Draw Lewis Dot structures for N2H4, N2, and N2F2 in a sheet of paper. Please answer the following questions here. a) Arrange the molecules in the increasing order of their nitrogen-nitrogen bond length. b) Which molecule has the strongest nitrogen-nitrogen bond?
The Correct Answer and Explanation is:
Correct Answers
a) Increasing order of nitrogen-nitrogen bond length:
N₂ < N₂F₂ < N₂H₄
b) The molecule with the strongest nitrogen-nitrogen bond is: N₂
300-Word Explanation
The bond length between atoms is influenced by the type and order of the bond. A triple bond is shorter and stronger than a double bond, which in turn is shorter and stronger than a single bond. Among the three given molecules — N₂H₄, N₂, and N₂F₂ — the bonding between the two nitrogen atoms varies in bond order, leading to differences in bond lengths and strengths.
In N₂ (dinitrogen), the two nitrogen atoms are connected by a triple bond. This triple bond consists of one sigma bond and two pi bonds. It results in a very short and strong bond due to the high electron density shared between the atoms. As a result, N₂ has the shortest and strongest nitrogen-nitrogen bond.
N₂F₂ (difluorodiazene) has a double bond between the two nitrogen atoms. It has one sigma and one pi bond, resulting in a bond length longer than that of N₂ but shorter than a single bond. The presence of electronegative fluorine atoms may affect electron distribution but does not change the bond order between the nitrogen atoms.
N₂H₄ (hydrazine) contains a single bond between the two nitrogen atoms. This bond is the longest and weakest among the three molecules because only one pair of electrons is shared. The lone pairs on each nitrogen may also cause repulsion, slightly increasing the bond length further.
Thus, the nitrogen-nitrogen bond lengths increase in the order N₂ < N₂F₂ < N₂H₄, and the bond strength decreases in the same order. Therefore, N₂ has the strongest nitrogen-nitrogen bond, a direct consequence of its triple bond.
