Draw Lewis structures for the AsCl4^+ and AsCl6^- ions. What type of reaction (acid-base, oxidat

Draw Lewis structures for the AsCl4^+ and AsCl6^- ions. What...

Key Concepts

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms in a molecule as well as the lone pairs of electrons that may exist. They are a basic tool used in understanding molecular geometry, reactivity, and the distribution of electrons, which is essential for drawing accurate representations of polyatomic ions and understanding how atoms interact in a molecule.

Expanded Octet/Hypervalency

For atoms in period 3 and beyond, such as arsenic, the ability to have more than eight electrons in their valence shell is known as hypervalency or an expanded octet. This concept is key to understanding the bonding in compounds where the central atom forms more than four bonds, such as in certain chlorides, because it explains how these species can accommodate extra electron pairs.

Lewis Acid-Base Reactions and Autoionization

Lewis acid-base reactions involve the donation of an electron pair from a Lewis base to a Lewis acid. In some systems, a molecule can simultaneously act as both a Lewis acid and a Lewis base, leading to autoionization where one molecule gives up an electron pair to another identical molecule. This self-ionization leads to the formation of a positive ion and a negative ion without changing the oxidation states of the central atom.

Reaction Classification (Acid-Base vs Redox)

In chemical reactions, it is important to distinguish between reaction types. An acid-base reaction is characterized by the transfer of electron pairs (or protons in the Brønsted–Lowry framework) without a change in oxidation states, whereas oxidation-reduction (redox) reactions involve a change in oxidation numbers due to electron transfer. Classifying a reaction correctly requires examining if oxidation numbers change; if they remain the same, the reaction might be better described as an acid-base or ionization process rather than a redox reaction.

Transcript

00:01 So in this video we're going to work through question 88 from chapter 20, which says to draw lewis structures for the ascl4 plus and ascl6 minus ions.

00:11 What type of reaction? acid base, oxidation reduction, or the like, is the following.

00:16 So our reaction is to asdl5 reacting to form ascl4, ascl6, the ionic compound of our positive ascl4 ion and our negative ascl6 ion.

00:31 Let's go ahead and start drawing lewis structures.

00:34 So for ascl 4 minus, we know that our arsenic is going to bring five electrons, and each of our chlorine atoms will bring seven electrons, and then because we have a negative one charge, we have one extra electrons.

00:46 So total, we have 32 electrons.

00:49 So to draw our lewis structure, we can just go ahead and position our four chlorine atoms around our central arsenic atom, and then fill in the octet of each of our chlorine atoms, and that actually uses up all 32 of our electrons.

01:07 So, and then we indicate the positive charge by putting the whole thing in brackets and putting a positive in the right, top right corner.

01:17 And then i want to go ahead and assign oxidation states in formal charges while we're here.

01:23 Now, remember that when we're assigning formal charges, electrons that make up bonds are shared by that two atoms in the bond, but in oxidation states, the more electronegative atom of a bond will take all of the electrons in that bond.

01:37 So let's start with oxidation states.

01:40 So since chlorine is more electronegative than arsenic, it has its six lone pair electrons plus two that make up the bond.

01:49 So we have seven valence electrons minus six lone pair electrons minus two electrons in my bond.

01:56 And that makes an oxidation state of minus 1.

02:01 Now for our arsenic, we have five valence electrons minus no lone pair electrons and minus no bonding electrons because all of our bonds are with chlorine, which is more electronegative.

02:14 So we have a plus five oxidation state.

02:17 Now for formal charges, we have for chlorine, seven valence electrons minus six lone pair electrons minus half of the bonding electrons, so minus one.

02:28 And that gives us a formal charge of zero.

02:30 Now for arsenic, we have five valence electrons minus half of our bonding electrons minus four.

02:39 So that gives us a formal charge of plus one.

02:44 So let's do the same thing for ascl6 minus.

02:48 So we know that our arsenic brings five electrons, and that each of our chlorides is going to bring seven electrons.

02:56 And then we have one fewer electron, one more electron, one extra electron, because we have a minus one charge.

03:04 So we have five plus six times seven plus one that gives us a total of 48 electrons.

03:09 And we can use all those electrons by just positioning all six of our chlorine atoms around our arsenic atom and then single bonding them and filling in the octet on the chlorine atoms.

03:20 And that should use up all 48 of your electrons.

03:22 And then the whole thing has a minus one charge.

03:25 It in brackets and put a minus 1 to the top right.

03:28 So assigning oxidation states, again, chlorine has seven valence electrons minus six lone pair electrons minus two bonding electrons that gives us an oxidation state of minus one.

03:40 And again, arsenic had five valence electrons and has no bonding electrons because it's less an electronegative than chlorine and no lone pair electrons.

03:52 So the oxidation state on arsenic is plus five...

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