Why isn't ClF3 trigonal planar?

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ClF3 exhibits a trigonal bipyramidal arrangement with a T-shaped molecular geometry due to the presence of non-bonding electron pairs that exert greater steric repulsion compared to bonding pairs. In VSEPR theory, non-bonding pairs are considered sterically more demanding, meaning they require more space and occupy a larger volume, which influences molecular shape. The discussion highlights that in the first case, non-bonding pairs are positioned with only two neighbors at 90-degree angles, while in the second scenario, they would have three neighbors at 120 degrees, leading to increased repulsion. This repulsion is crucial in determining the actual geometry of ClF3, reinforcing why it does not adopt the second case's trigonal planar shape. The term "sterically more demanding" refers to the need for non-bonding pairs to have more space due to their larger volume compared to bonding pairs.
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I know ClF3 has triagonal bi-pyramidal arrangement and T-shape molecular geometry. (as shown in first diag.). However, it can also be 2nd case. In this one, the shape will be triagonal planar. Also, electrons will be farthest. So why isn't ClF3 like second case?
 

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In VSEPR theory, non-bonding electron pairs are sterically more demanding than bound electron pairs. In the first formula, each non-bonding electron pair has only two neighbours in a 90 degree separation while in the 2nd one, each has 3. Whether the distance between the two non-bonding pairs is 120 or 180 degrees will be rather irrelevant compared to the next neighbour repulsion.
 
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I get what you mean. Thank you.
But would you be kind enough to explain what
DrDu said:
'sterically more demanding than bound electron pairs'
means?
[I did look up the meaning of steric but the meaning still isn't clear]
Thanks.
 
Need more space for themselves, occupy larger volume.
 

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