Lewis Structure of Polyatomic Molecules with Central Element Halogen

General Chemistry • Chemical Bonds

Written by STEM Calculators Team Published December 6, 2025 Updated April 21, 2026

Polyatomic Ions with a Central Halogen (Cl, Br, I)

Choose the central halogen and the number of oxygens to identify the oxoanion formula, traditional name, oxidation state, valence-electron count, resonance pattern, and a readable Lewis-structure schematic.

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Overview of halogen oxoanions XO_n⁻

Chlorine, bromine, and iodine form a characteristic family of oxoanions of the type XO_n⁻, where X is the halogen and n is the number of oxygens (from 1 to 4). For example, the chlorine series is: ClO⁻, ClO₂⁻, ClO₃⁻, and ClO₄⁻.

Ion	Traditional name	Oxidation number of X
ClO⁻	hypochlorite	+1
ClO₂⁻	chlorite	+3
ClO₃⁻	chlorate	+5
ClO₄⁻	perchlorate	+7

The same pattern applies to bromine (BrO⁻, BrO₂⁻, BrO₃⁻, BrO₄⁻) and iodine (IO⁻, IO₂⁻, IO₃⁻, IO₄⁻), with names such as hypobromite, bromite, bromate, perbromate, and the analogous iodite / iodate / periodate ions.

Oxidation numbers in XO_n⁻

In these oxoanions, oxygen is assigned an oxidation number of −2 and the overall charge is −1. Let the oxidation number of the halogen X be x. For a general ion XO_n⁻:

Step 1. Write the oxidation-number equation.

\[ \begin{aligned} x + n\cdot(-2) &= -1 \end{aligned} \]

Step 2. Solve algebraically for \(x\).

\[ \begin{aligned} x - 2n &= -1 \\ x &= -1 + 2n \\ x &= 2n - 1 \end{aligned} \]

Result. The oxidation number of the halogen in XO_n⁻ is \(x = 2n - 1\).

Substituting concrete values of n gives the familiar sequence of oxidation numbers: +1 (n = 1), +3 (n = 2), +5 (n = 3), and +7 (n = 4). This is why the series of names progresses from “hypo–ite” (lowest oxidation state) through “–ite” and “–ate” to “per–ate” (highest oxidation state).

Name pattern for ClO_n⁻, BrO_n⁻, IO_n⁻

The traditional names follow a simple prefix–suffix pattern based on the number of oxygens:

n = 1: “hypo–ite” → ClO⁻ is hypochlorite, BrO⁻ is hypobromite.
n = 2: “–ite” → ClO₂⁻ is chlorite, IO₂⁻ is iodite.
n = 3: “–ate” → ClO₃⁻ is chlorate, BrO₃⁻ is bromate, IO₃⁻ is iodate.
n = 4: “per–ate” → ClO₄⁻ is perchlorate, IO₄⁻ is periodate.

The calculator mirrors this logic: given the halogen symbol and the number of oxygens, it constructs the correct formula, applies the naming pattern, computes the oxidation number using \(x = 2n - 1\), and displays a representative Lewis-structure schematic. For Br and I, the Lewis structures are analogous to those of the corresponding chlorine oxoanions, differing only in the identity of the central atom.

Frequently Asked Questions

How do you count total valence electrons for an oxyhalogen ion?

Add the valence electrons from each atom and then adjust for the charge. A negative charge adds electrons and a positive charge removes electrons (total valence electrons = sum of atomic valence electrons - ionic charge).

Why do many halogen-oxygen polyatomic ions show resonance in Lewis structures?

More than one valid Lewis structure can be drawn by placing the multiple bond and charge on different oxygen atoms while keeping the same total electron count. These equivalent arrangements are resonance forms and represent delocalized bonding in the real species.

What is the formal charge formula used to evaluate a Lewis structure?

Formal charge can be written as FC = V - (N + B/2), where V is the atom's valence electrons, N is nonbonding electrons, and B is bonding electrons. Preferred structures usually minimize formal-charge magnitude and place negative charge on oxygen.

Why can a central halogen appear to have an expanded octet in some Lewis structures?

For halogens in period 3 or below, common general-chemistry Lewis descriptions may place more than 8 electrons around the central atom to reduce formal charges in oxyhalogen species. This is a simplified model used to represent bonding and charge distribution.

Polyatomic Ions with a Central Halogen (Cl, Br, I)

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