To draw the Lewis structure for ClF3 (chlorine trifluoride), we start by calculating the total number of valence electrons. Chlorine (Cl) has 7 valence electrons, and each fluorine (F) has 7 valence electrons. Since there are three fluorine atoms, the total number of valence electrons is:
7 (from Cl) + 3 × 7 (from F) = 28 valence electrons.
Next, we place the chlorine atom in the center since it’s the least electronegative, and arrange the three fluorine atoms around it. We connect each fluorine atom to the chlorine with a single bond, using 6 of the total 28 electrons (2 electrons per bond). This leaves us with 22 electrons.
Now, we need to satisfy the octet rule for the fluorine atoms. Each fluorine will have 3 lone pairs (6 electrons) to complete their octet. Assigning these lone pairs uses up 18 electrons (3 fluorine × 6 electrons). After this, we have:
28 – 6 (bonds) – 18 (lone pairs) = 4 electrons remaining, which will be placed as 2 lone pairs on the chlorine atom itself.
The Lewis structure can be illustrated as follows:
F F | | Cl ||
In the Lewis structure, chlorine has 2 lone pairs and forms 3 single bonds with the fluorine atoms.
Now, let’s determine the electron geometry. With chlorine having 3 bonded pairs and 2 lone pairs, we use the VSEPR theory. The arrangement of these electron pairs is generally trigonal bipyramidal. However, considering the lone pairs, the molecular geometry is T-shaped.
Regarding the polarity of the molecule, since chlorine is less electronegative than fluorine, each Cl-F bond is polar. Given that the molecular shape is not symmetrical due to the T-shaped arrangement, this leads to an unequal distribution of electron density which confirms that ClF3 is a polar molecule.
In summary:
- Lewis Structure: Cl with 3 F atoms and 2 lone pairs.
- Electron Geometry: Trigonal bipyramidal; Molecular Geometry: T-shaped.
- Non-bonding Domains: 2 on the central atom (Cl).
- Polarity: Polar molecule.