. \( \ce{S + 6HNO3 -> H2SO4 + 6NO2 + 2H2O} \). Aqua regia (3 HCl : 1 HNO3) dissolves gold and platinum. Concentrated nitric acid turns yellow-brown on standing because of dissolved \( \ce{NO2} \). Nitrous acid \( \ce{HNO2} \) is unstable, acts as an oxidant and a reductant, and disproportionates: \( \ce{3HNO2 -> HNO3 + 2NO + H2O} \).

Compounds of Phosphorus

Phosphine \( \ce{PH3} \): \( \ce{P4 + 3NaOH + 3H2O -> PH3 + 3NaH2PO2} \) (white P with hot alkali) or \( \ce{Ca3P2 + 6H2O -> 3Ca(OH)2 + 2PH3} \). It is a colourless, very poisonous gas with a fishy smell, a weaker base than ammonia (forms \( \ce{PH4I} \) and \( \ce{PH4Br} \) only with anhydrous HI and HBr), burns in air to give \( \ce{P4O10} \), and gives black \( \ce{Cu3P2} \) with copper sulphate. Traces of \( \ce{P2H4} \) make it inflame spontaneously.
Halides: \( \ce{P4 + 6Cl2 -> 4PCl3} \) (excess P) and \( \ce{P4 + 10Cl2 -> 4PCl5} \) (excess chlorine). \( \ce{PCl3 + 3H2O -> H3PO3 + 3HCl} \) (trigonal pyramidal, fumes in moist air); \( \ce{PCl5 + H2O -> POCl3 + 2HCl} \) and then \( \ce{POCl3 + 3H2O -> H3PO4 + 3HCl} \). In the solid state PCl5 is ionic, \( \ce{[PCl4]+[PCl6]-} \), and in the vapour it is trigonal bipyramidal (sp3d). \( \ce{PCl5} \) converts R-OH and carboxylic acids into chlorides.
Oxides: \( \ce{P4O6} \) (in limited air, forms \( \ce{H3PO3} \) with cold water) and \( \ce{P4O10} \) (in excess air; a very strong drying agent that dehydrates \( \ce{HNO3} \) and \( \ce{H2SO4} \)).

Acidity and reducing power: reducing power falls \( \ce{H3PO2 > H3PO3 > H3PO4} \) (number of P-H bonds 2, 1, 0). Basicity equals the number of P-OH groups, so \( \ce{H3PO3} \) is dibasic.

Worked Examples

Example 1. How many moles of NaOH are needed to neutralise the products of complete hydrolysis of one mole of \( \ce{PCl5} \)?
Answer:
PCl5 + 4H2O gives \( \ce{H3PO4} \) + 5HCl. Phosphoric acid is tribasic (3 mol) and HCl needs 5 mol: total 8 mol of NaOH.

Example 2. Find the oxidation number of P in \( \ce{H4P2O6} \) and in \( \ce{H4P2O7} \).
Answer:
In \( \ce{H4P2O6} \): 4 + 2x - 12 = 0, so x = +4. In \( \ce{H4P2O7} \): 4 + 2x - 14 = 0, so x = +5.

Example 3. Why does phosphine not form a salt with aqueous HCl although ammonia does?
Answer:
The lone pair on phosphorus sits in a larger, more diffuse orbital than the one on nitrogen (and N-H bonds form hydrogen bonds with water), so it is much less available for protonation and \( \ce{PH3} \) is a far weaker base. Its phosphonium salts exist only with anhydrous HBr and HI, and they hydrolyse at once in water.