How do I draw the lewis structure for [PF5Cl]- i'm getting confused on this one!! Please explain answer Thanks =))

thanks........but that didn't help me at all =(

To draw the Lewis structure for [PF5Cl]-, follow these steps:

Step 1: Count the total number of valence electrons in the molecule/ion:
- Phosphorus (P) contributes 5 valence electrons
- Fluorine (F) contributes 7 valence electrons x 5 = 35 valence electrons
- Chlorine (Cl) contributes 7 valence electrons
- Negative charge contributes 1 electron

Total valence electrons = 5 + 35 + 7 + 1 = 48

Step 2: Identify the central atom. In this case, it is phosphorus (P).

Step 3: Connect the outer atoms to the central atom with single bonds:
- Phosphorus (P) has 5 single bonds with Fluorine (F)

Step 4: Distribute the remaining electrons to fulfill the octet rule:
- Start by placing the remaining electrons on the outer atoms (F and Cl) to complete their octets.
- Distribute the remaining electrons around the central atom (P) following the octet rule as much as possible.

Step 5: Check if all atoms have achieved an octet:
- Count the electrons around each atom; make sure that they all have an octet (except hydrogen, which can have 2 electrons).

In the case of [PF5Cl]-, the Lewis structure will have the following arrangement:

F
/
F - P - F
\
F

|
Cl-

Note: The negative charge means that there is one extra electron, which will be placed on the Cl atom in this case.

To draw the Lewis structure for [PF5Cl]-, follow these steps:

Step 1: Count the total number of valence electrons.

The phosphorus atom (P) contributes 5 valence electrons, and each of the fluorine atoms (F) contributes 7 valence electrons (since fluorine is in Group 7A or 17), and the chlorine atom (Cl) contributes 7 valence electrons. The negative charge (-) also adds one additional valence electron.

Total valence electrons = 5 + (5 x 7) + 7 + 1 = 48

Step 2: Identify the central atom.

In this case, the phosphorus atom (P) is the central atom, as it is the least electronegative element.

Step 3: Connect the atoms with single bonds.

Connect each fluorine atom (F) and the chlorine atom (Cl) to the central phosphorus atom (P) using single bonds. This accounts for 6 electrons (2 for each bond).

Step 4: Distribute the remaining electrons.

Place the remaining electrons around the atoms in pairs to complete their octets, except for the central atom since it can have more than 8 electrons due to its position in the periodic table.

Start by placing electron pairs on the outer atoms (F and Cl) until all electrons are used or every atom has at least an octet.

Place the electrons around the remaining atoms as lone pairs.

Step 5: Check if the central atom has exceeded the octet rule.

In this case, phosphorus (P) can expand its octet and accommodate more than 8 electrons since it is in the third period or beyond.

Step 6: Determine the formal charges.

Calculate the formal charges for each atom by considering the difference between the number of valence electrons and the assigned electrons (lone pairs + half of the shared electrons).

Step 7: Verify that the total number of valence electrons equals the sum of the formal charges and the overall charge.

If the total number of valence electrons equals the sum of the formal charges and the overall charge (if present), then your structure is correct.

Remember to consider the formal charge on the molecule as well.

I hope this clears up the confusion in drawing the Lewis structure for [PF5Cl]-. If you need further assistance or clarification, feel free to ask.

http://en.wikipedia.org/wiki/Octahedral_molecular_geometry