Shape vs geometry MCAT question

[br2pi5]

I thought MCAT questions about geometry and shape were pretty direct. geometry meant only looking at number of bonds while shape meant looking at both bonds and lone pairs. however, I came across this question (below). In the explanation of choosing between SiCl4 and XeF4 they mention the lone pairs too. I guess I wonder what I should focus on when asked geometry or shape: the bonds, the lone pairs, the electron groups?

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[JustinM88]

Kind of need to draw out each one's structure and take an overall look at what's going on.

Si is right below C in the periodic table of elements so it's going to be quite similar to that C in methane (ie. it has the same # of valence electrons) so when it's the central atom it will have the same number of lone pairs as C (ie. none in this case).

Xe's extra lone pairs repel the bonds (google its structure and you'll see what I mean) so the lone pairs do affect were the bonds are.

 

[aldol16]

You should focus on visualizing the 3D orientation of a molecule. Forget the shape vs. geometry distinction - it doesn't matter. It's invented to help us teach an often difficult concept to students. You just have to understand the concept. For instance, what does water look like? Well, it's only got three atoms and without taking into account electrons, you might think that it would want to be in a straight line. But you know that oxygen's two lone pairs distort this and force the molecule into a bent form. If you can understand that, you should be fine.

 

[NextStepTutor_1]

Hi @br2pi5 -

I definitely agree w/ @aldol16 that the specific terminology is less important than being able to understand and explain in your own words the factors that give a molecule a certain shape.

An analogy that I find useful (but cannot take credit for myself -- it's from an excellent chemistry professor I had way back when) is to think of lone pairs and bonding orbitals to basically be "electron buckets" that want to be as far away from each other as possible. So, if you have 4 "electron buckets", their baseline orientation will be tetrahedral because that is the orientation that maximizes the separation of these electron buckets. If you have 3 electron buckets, they will be arranged in a planar manner, at 120°. If you have only 2 electron buckets, they will be 180° apart. Now let's work through what happens if you have various combinations of lone pairs and bonding orbitals.

If you start with 4 "electron buckets" and all 4 are used to form a bond, you'll have a classic tetrahedral geometry like that of CH4. If only 3 bonds are formed, and 1 lone pair remains, you're essentially just filling in 3 of the 4 available slots with bonds. These three bonds will be arranged in a pyramidal shape (NH3 is an example). If you have 2 bonds and 2 lone pairs, you're only filling in 2 of the slots, so you will get a bent shape (H2O is the classic example of this). The bond angles in the tetrahedral orientation are ~109.5°, but in the pyramidal and bent orientations, they are further compressed to about ~107° and ~105° due to the lone pairs (but that's not a highly important factoid for the MCAT).

If you start with 3 "electron buckets" and use all 3 to form a bond, you'll have a trigonal planar shape, like AlBr3. (Note that this can happen with a molecule like AlBr3 that has 3 single bonds and no lone pairs, or in a molecule like formaldehyde (H2CO) that has two single bonds and a double bond -- in this metaphor, an "electron bucket" can correspond to a double or triple bond). However, if you only use 2 of these "electron buckets" and have a lone pair left over, like in SO2, you'll have a bent geometry, although this time the bond angle will be slightly less than 120°, because you're starting with a trigonal planar orientation with a baseline angle of 120°.

Finally, if you start with only 2 "electron buckets," you're bound to have a linear geometry, because they're located at 180° from each other. This can happen if both "buckets" are used to form bonds, as in CO2, or if a lone pair remains, as in the negatively charged carbon in an acetylide anion.

This way of looking at things is somewhat simplified and doesn't build connections with orbital hybridization, but I've found it to be a simple and spatially intuitive way of visualizing the shape of a molecule without getting bogged down in terminology that may be more of a hindrance than a help. Hope this clarifies things, and best of luck!

 

[NextStepTutor_1]

[accidentally posted a second reply, please disregard]

 

[br2pi5]

Hi @br2pi5 -

I definitely agree w/ @aldol16 that the specific terminology is less important than being able to understand and explain in your own words the factors that give a molecule a certain shape.

An analogy that I find useful (but cannot take credit for myself -- it's from an excellent chemistry professor I had way back when) is to think of lone pairs and bonding orbitals to basically be "electron buckets" that want to be as far away from each other as possible. So, if you have 4 "electron buckets", their baseline orientation will be tetrahedral because that is the orientation that maximizes the separation of these electron buckets. If you have 3 electron buckets, they will be arranged in a planar manner, at 120°. If you have only 2 electron buckets, they will be 180° apart. Now let's work through what happens if you have various combinations of lone pairs and bonding orbitals.

If you start with 4 "electron buckets" and all 4 are used to form a bond, you'll have a classic tetrahedral geometry like that of CH4. If only 3 bonds are formed, and 1 lone pair remains, you're essentially just filling in 3 of the 4 available slots with bonds. These three bonds will be arranged in a pyramidal shape (NH3 is an example). If you have 2 bonds and 2 lone pairs, you're only filling in 2 of the slots, so you will get a bent shape (H2O is the classic example of this). The bond angles in the tetrahedral orientation are ~109.5°, but in the pyramidal and bent orientations, they are further compressed to about ~107° and ~105° due to the lone pairs (but that's not a highly important factoid for the MCAT).

If you start with 3 "electron buckets" and use all 3 to form a bond, you'll have a trigonal planar shape, like AlBr3. (Note that this can happen with a molecule like AlBr3 that has 3 single bonds and no lone pairs, or in a molecule like formaldehyde (H2CO) that has two single bonds and a double bond -- in this metaphor, an "electron bucket" can correspond to a double or triple bond). However, if you only use 2 of these "electron buckets" and have a lone pair left over, like in SO2, you'll have a bent geometry, although this time the bond angle will be slightly less than 120°, because you're starting with a trigonal planar orientation with a baseline angle of 120°.

Finally, if you start with only 2 "electron buckets," you're bound to have a linear geometry, because they're located at 180° from each other. This can happen if both "buckets" are used to form bonds, as in CO2, or if a lone pair remains, as in the negatively charged carbon in an acetylide anion.

This way of looking at things is somewhat simplified and doesn't build connections with orbital hybridization, but I've found it to be a simple and spatially intuitive way of visualizing the shape of a molecule without getting bogged down in terminology that may be more of a hindrance than a help. Hope this clarifies things, and best of luck!

thank you so very much!!

 

[DetroitDoc1993]

I thought MCAT questions about geometry and shape were pretty direct. geometry meant only looking at number of bonds while shape meant looking at both bonds and lone pairs. however, I came across this question (below). In the explanation of choosing between SiCl4 and XeF4 they mention the lone pairs too. I guess I wonder what I should focus on when asked geometry or shape: the bonds, the lone pairs, the electron groups?

You should think of it as Electron geometry influences molecular shape, everything depends on how many electrons then it is followed by how many bonded atoms. You should always looks at electrons first they dictate just about everything including molecular shape so do not think of it as how many bonds you have you should think how many electrons, followed by lone pairs and bonds. It is a key concept in bonding and if you understand the concept you really don't have to memorize anything. For example, it is the electron geometry of water that influences its molecular "bent" shape and influences its polar properties. In your question Si has the same number of valence electrons as Carbon, where as Xe has more valence electrons which effects its molecular shape differently and that is why the answer was D.