Write Lewis formulas for the resonance forms of the thiosulfate ion, S2 O3^2- . Indicate formal

Write Lewis formulas for the resonance forms of the...

Key Concepts

Resonance Structures

Resonance structures are multiple valid Lewis structures for a molecule or ion that cannot be adequately represented by a single bonding pattern. These structures differ only in the arrangement of electrons, not in the arrangement of atoms, and collectively describe the delocalization of electrons within the ion. The true structure is a hybrid of these resonance forms.

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They are especially useful for understanding the geometrical arrangement and electron sharing in polyatomic ions. In drawing these structures, one shows each atom, the bonds between them, and assigns electrons to satisfy the octet rule where applicable.

Formal Charge

Formal charge is a bookkeeping tool used in Lewis structures to indicate the hypothetical charge an atom would have if all bonding electrons were shared equally between the bonded atoms. It is calculated by subtracting the number of electrons assigned to an atom in the structure from its valence electrons. Tracking formal charges helps identify the most stable resonance forms and the distribution of charge in polyatomic ions.

Delocalized Bonding

Delocalized bonding refers to the phenomenon in certain molecules or ions where electrons are spread over several adjacent atoms, rather than localized between a pair of atoms. This results in bond orders that are averaged over several bonds, contributing to extra stability. In the context of polyatomic ions, like the thiosulfate ion, delocalization is reflected in equivalent or near-equivalent bonding environments found in the resonance structures.

Transcript

00:01 In this video, we're going to draw the lewis structures for the s2 -2 -2 -minus or thiosulfate ion, and we're told it has resonance.

00:09 So let's see.

00:10 So we have two sulfurs, and each of them is going to have six valence electrons, plus three oxygens, and each of those is going to have six valence electrons.

00:17 And then the negative two indicates that we've gained two electrons.

00:20 So the available electrons for our lewis structure are going to be 12 plus 18 is 32.

00:29 Let's draw the skeleton structure.

00:30 So we have sulfur is going to be bonded to sulfur, and then we'll have, let's see, an oxygen, an oxygen like that, and an oxygen like that, probably.

00:56 Or more likely, actually, because the sulfur and the oxygen have similar bonding patterns if we put the third oxygen here and then a bonded sulfur like that, treating it like an oxygen.

01:08 Okay.

01:09 And then we want 32 electrons.

01:18 So one way we could do that is if we leave it like this, but then everything's going to have a very negative formal charge.

01:26 So we can also just add in two double bonds.

01:31 So we can do that to two oxygens at first.

01:35 Let's see.

01:36 We do that.

01:37 Let's do that one.

01:39 Like that and like that...