Answer
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Step 1:: Determine the total number of valence electrons in the molecule.
Phosphorus (P) is in the fifth group of the periodic table, so it has 5 valence electrons. Chlorine (Cl) is in the seventh group, so it has 7 valence electrons. The Lewis structure for PCl^5 requires us to determine the total number of valence electrons in a phosphorus atom and five chlorine atoms. Total valence electrons = 5 (from P) + 5 × 7 (from Cl) = 5 + 35 = 40 valence electrons
Step 2:: Draw the skeletal structure of the molecule.
Phosphorus is the central atom, and chlorine atoms will surround it. Place one electron from phosphorus and seven electrons from each chlorine atom to form P-Cl bonds.
Step 3:: Distribute the remaining valence electrons as lone pairs on each atom.
Distribute the remaining electrons to complete the octets for chlorine atoms first, then for the phosphorus atom. In this case, there are no remaining electrons, so all 40 valence electrons are used up.
Step 4:: Check the formal charges on each atom.
For each atom, calculate the formal charge as follows: Formal charge = (valence electrons - non-bonding electrons - 1 / 2 × bonding electrons) Phosphorus: (5 - 0 - 1 / 2 × 10) = + 1 Chlorine: (7 - 6 - 1 / 2 × 2) = 0 for each atom Since the phosphorus atom has a formal charge of + 1, we need to move one of the bonding electrons to form a double bond between phosphorus and one of the chlorine atoms. This will give phosphorus a formal charge of zero and maintain the octet rule for all atoms.
Step 5:: Draw the final Lewis structure.
The final Lewis structure for PCl^5 has one phosphorus atom and five chlorine atoms. Phosphorus forms five single bonds with chlorine atoms, and one of these bonds is a double bond. The double bond satisfies the octet rule for phosphorus and results in a formal charge of zero for all atoms.
Final Answer
The Lewis structure for PCl^5 is: P-Cl | | \*/\* | | Cl Cl where the double bond is represented by the asterisks (*).
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