Sulphur and phosphorus bear similar relationships to nitrogen and oxygen.
The principal differences between S/P and O/N are as follows.
Oxidation states, coordination numbers and stereochemistries
Valence shell expansion, and the resultant formation of species like PX5, PX6–, R3P=CR2, and SF6, SF5Cl, SF4 have no N or O counterparts, the highest coordination being NX3 and OX2.
This is due to the increased size of the second row elements relative to those of the first row, and also the availability of the empty 3d orbitals to take part in bonding (there are no 2d orbitals).
There are only very weak double bonds formed by S and P, eg. S=C=S and RP=PR, whereas there are many examples of double bond formation with O and N.
N/O from strong bonds to electropositive elements, whereas P/S form strong bonds to electronegative elements, eg. NH3/OH2 compared to PF5/SF6.
N(PF2)3 and N(SCF3)3 are planar at the central N rather than pyramidal like N(CH3)3. This is due to the presence of favourable Npπ-Sdπ, or Npπ-Pdπ, bonding interactions, which require a planar orientation, rather than obeying the VSEPR rules in N(CH3)3 when this bonding cannot occur due to the absence of 2d-orbitals in the first row elements.
Acidity/Basicity of PX3 and SX2
PX3 and SX2 have HOMO‘ s at higher energies than NX3 and OX2.
PX3 and SX2 show π-acid properties associated with the electron pair acceptor nature of the vacant 3d orbitals.
Therefore, PX3 resembles CO in its bonding with transition metals. There is no such bonding with N.
The relative strength of single bonds in the second row of the periodic table compared to those in the first row is shown in the instability of compounds such as N2H4 and H2O2. However, the multiple formation of single bonds is an important feature of the chemistry of P and S, eg. in compounds like cyclic-(RP)n.
Typical compounds formed
Hydrogen bonding occurs for N/O but not for P/S.
The H-E-H bond angles of EH2/EH3 reflect the increased p-character of the E-H bonds formed by P/S compared to N/O.
OH2/NH3 are more basic than SH2/PH3.
The stability of the highest oxidation states increases down the groups. Thus we see PF5/SF6 but only NF3/OF2.
Nitrogen forms oxides N2O up to N2O5, but Phosphorus forms oxides P4O6 and P4O10.
This reflects the change in relative strengths of the single and double bonds in going down the group, ie. B(N=O)>2B(N-O) but 2B(P-O)>B(P=O). The is also reflected by the increased tendency for catenation to occur.