Nitrogen formal charge

It will always be true that a nitrogen with four covalent bonds has a formal charge of + 1 (A nitrogen with four co valent bonds cannot have unshared pairs because of the octet rule)... 

Moving now to nitrogen we see that it has four covalent bonds (two single bonds + one double bond) and so its electron count is 5(8) = 4 A neutral nitrogen has five electrons m its valence shell The electron count for nitrogen m nitric acid is one less than that of a neutral nitrogen atom so its formal charge is +1... 

The green oxygen m Figure 1 5 owns three unshared pairs (six electrons) and shares two electrons with nitrogen to give it an electron count of seven This is one more than the number of electrons m the valence shell of an oxygen atom and so its formal charge is —1... 

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

The electron counts of nitrogen in ammonium ion and boron in borohydride ion are both 4 (half of eight electrons in covalent bonds) Because a neutral nitrogen has five electrons in its valence shell an electron count of 4 gives it a formal charge of +1 A neutral boron has three valence electrons so that an electron count of 4 in borohydride ion corresponds to a formal charge of -1... 

Both protonated forms place the formal positive charge on one of the nitrogens. Is charge delocalization more effective for one of the structures over the other, making it the more stable ... 

The same is true for the nitrogen atom in ammonia, which has three covalent N-H bonds and two nonbonding electrons (a lone pair). Atomic nitrogen has five valence electrons, and the ammonia nitrogen also has five—one in each of three shared N-H bonds plus two in the lone pair. Thus, the nitrogen atom in ammonia has no formal charge. 

For such odd electron species (sometimes called free radicals) it is impossible to write Lewis structures in which each atom obeys the octet rule. In the NO molecule, the unpaired electron is put on the nitrogen atom, giving both atoms a formal charge of zero ... 

EXERCISE 1.33 Consider the nitrogen atom in the structure below and determine if it has a formal charge ... 

PROBLEMS For each of the structures below determine if the oxygen or nitrogen atom has a formal charge. If there is a charge, draw the charge. 

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 ... 

Now let s look at the common situations for nitrogen atoms. When nitrogen has no formal charge, it will have three bonds and one lone pair ... 

If nitrogen has a negative formal charge, then it must have two bonds and two lone pairs ... 

For each of the remaining nitrogen atoms, there is a negative formal charge. That means that each of those nitrogen atoms has one extra electron, 5 + 1=6 electrons. Each nitrogen atom has two bonds, which means that each nitrogen atom has... 

C09-0108. Carbon, nitrogen, and oxygen form two different polyatomic ions cyanate ion, NCO, and isocyanate ion, CNO". Write Lewis stmctures for each anion, including near-equivalent resonance structures and indicating formal charges. 

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. 

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