Iodine forms a series of fluorides (listed here). Write...

Key Concepts

Lewis Structures

Lewis structures are diagrams that represent the valence electrons of atoms within a molecule, showing how these electrons are arranged as bonding pairs and lone pairs. They offer a visual guide for predicting bonding interactions, molecular geometry, and reactivity by allocating electrons to each atom in accordance with the molecule’s overall electron count.

Formal Charge

Formal charge is the computed charge on an atom in a molecule derived from the difference between the valence electrons in the free atom and the electrons assigned to it in the molecule's Lewis structure. Calculating formal charges helps in identifying the most stable or likely structure among several possibilities, ensuring the electrons are distributed according to electronegativity and bonding considerations.

Octet Rule and Expanded Octet

The octet rule is a principle that atoms tend to form bonds until they are surrounded by eight valence electrons, mimicking the electron configuration of noble gases. However, elements that are in the third period or beyond, such as iodine, can have an expanded octet, meaning they are able to accommodate more than eight electrons in their valence shell due to available d-orbitals.

Molecular Geometry and VSEPR Theory

VSEPR (Valence Shell Electron Pair Repulsion) theory uses the concept that electron pairs repel each other to determine the three-dimensional shape of a molecule. This theory helps predict molecular geometry by considering both bonding pairs and lone pairs around a central atom, which is critical for understanding the structural arrangement in compounds such as the series of iodine fluorides.

Transcript

0:00 Good day.

00:01 In this question we are told that iodine reacts with fluorine to produce several molecules.

00:12 Now we are told that we need to write down the lower structures for each of these molecules and also find the formal charge for iodine in each of those molecules.

00:25 So we'll begin with a iodine fluoride.

00:34 So to begin with for the lewis structure, we need to calculate the total number of valence electrons by adding the individual valence electrons.

00:42 So for iodine, it is 7, fluorine is 7 as well.

00:47 That gives us a total of 14 valence electrons.

00:51 So we write our iodine and our fluorine.

00:55 Then the following step is to put our bonding electrons, bonding.

01:01 Pair, that's two, and then we then fill the outer atom octates.

01:15 And counting them, we've got 2468, 2468, and that's total of 14 electrons and both iodine and fluorine have stable octates.

01:28 Then now to calculate the formal charge, formal charge is equal to the number of valence electrons minus the non -bonding valence electrons minus the bonding electrons divided by two.

01:39 So we want to calculate the formal charge for iodine.

01:43 Okay, so for iodine, so the number of valence electrons is 7 minus the number of non -bonding valence electrons, which is 2 ,4, 6, minus the bonding electrons, which is 2 over, and the answer is 0.

02:12 Moving on to b, which is iodine trifluoride.

02:23 So to calculate the valence electrons, we've got 7 plus 7 times 3, there are 3 fluorine atoms, and that gives us a total of 28 valence electrons.

02:41 So in the center, we write down the least electronegative atom, which is iodine.

02:47 And then the surrounding atoms are fluorine atoms.

02:54 Then the following step is to put the bonding electrons and then we fill the octates of the surrounding fluorine atoms.

03:14 Now when we count these, we find out that we now have a total of 24 and that leaves us with four more electrons.

03:23 Now iodine is in period 5, so it can hold more than 8 electrons.

03:27 And this is an exception to the octet rule.

03:29 So we can just put the two more electrons on iodine.

03:34 Then now to calculate the formal charge for iodine, it will be the total number of valence electrons, which is seven, minus the bonding electrons, which is equal to the non -bonding electrons, which is 2 -4, minus the number of electrons that are bonding divided by 2, which is 2, 4, 6.

04:05 So that is 6 divided by 2.

04:08 So that's 7 minus 4 minus 3, which is equal to 0.

04:12 So the formal charge for iodine is 0.

04:16 And then for c, this is iodine pentafluoride...

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