ICl₄⁻ Lewis Structure

Tetrachloroiodate ion · An interhalogen anion formed when iodine trichloride meets chloride.

• 36 valence e−
• square planar
• 90° bond angle
• sp³d²
• 1− charge

ICl₄⁻ molecular geometry

Calculated properties

I Iodine 7 valence e− Cl Chlorine 4 atoms, 7 valence e− each

How to draw the ICl₄⁻ Lewis structure

Step through it. The bar shows how much of the 36-electron budget is spent at each stage - it never grows, which is the whole constraint.

The reasoning at each step

1. Count the valence electrons

   Add up the outer-shell electrons every atom brings, then account for the charge. A 1- ion has gained one electron, so it is added to the budget. Skip this and every count that follows will be wrong.

   7 (I) + 4 x 7 (Cl) + 1 (1- charge) = 36 valence electrons

2. Choose the central atom

   I is the least electronegative atom here, so it takes the centre. Hydrogen is never central: it can only form one bond.

3. Join everything with single bonds first

   Every connection starts as one shared pair. Draw them all before worrying about double bonds - the arithmetic in the next step is what tells you where the double bonds have to go.

   4 bonds x 2 = 8 electrons used, 28 left

4. Work out how far short you are

   Add up what every atom still needs to fill its shell - eight electrons for most atoms, two for hydrogen - minus what its single bonds already give it. Compare that with the electrons you have left. The difference decides everything that follows.

   needs 24, has 28, 4 to spare

5. Park the spare electrons on the central atom

   There are 4 electrons more than the outer atoms can use. They become 2 extra lone pairs on I, which is in period 5 and can hold more than eight electrons. A period-2 atom could not do this.

   12 electrons around I

6. Fill in the lone pairs

   Every electron not in a bond sits as a lone pair on an atom. Each atom takes exactly what it needs to finish its shell - there is no choice left at this point.

   I: 2, Cl1: 3, Cl2: 3, Cl3: 3, Cl4: 3 (28 electrons)

7. Check the central atom

   I ends up with 12 electrons, an expanded octet. Only period 3 and below can do this.

   I: 8 bonding + 4 nonbonding = 12

8. Check the formal charges

   Formal charge is valence electrons, minus lone-pair electrons, minus the number of bonds. They are not real charges, and they are not the ion's charge either - but they must add up to it. Adding these up gives 1-, which is the check that the structure is complete.

   I: 7 - 4 - 4 = -1

9. Work out the shape

   Count the electron domains on I: 4 bonded groups plus 2 lone pairs. A double or triple bond still counts as one domain, because it points in one direction. That gives octahedral electron geometry; ignore the lone pairs and the atoms themselves sit in a square planar arrangement. The two lone pairs sit opposite each other, above and below the plane, so the four bonds stay at a clean 90 degrees.

   steric number 6 -> sp3d2 -> square planar, bond angle 90

10. Decide whether it is polar

    The bond dipoles and lone pairs are arranged symmetrically, so the vectors cancel and the molecule has no net dipole.

    dipoles cancel -> nonpolar

Where the 1− charge sits

The I carries a 1− formal charge, and that is the ion's 1− charge. The charge belongs to the ion as a whole, which is what the brackets around the drawing mean.

People also ask whether ICl₄⁻ is polar. Strictly it is neither: polar and nonpolar sort neutral molecules by whether their bond dipoles cancel. A 1− ion is pulled towards any nearby charge whatever its shape, which is why its salts dissolve in water regardless of geometry. The table below is still worth reading - it shows how unevenly each bond holds its electrons, which is what decides whether the charge ends up spread or concentrated.

Common questions

How many valence electrons does ICl4- have?

ICl4- has 36 valence electrons. The atoms bring 35 between them, and the 1− charge adds one. 8 of them are in bonds and 28 sit in lone pairs.

Why does ICl4- have a 1− charge?

The I carries a 1− formal charge, and that is the ion's 1− charge. The charge belongs to the ion as a whole, which is what the brackets around the drawing mean. The brackets around the drawing are part of the answer: they say the charge belongs to the whole ion rather than to whichever atom happens to be nearest.

What is the molecular geometry of ICl4-?

ICl4- is a square planar ion. The central I has 4 bonded groups and two lone pairs, a steric number of 6, which gives octahedral electron geometry and a square planar ion.

What is the bond angle in ICl4-?

The bond angle in ICl4- is 90 degrees. The two lone pairs sit opposite each other, above and below the plane, so the four bonds stay at a clean 90 degrees.

What is the hybridization of ICl4-?

The central I in ICl4- is sp³d² hybridized. Steric number 6 means 6 orbitals have to be mixed, which is exactly what sp³d² gives you.

Is ICl4- polar or nonpolar?

Neither, strictly. Polar and nonpolar describe neutral molecules, by asking whether their bond dipoles cancel out. ICl4- carries a net 1− charge, so it is pulled towards any nearby charge whatever its shape - which is why ionic compounds dissolve in water regardless of geometry. The question the shape does answer is whether the charge is spread evenly: it is, because the square planar arrangement puts every identical bond at the same angle to the others.

Related structures

• I₃⁻ also a 1− ion
• IO₃⁻ also a 1− ion
• ICl₂⁻ also a 1− ion
• XeF₄ also square planar
• IF₅ also built around I
• NO₃⁻ also a 1− ion