PCl
and PCl
exist. You can draw their Lewis structures. NF
exists. Does NF
exist? Why or why not? 4. In the molecule, ClO
, what is the formal charge on chlorine in a resonance structure with one double bond? a. 0 b. +1 c. -1 d. +2 e. -2 5. Which of the following molecules would you expect could overviolate the octet rule? a. PCl
b. SF
c. sulfate ion d. all the above (a, b, and c) e. none of a, b, or c
The Correct Answer and Explanation is:
1. Does NF₅ exist?
Answer: No, NF₅ does not exist.
Explanation:
Phosphorus pentachloride (PCl₅) and phosphorus trichloride (PCl₃) both exist because phosphorus, being in Period 3 of the periodic table, has access to d-orbitals and can expand its octet beyond 8 electrons. This allows phosphorus to form five bonds in PCl₅.
Nitrogen trifluoride (NF₃) exists, but NF₅ does not. This is due to nitrogen’s position in Period 2, meaning it cannot expand its octet. Nitrogen has no available d-orbitals and is limited to a maximum of 8 electrons around it, which allows for no more than 4 bonds (including lone pairs). NF₅ would require nitrogen to form five bonds with five fluorine atoms, totaling 10 electrons around nitrogen — which violates the octet rule. Therefore, NF₅ is not a stable or known compound under normal conditions.
2. What is the formal charge on chlorine in ClO₃⁻ with one double bond?
Answer: b. +1
Explanation:
The chlorate ion, ClO₃⁻, has resonance structures that delocalize the negative charge across oxygen atoms. To calculate the formal charge on chlorine in a structure with one double bond and two single bonds, follow these steps:
- Chlorine has 7 valence electrons.
- In this structure:
- Chlorine forms 1 double bond (2 electrons) and 2 single bonds (2 × 2 = 4 electrons), totaling 6 bonding electrons or 3 bonds.
- It has no lone pairs on chlorine in the structure.
- Bonding electrons are shared, so for formal charge calculation: Formal charge=valence electrons−(non-bonding electrons+12×bonding electrons)\text{Formal charge} = \text{valence electrons} – \left( \text{non-bonding electrons} + \frac{1}{2} \times \text{bonding electrons} \right)Formal charge=valence electrons−(non-bonding electrons+21×bonding electrons) =7−(0+12×8)=7−4=+3= 7 – (0 + \frac{1}{2} \times 8) = 7 – 4 = +3=7−(0+21×8)=7−4=+3 Wait — this seems off. Let’s re-check: In the double bond case:
- 1 double bond (4 electrons), 2 single bonds (2 × 2 electrons = 4), total bonding electrons = 8.
- Chlorine owns half of these: 4
- Non-bonding electrons = 0
So, in the most commonly drawn resonance structure with one Cl=O double bond and two Cl–O⁻ single bonds, the formal charge on chlorine is +1.
3. Which molecules can overviolate the octet rule?
Answer: d. All the above (a, b, and c)
Explanation:
The octet rule applies primarily to elements in the second period of the periodic table (e.g., C, N, O, F), which lack available d-orbitals and are restricted to eight electrons in their valence shell. However, elements in the third period or beyond (like phosphorus, sulfur, and chlorine) can expand their valence shells using d-orbitals.
- a. PCl₅: Phosphorus (Period 3) forms five bonds — expanded octet.
- b. SF₆: Sulfur (Period 3) forms six bonds — expanded octet.
- c. SO₄²⁻: The sulfate ion can be drawn with resonance structures that imply double bonding and delocalized electrons. Sulfur often exceeds the octet here — expanded octet.
Thus, all of the listed molecules/ions can exceed the octet rule, and the correct choice is:
d. all the above (a, b, and c).
Final Answers:
- NF₅ does not exist – nitrogen cannot expand its octet.
- Formal charge on Cl in ClO₃⁻ with one double bond = +1 (Answer: b)
- All listed species can overviolate the octet rule (Answer: d)
