Draw all possible resonance structures for NO2, NO2, and NO. Use the resonance structures to solve the problems below.
(b) Match each species with the number of covalent bonds predicted by Lewis structures to exist between an N atom and an O atom bonded to this N atom. (Hint: Average the number of bonds between N and an attached oxygen atom in a particular position using all of your resonance structures for the species that you are working on.)
NO:
NO2:
NO:
bonds between N and O.
bonds between N and O.
bonds between N and O.
(c) Match each species with the correct formal charge on the central N atom.
NO:
NO:
NO:
charge on N.
charge on N.
charge on N.
(d) Match each species with the average formal charge on an outside oxygen atom predicted by Lewis structures.
NO:
NO
NO2:
average charge on outside O atom.
average charge on outside O atom.
average charge on outside O atom.
The Correct Answer and Explanation is :
Resonance Structures of NO, NO₂⁺, and NO₂⁻
- Nitric oxide (NO)
NO has one major Lewis structure with a single bond between N and O and an unpaired electron on nitrogen. - Nitronium ion (NO₂⁺)
NO₂⁺ has two resonance structures, each with a double bond between N and one O, with formal charges distributed. - Nitrite ion (NO₂⁻)
NO₂⁻ has two resonance structures, each with one single and one double bond between N and O, leading to an average bond order.
(b) Number of Bonds Between N and O
Using resonance structures, we calculate the bond order:
- NO: The molecule has a bond order of 2.5 (one single bond and one partial double bond due to the unpaired electron).
- NO₂⁺: Since both resonance structures contribute equally and both have double bonds, the bond order is 2.0.
- NO₂⁻: Each resonance structure has one single bond and one double bond. The average bond order is 1.5.
Final Answer:
- NO: 2.5 bonds
- NO₂⁺: 2.0 bonds
- NO₂⁻: 1.5 bonds
(c) Formal Charge on the Central N Atom
Formal charge = (Valence electrons in free atom) – (Nonbonding electrons) – (1/2 * Bonding electrons)
- NO: +1
- NO₂⁺: +1
- NO₂⁻: 0
Final Answer:
- NO: +1 charge on N
- NO₂⁺: +1 charge on N
- NO₂⁻: 0 charge on N
(d) Average Formal Charge on Outside O Atoms
Formal charge on O = (6 valence electrons) – (nonbonding electrons) – (1/2 * bonding electrons)
- NO: The oxygen has a formal charge of -1.
- NO₂⁺: Each oxygen is 0.
- NO₂⁻: Each oxygen has -0.5 due to resonance.
Final Answer:
- NO: -1 average charge on outside O
- NO₂⁺: 0 average charge on outside O
- NO₂⁻: -0.5 average charge on outside O
Explanation (300 Words)
The resonance structures of NO, NO₂⁺, and NO₂⁻ help us determine bond orders and formal charges. Nitric oxide (NO) is a radical molecule, meaning it has an unpaired electron. Its Lewis structure shows a bond between nitrogen and oxygen that is stronger than a double bond but not quite a triple bond, leading to a bond order of 2.5. The formal charge on nitrogen is +1, and the oxygen has a formal charge of -1.
For nitronium ion (NO₂⁺), resonance occurs between two equivalent structures where nitrogen forms a double bond with each oxygen. Since both resonance structures contribute equally, the bond order is 2.0, meaning each N-O bond is a double bond. The nitrogen carries a +1 formal charge, while each oxygen remains neutral with 0 formal charge.
For nitrite ion (NO₂⁻), resonance also occurs between two structures where nitrogen forms one single bond and one double bond with oxygen, but they switch positions. This results in an average bond order of 1.5. The nitrogen has a formal charge of 0, while each oxygen gets an average formal charge of -0.5 due to delocalization of negative charge.
In summary, resonance is crucial in predicting bond strengths and charge distributions. NO has a high bond order due to its unpaired electron, NO₂⁺ has stable double bonds, and NO₂⁻ exhibits delocalized bonding, reducing the formal charge on each oxygen.