Equatorial and Axial Positions - Video Tutorials & Practice Problems
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Covalent compounds with 5 or 6 electron groups have equatorial and axial positions for surrounding elements
Equatorial and Axial Positions
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Equatorial and Axial Positions Concept 1
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Covalent compounds with 5 or 6 electron groups have equatorial and axial positions for their surrounding elements. Now when we say the equatorial position, we're gonna say the equatorial position. These are your surrounding elements position around the equator of a compound. So here we take a look at these two illustrations, we can say that the equator of this sphere is right here. And here we have 5 electron groups, 3 of them are along the equator. Here we have 6 electron groups and 4 of them are along the equator of this sphere. Now if we say axial or apical position, this is basically a surrounding element's position above or below the equatorial position. So, remember, we have our equator here, so our axial positions are above it or below it. And here, above it or below it. Now these arrangements themselves, they increase repulsion between elements. This in turn causes a decrease in energy for the compounds. And just remember in chemistry, if we decrease our energy, that's a good thing. That leads to greater stability. Now a rule of thumb is we're gonna say that the more electronegative element tends to prefer the axial position over the equatorial position. Okay. So this again ties into the whole idea of energy and stability. So just remember these 5 points when we're talking about 5 electron group molecules and 6 electron group molecules.
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Equatorial and Axial Positions Example 1
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Based on your knowledge of axial and equatorial positions, draw the most likely structure of 2 c l 3. Alright. So remember, when it comes to drawing Lewis dot structures, we place the least electronegative element in the center. In this case, it'd be phosphorus. Phosphorus has attached to it 5 bonding groups, 2 fluorines and 3 chlorines, so it's a 5 electron group system. Now, phosphorus is in group 5 a, so this makes sense. Remember that 3 of them would be along the equator, and 2 of them would be in the axial positions. Now remember, in terms of stability and energy, we want the more electronegative element to be in the axial positions. Fluorine is more electronegative than chlorine, so fluorine we place in the axial positions. So remember, fluorine is in group 7a, so it only has 7 valence electrons. Halogens, like chlorine and fluorine, when they're not in the center, they only make single bonds as surrounding elements. And then we draw the chlorines. Chlorines are in group 7 a as well, so they have 7 valence electrons. Then here we're drawing them here. They're along the equator. So this would be the most stable, most likely structure of 2 c l 3. We have a 5 electron system and therefore we have equatorial and axial positions. Fluorine is more electronegative than chlorine, So remember, fluorine would go into the axial positions, and chlorine will go into the equatorial positions.
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Equatorial and Axial Positions Concept 2
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Lone pair positions are extremely important when trying to draw the most accurate molecule. Here we're going to say that lone pairs will orient themselves in order to decrease the interactions between surrounding elements. Now here when we're drawing different structures we're paying more attention to those with 6 electron groups and those with 5 electron groups. We're going to say here that with 6 electron group systems, lone pairs are most stable in the axial position. Within 5 electron group systems, lone pairs are most stable in the equatorial position. Now you might ask, how do I remember this? Well, we have a memory tool for that and the memory tool is, just remember it's a lock as long as you remember the hands of the clock. So we take a look here at the first clock, we have the hands giving us a time of 6 o'clock. Now let's just imagine that that circle in the middle, that dark circle represents our central element. And here, this would represent our equator. We're going to say, we said when it deals with 6 electron groups they are in the axial positions. So remember, axial positions are straight up and straight down, so above and below the equator of our structure. And if we look they're pointing at 6 o'clock. So we're going to say 6 o'clock is for 6 electron groups. The hands are pointing straight up and straight down just like axial positions point straight up and straight down. For fine electron group systems, again, we imagine that we have our equator here. Both the hands are not pointing straight up and straight down, so they wouldn't represent axial positions they represent equatorial positions. So here we'd say that 5 o'clock is for 5 electron groups. Right? So just remember, if you can remember 6 o'clock straight up and straight down those are axial positions, then you remember lone pairs are most stable in the axial positions. If it's at 5 o'clock that means both hands are not straight up and straight down, so they couldn't be axial positions they would be equatorial positions. So in those cases, lone pairs should be oriented in the equatorial positions to draw the best possible structure. So keep that in mind when we deal with 5 electron group systems and 6 electron group systems.
It's a lock as long as you remember the hands of the clock.
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Equatorial and Axial Positions Example 2
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For this example question it says, Determine the molecular geometry of the following ions. So here we're dealing with SCl3 negative. Alright, so here we're going to place sulfur as our central element. Sulfur itself is in group 6A so it has 6 valence electrons. And here it's going to use 3 out of the 6 to make connections to our 3 chlorines. Chlorines are in group 7a, so they each have 7 valence electrons. They use one of their electrons to make a single bond to our sulfur atom and we're going to say this is what our structure would look will look like. So the chlorines have used all their electrons available, sulfur has only used 3 out of its 6 total electrons. So here we're gonna draw its remaining 3. Now, we're going to go back and clean this up based on our understanding of 5 electron systems versus 6 electron systems. So right now, let's just draw the basic framework of our structure. So there goes 5 valence electrons and here's the 6. -one means we've gained an outside electron. The chlorines no longer need an extra electron. They're all, making the right amount of bonds that they want. So that extra electron that we're gaining, because it's minus 1, it's going to go to the sulfur. And because our structure has a charge we need to draw it in brackets with the charge on the outside. Now we need to determine, is this a 5 electron system or a 6 electron system? Well if we look we're going to say that we have 1, 2, 3, 4, 5 electron groups. So this is a 5 electron group system. So think about our memory tool. It's a long as long as you remember the hands of the clock. So 5 groups means we're looking at 5 o'clock. Both hands aren't pointing straight up or straight down so they're not in the axial position. They'd be in the equatorial position. That means our lone pairs should be oriented in the equatorial position to draw the best possible structure. So here we're going to place sulfur in the center, so we're going to draw the correct structure over here. Sulfur goes in the center. The 2 lone pairs that it has, we have to draw them in the equatorial position. And one of the chlorines will also be, in the equatorial position as well. So there it goes. And then the other 2 are going to be in the axle positions, pointing straight up and straight down. Okay. So there goes our structure. Now remember, draw this one as our extra electron. Remember because the structure has a charge we put it in brackets with the charge on the outside. So this would be the correct way of drawing Scl3 negative. Right? So just remember, when you're dealing with a 6 electron group system or a 5 electron group system it's important to remember where should my lone pairs go. Knowing that gives you the best possible structure.
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Problem
Problem
Draw the most likely shape for the following compound:XeF4
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Problem
Problem
Draw and determine the geometry for the following molecule:Br2CO
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Problem
Problem
How many lone pairs reside in the equatorial position of the KrCl5+ ion.