Use formal charges to predict the most plausible Lewis structure for thiosulfate ion, S2 O3^2-.

Use formal charges to predict the most plausible Lewis...

Key Concepts

Lewis Structures

A Lewis structure is a diagram that represents the bonding between atoms of a molecule and the lone pairs of electrons that may exist. It is used to visualize how atoms are connected, how electrons are shared among them, and to provide insight into the molecule’s overall geometry. In constructing these structures, one must account for all the valence electrons and arrange them in a way that reflects likely bonds and lone pairs while following established bonding rules.

Valence Electrons

Valence electrons are the outermost electrons of an atom available for bonding. When drawing a Lewis structure, counting the total number of valence electrons is a critical first step, including any adjustments for ions by adding electrons for negative charges or subtracting electrons for positive charges. This count ensures that no electron is unaccounted for in the structure and helps in determining the connectivity and distribution of electrons among atoms.

Formal Charge

Formal charge is an assigned charge on an atom in a Lewis structure, calculated using the formula: formal charge = (valence electrons of the atom) - (nonbonding electrons + one-half the bonding electrons). The most favorable Lewis structure typically minimizes the formal charges on atoms as much as possible, thereby contributing to the overall stability of the ion or molecule. It serves as a guide for evaluating different possible structures.

Octet Rule

The octet rule is a chemical rule of thumb that states atoms tend to form bonds in such a way that they have eight electrons in their valence shell, thus achieving a stable electron configuration similar to the noble gases. While most second-period elements strictly follow the octet rule, elements in the third period and beyond can have expanded octets. In forming Lewis structures, this rule provides a guideline for the distribution of electrons around each atom and helps verify that the most likely structures are considered.

Resonance Structures

Resonance structures are different valid Lewis structures that represent the same molecule or ion by showing the possible arrangements of electron pairs around the atoms. These structures illustrate the delocalization of electrons within a molecule, contributing to its overall stability. In cases where multiple arrangements of electrons yield comparable formal charges, resonance structures help in depicting the actual electron distribution, which is a hybrid of the individual forms.

Transcript

00:01 We are using formal charges to predict the most plausible loo structure for thiosulfate, which is s2o3 with a minus 2 charge.

00:10 Now, the previous question number 65 already had you draw out the resonance structures.

00:15 So it's possible that a large part of this you've already done.

00:17 But if not, we're going to go ahead and draw the lewis structure and then use formal charge to choose which one would be the most stable form of it.

00:25 Ah, where to go? come back.

00:29 There it is.

00:30 All right.

00:31 So to start with, we need to know how many very very.

00:33 Valence electrons we're working with.

00:34 I've already drawn out the columns with periodic table.

00:38 Sulfur would have six valence electrons, and there's two of those.

00:46 Oxygen would also have six, and there's three of those.

00:51 And then the plus two charge means that there's two electrons that were added to the system on top of that.

00:56 So your total is going to be 32 electrons.

01:01 We're putting in our drawing.

01:04 All right.

01:05 Now drawing this thing is the tough part.

01:09 All right.

01:10 This has kind of a weird structure.

01:14 It's really similar to sulfate where one of the oxygens is replaced by a sulfur.

01:20 Looks like that.

01:24 Now we've got how many electrons so far? 1, 2, 3, 4, 5, 6, 7, 8 that we've used so far, 2 per bond.

01:31 So we have 24 left to go.

01:35 We'll fill in our outer atoms before we fill in the inner one.

01:40 So each of those needs, six more to get up to their desired eight.

01:50 Okay, six times four is 24.

01:52 We are now out of electrons.

01:56 Now, based on just the lewis structure rules, this would be complete.

01:59 The inner sulfur has a full octet.

02:01 The outer ones all have octets.

02:03 We're fine.

02:04 But this is asking us to use formal charges to predict the most plausible form.

02:10 Now, formal charges take a look at how many electrons each atom has started with and decides if it still has possession of the form.

02:18 Those, if it's sharing equally, if it's trying to steal from neighbors.

02:22 And the most stable situation is when your formal charges all come out as zero.

02:28 Very quickly, your formal charge equation is for a single atom, it's how many valence electrons that atom has based on the periodic table, minus half of its bonding electrons, and then all of its non -bonding electrons.

02:47 Bonding electrons are the ones tied up in bonds.

02:50 Non -bondings are the individual dots.

02:52 If you have a hard time remembering that, i always just use formal charge equals the valence minus how many sticks are attached to that atom, like how many bonds, minus how many dots are on that atom, the individual electrons.

03:10 Okay, so for the oxygens on the outside, okay, they each, an oxygen starts with six valence, minus, there's one bond on that oxygen, one stick, so that's minus one, minus six dots, six separate lone electrons.

03:30 So that's a formal charge of minus one.

03:33 So all of these oxygens have a minus one.

03:39 Then we look at the sulfur out here.

03:42 Sulfur starts with six valence, minus one bond and six dots.

03:50 Again, equals minus one.

03:52 Zero's would be better.

03:54 We're working with what we got.

03:57 Now, lastly, that sulfur in the middle starts with six minus four bonds and no dots, which is plus two.

04:11 Having four minus ones and then a plus two really isn't ideal, it'd be better if we could get more zeros in place...