Formal charge on nitrogen

Verify that the formal charges on nitrogen in ammonium ion... 

Because HCN has formal charges of zero on all the atoms, whereas HNC has a formal charge on nitrogen of +1 and a formal charge on carbon of — 1, we would expect HCN to be more stable than HNC. Indeed, this is the case. We will encounter many cases of constitutional isomers in which the difference in stability is much less than in this example. In such cases, much more subtle arguments must be used to predict which is more stable, if a prediction can be made at all without experimental measurements. 

The W—N bond of 174.3(15) pm is a little longer than that in [Cl5WNC2Cl5] , which was considered to be a triple bond. Since the geometry at nitrogen is almost linear [176.9(14)°], a triple-bonded resonance extreme with a positive formal charge on nitrogen may make a significant contribution. 

The formal charge is zero if the number of unshared electrons, plus the number of shared electrons divided by two, is equal to the number of valence shell electrons in the neutral atom (as ascertained from the group number in the periodic table). As the number of bonds formed by the atom increases, so does the formal charge. Thus, the formal charge of nitrogen in (CH3)j N is zero, but the formal charge on nitrogen in (CH3)4N is -l-l. 

For nitrogen valence electrons of free atom subtract assigned electrons Formal charge on nitrogen Overall charge on Ion = 4(0) + 1 = +1... 

From all of the cases above (oxygen, nitrogen, carbon), you can see why you have to know how many lone pairs there are on an atom in order to figure out the formal charge on that atom. Similarly, you have to know the formal charge to figure out how many lone pairs there are on an atom. Take the case below with the nitrogen atom shown ... 

In this case, there is a simple experiment that will determine whether this is correct. Structure I places a negative formal charge on the terminal nitrogen atom, while structure II places a negative formal charge on the oxygen atom on the opposite end of the molecule. If the two structures contribute equally, these effects should cancel, which would result in a molecule that is not polar. In fact, the dipole moment of N20 is only 0.17 D, so stmctures I and II must make approximately equal contributions. 

The resulting structure has a formal charge of 1 - on the terminal nitrogen atom and a 1 + formal charge on the NH2 nitrogen atom. 

The features of this resonance structure could be taken into consideration in the over-all description of the molecule, but it would not carry nearly the same weight as the previous resonance structures. For instance, in averaging the formal charges the +2 charge on nitrogen would not be taken as one out of four, but as much less than that. 

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