Homework 2E.5

(Polar molecules, Non-polar molecules, etc.)

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WYacob_2C
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Joined: Sat Jul 20, 2019 12:16 am

Homework 2E.5

Postby WYacob_2C » Thu Nov 14, 2019 11:45 am

Can someone help me with this problem? It states, " (a) What is the shape of a ClO2+ ion? (b) What is the expected OClO bond angle?"

FDeCastro_1B
Posts: 54
Joined: Thu Jul 25, 2019 12:16 am

Re: Homework 2E.5

Postby FDeCastro_1B » Thu Nov 14, 2019 12:48 pm

To find the shape you'll have to start with the Lewis structure of the ion. From there, you'll find the amount of bonded pairs there are around the central atom which will give you the shape of the atom (trigonal planar, tetrahedral, etc.).

To find the expected bond angle you'll have to start with the Lewis structure. From the Lewis structure you'll be able to determine the shape of the compound. From the shape, you'll be able to determine the bond angles.

Brian_Ho_2B
Posts: 221
Joined: Fri Aug 09, 2019 12:16 am

Re: Homework 2E.5

Postby Brian_Ho_2B » Thu Nov 14, 2019 12:51 pm

WYacob_1F wrote:Can someone help me with this problem? It states, " (a) What is the shape of a ClO2+ ion? (b) What is the expected OClO bond angle?"

To start, we will draw the lewis structure of this molecule. The total number of valence electrons is 18, and since Cl isn't as electronegative as the oxygen atoms, we can assign the positive formal charge to the Cl atom. Placing Cl in the center, we have 2 double bonds connecting the oxygen atoms to Cl and Cl is left with a lone pair. Since there are three regions of electron density (2 double bonds and one lone pair), we would expect the shape to be trigonal planar; however, one of those regions is just a lone pair, so we do not include it in naming the shape (the shape is still influenced by the lone pair!). The shape looks trigonal planar except one region of electron density is invisible because there technically isn't a bond there, so the shape is considered "angular", with two oxygen atoms and a Cl in the middle forming an angle less than 180 degrees. The expected angle created by the O-CL-O bond should be less than 120 degrees because the lone pair has stronger repulsion power, so it pushes the two oxygens slightly closer to each other. It is specifically less than 120 degrees because the three regions of electron density form a trigonal planar shape (except it is actually just "angular" due to the lone pair).


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