Nitrate is a polyatomic ion with the chemical formula NO⁻ ₃. Salts containing this ion are called nitrates. Nitrates are common components of fertilizers and explosives. Almost all inorganic nitrates are soluble in water. An example of an insoluble nitrate is bismuth oxynitrate.
In the lewis structure, the single bonded oxygens all had 3 lone pairs, giving them 6 electrons. Their single bond with nitrogen means that they share 2 electrons with it–these electrons belong to both the oxygens and the nitrogen. These two electrons serve to complete the octet.
To draw the Lewis structure for the nitrate ion (NO3^-), follow these steps:
Count the Total Valence Electrons:
- Nitrogen (N) has 5 valence electrons.
- Oxygen (O) has 6 valence electrons, and there are three oxygen atoms, so 6 * 3 = 18 electrons.
- The ion has a -1 charge, which adds one more electron.
Total = 5 (from N) + 18 (from O) + 1 (for the charge) = 24 electrons.
Choose the Central Atom:
- Nitrogen is less electronegative than oxygen and can form more bonds, so it is the central atom.
Draw a Skeleton Structure:
- Connect the nitrogen to the three oxygen atoms with single bonds. This uses up 6 electrons (2 per bond).
Distribute Remaining Electrons:
- You have 24 – 6 = 18 electrons left.
- Place these electrons around the oxygen atoms to complete their octets. Each oxygen will get 6 more electrons in pairs (three pairs each), which accounts for 18 electrons.
Check Octets:
- Each oxygen now has an octet (8 electrons around each, counting the bond to nitrogen).
- Nitrogen has only 6 electrons from the bonds; it needs 2 more.
Form Double Bonds:
- Since all oxygens have their octet but nitrogen doesn’t, convert one of the single bonds to a double bond. This can be done with any of the oxygen atoms. However, in reality, nitrate exhibits resonance, meaning the double bond can be on any of the three N-O positions. Here, we’ll arbitrarily choose one:
- N=O
- N-O
- N-O
Now, nitrogen shares 8 electrons, fulfilling its octet.
Resonance Structures:
- Draw all equivalent resonance structures by moving the double bond between nitrogen and the different oxygen atoms. This shows that the nitrate ion has three equivalent resonance structures.
Formal Charges:
- To ensure the structure is the most stable, calculate formal charges:
- For N in any structure: Formal charge = 5 (valence e-) – 0 (unshared e-) – 8/2 (shared e-) = +1
- For O with single bond: 6 – 6 – 1 = -1
- For O with double bond: 6 – 4 – 4/2 = 0
The structure where one oxygen has a double bond and two have single bonds with the formal charges (+1 on N, -1 on one O, 0 on the other two O’s) is correct, but remember, due to resonance, these charges are delocalized over all three oxygen atoms.
Here’s how you might depict one of the Lewis structures (without drawing, which I can’t do here):
Remember, in practice, you would show all three resonance forms with double-headed arrows between them, indicating that the actual structure is an average of these forms.
Hybridization & Geometry
- Hybridization: The nitrogen atom in NO3^- undergoes sp² hybridization. This is because it forms three sigma bonds (one with each oxygen atom) and has no lone pairs on the nitrogen, which leaves one p-orbital available for π bonding.
- Geometry: The molecular geometry of NO3^- is trigonal planar because all four atoms (N and three O’s) lie in the same plane with bond angles of approximately 120 degrees. This shape arises from the sp² hybridization and the arrangement to minimize electron repulsion.