Final answer:
The Lewis structure for the ICl₂⁻ ion consists of a central Iodine atom bonded to two Chlorine atoms, each with three lone pairs of electrons, and the Iodine with two lone pairs and an extra electron to account for the -1 charge of the ion. The formal charge is -1 for Iodine and 0 for each Chlorine.
Step-by-step explanation:
To draw the Lewis structure for the ICl₂⁻ ion, start by arranging the Iodine (I) atom in the center with the two Chlorine (Cl) atoms on either side. Connect each Cl to the I atom with a single line to represent a covalent bond. According to the octet rule, each Cl atom will have three lone pairs, as they each need eight electrons to complete their octet (which is already achieved with one shared pair and six non-bonding electrons).
The Iodine atom, on the other hand, will have two lone pairs along with the two single covalent bonds with the chlorine atoms. Since the molecule carries a -1 charge, one more electron is added to the Iodine, which now has a formal charge of -1. The Lewis structure should be enclosed in brackets with the charge (-1) written outside the bracket to denote the overall charge of the ion. Here's how you calculate the formal charges:
For Iodine (I): 7 valence electrons (for neutral atom) - 8 electrons (6 nonbonding + 2 bonding) = -1
For each Chlorine (Cl): 7 valence electrons (for neutral atom) - 7 electrons (6 nonbonding + 1 bonding) = 0
The formal charge for each Chlorine is zero, so no need to mark them. The sum of the formal charges of all the atoms equals -1, which matches the charge of the ion.