For the overlap efficiency is close to 1.0 (100% -bonds overlap), which nicely explains why are stronger -bonds usually than -bonds. Overlap efficiency also applies to between atoms. Bond -bonds strengths can vary as the bond overlap spreads out over a larger region of space. This is not simple to qualitatively predict. For example, the dihalogen bond strengths increase, then decrease as follow: F2 < Cl2 > Br2 > I2. They steadily decrease as one goes from C-C > Si-C > Ge-C > Sn-C > Pb-C. For most transition metal M-M single bonds the trend is fairly consistent: first row < second row < third row. But for M-M quadruple bonds one has: Cr-Cr << Mo-Mo > W-W. In the case of strong multiple bonding and shorter bond distances one can get TOO much overlap between orbitals. This can lead -bonding to destabilizing anti-bonding type interactions that weaken the -bond. N2 is a good example of this: The orbital is higher in energy (& the HOMO) than the -bonding p- orbitals. Bringing two orbitals together to the N≡N bond distance p- leads to anti-bonding interactions due to the extending over into p-lobe the other phased orbital region as shown qualitatively below: To minimize this unfavorable antibonding interaction N2 forms a sp hybrid orbital combination that reduces the overlap in favor -bonding of the two stronger (overall) -bonds. Square Planar
For -bonds the overlap efficiency is close to 1.0 (100% overlap), which nicely explains why -bonds are usually stronger than -bonds. Overlap efficiency also applies to -bonds between atoms. Bond strengths can vary as the bond overlap spreads out over a larger region of space. This is not simple to qualitatively predict. For example, the dihalogen bond strengths increase, then decrease as follow: F2 < Cl2 > Br2 > I2. They steadily decrease as one goes from C-C > Si-C > Ge-C > Sn-C > Pb-C. For most transition metal M-M single bonds the trend is fairly consistent: first row < second row < third row. But for M-M quadruple bonds one has: Cr-Cr << Mo-Mo > W-W. In the case of strong multiple bonding and shorter bond distances one can get TOO much overlap between -bonding orbitals. This can lead to destabilizing anti-bonding type interactions that weaken the -bond. N2 is a good example of this: The p- orbital is higher in energy (& the HOMO) than the -bonding orbitals. Bringing two p- orbitals together to the N≡N bond distance leads to anti-bonding interactions due to the p-lobe extending over into the other phased orbital region as shown qualitatively below: To minimize this unfavorable antibonding interaction N2 forms a sp hybrid orbital combination that reduces the -bonding overlap in favor of the two stronger (overall) -bonds. Square Planar