When will hybridization states end?!

[ArkansasRanger]

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So I'm taking university chemistry I otherwise known as general chemistry on the boards here, but at school only nurses and ag. business majors get credit for general chemistry. It's considered a diluted course.

We're on chapter four: covalent bonds. I took general chemistry I nine years ago at a school where science majors received credit. We didn't even touch on hybridization. I've never heard of it before. I understood the octet rule and Lewis Structures, but now there are so many violations of the rule that I can't draw one diagram without blowing a gasket. I understood electron domains. That was cake, but molecular geometry was thrown in on it, and maybe I'm just stupid. However, I do not see a seesaw in any of this. Now, we've got hybridization and despite my text book, the instruction of Dr. Useless, and General Chemistry I as a Second Language I just don't get where it fits. Why are we learning it? What's it do? Is it connected to part II of this course or organic chemistry?

I got my first exam back today. I made a 78 which was six points higher than the class average. I think everyone in there is some kind of pre-doctor something. Basically, I got every equation problem wrong, but hell I've always sucked at applying math to a given situation so I figured that was coming.

What joys are next?

End of rant.

 

[crazybob]

The single bond, double bond, and triple bond are where you try to make sense of the hybridization. It also comes in to use with the bond angles.

 

[phathead]

The single bond, double bond, and triple bond are where you try to make sense of the hybridization. It also comes in to use with the bond angles.

Yep, being able to see where free electrons are and whether or not they are in a sigma or pi bond is quite helpful when tracking electron movement during reactions in o chem.

At least I thought so

 

[ArkansasRanger]

Yep, being able to see where free electrons are and whether or not they are in a sigma or pi bond is quite helpful when tracking electron movement during reactions in o chem.

At least I thought so

Yeah, sigma and pi bond was covered today as well. I hope he goes over it again. He gave it all of 15 minutes.

 

[qmoto]

We learned hybridization schemes, bonding schemes, and orbital overlap diagrams in gchem I. I think understanding the bonding schemes is the key. Also remember that for every double bond you have 1 sigma and 1 pi bond, for a triple bond you have 1 sigma and 2 pi bonds.

Take this question for an example, it was one of my quiz questions last semester.

Draw the bonding scheme for the molecule HCONH2.

Do your 6👎+2 to get 6(3)+2 = 20 - 18 #VE = +2

Draw your lewis structure followed by the VSEPR structure. You should get a trigonal plannar and trigonal pyramid with C and N being the central atoms and a double bond on C to O. Once you have your VSEPR structure you can do the bonding scheme.

C-H sigma bond: C(sp2) - H(1s)
C=O sigma bond: C(sp2) - O(2p)
C=O pi bond: C(2px) - O(2px)
C-N sigma bond: C(sp2) - N(sp3)
2x N-H sigma bond: N(sp3) - H(1s)

 

[hopeful4]

That was cake, but molecular geometry was thrown in on it, and maybe I'm just stupid. However, I do not see a seesaw in any of this. Now, we've got hybridization and despite my text book, the instruction of Dr. Useless, and General Chemistry I as a Second Language I just don't get where it fits. Why are we learning it? What's it do? Is it connected to part II of this course or organic chemistry?

What joys are next?

End of rant.

You are not stupid, just lost in the world of chemistry! If you plan on going on in chemistry it will start to make sense. It definitely was a huge part of organic. I suggest getting a model kit and building the appropriate molecules to help visualize the geometry. It is much harder to see on paper and my model was my best friend to try and get through all of that crap! The more I learned in chemistry the more all of the crazy stuff made sense. Just hang in there!

 

[b1234]

We learned hybridization schemes, bonding schemes, and orbital overlap diagrams in gchem I. I think understanding the bonding schemes is the key. Also remember that for every double bond you have 1 sigma and 1 pi bond, for a triple bond you have 1 sigma and 2 pi bonds.

Take this question for an example, it was one of my quiz questions last semester.

Draw the bonding scheme for the molecule HCONH2.

Do your 6👎+2 to get 6(3)+2 = 20 - 18 #VE = +2

Draw your lewis structure followed by the VSEPR structure. You should get a trigonal plannar and trigonal pyramid with C and N being the central atoms and a double bond on C to O. Once you have your VSEPR structure you can do the bonding scheme.

C-H sigma bond: C(sp2) - H(1s)
C=O sigma bond: C(sp2) - O(2p)
C=O pi bond: C(2px) - O(2px)
C-N sigma bond: C(sp2) - N(sp3)
2x N-H sigma bond: N(sp3) - H(1s)

Wow, somehow you managed to make that way more complicated than it really is...

Also, you're wrong... a 2p orbital from oxygen will not form a sigma bond with anything...

 

[qmoto]

I appreciate the *constructive* criticism. If you have a easier way of answering the question by all means share it. I am simply showing the method I was taught and we had to show all these steps to get full credit. I am curious as to what you would put for the C=O in that molecule since you say it is wrong. On my quiz I was marked correct for this answer.

 

[Project2015]

I think it's 2s while the p's are pi bond and lone pairs...

Or... sp2 for sigma and two lone pairs with remaining p as pi bond so that it's trigonal planar?

 

[b1234]

I appreciate the *constructive* criticism. If you have a easier way of answering the question by all means share it. I am simply showing the method I was taught and we had to show all these steps to get full credit. I am curious as to what you would put for the C=O in that molecule since you say it is wrong. On my quiz I was marked correct for this answer.

Maybe you made a typo in your first post or you instructor just missed it. Sigma bonds are made from MO's with at least some S character - not from P orbitals. Pure P orbitals make pi bonds. They can also have lone pairs used in conjugated pi systems (very similar to pi bonds) and they can be used to backbond when donating to the d orbitals in a metal-ligand complex.

I really don't want to go through that molecule - but I know I would never put down a p orbital for a sigma bond. The oxygen in a carbonyl will normally be sp2 hybridized - 3 sp2 MO's and one P orbital. The P orbital will be the pi bond and the sp2 orbitals will make up the sigma bond and two lone pairs.

The whole two equation thing you used to figure out how many valence electrons just looked very tedious to me - sucks that you were forced to do that to get full credit. I read through your method and thought - WTF? If it works for you great, but trying to teach that method to someone else... It seems way more complicated than it has to be.

I would normally look at a molecule and draw in all the bonds, then the lone pairs - making sure each atom had the correct number of electrons. Then I'd count the total electrons just to make sure I had everything right. From there - pi bonds are P orbitals and everything else is spx. with x being 3 minus the number of pi bonds that atom made (ie the number of p orbitals that weren't used to form pi bonds). Hydrogen of course only has s orbitals.

Once you have the bonds and lone pairs and all the atoms drawn in. The geometry should be fairly easy to determine. For any given atom just count up how many constituents - I usually treat lone pairs like atoms for this part... not really going to go into detail on this - either you understand it or you don't.

 

[b1234]

I think it's 2s while the p's are pi bond and lone pairs...

Or... sp2 for sigma and two lone pairs with remaining p as pi bond so that it's trigonal planar?

The latter for the oxygen in a carbonyl. P orbitals don't normally make lone pairs unless they are part of a conjugated pi system or being used to bond in metal-ligand complexes. You need to get farther into MO theory to deal with pi back bonding in metal-ligand complexes though.

 

[WVUPharm2007]

Thank god I don't have to know any of this stupid **** in the real world.

 

[byork]

Wow. That looks really hard. We covered it in detail in Grade 12 but barely talked about it in general chem. Thank god I'm done all that crap.

 

[By3Times1Minus1]

you will definitely need to know about hybridization once you get into organic. i had a fair amount of questions on the PCAT as well, so i'm sure it'll be on the MCAT if you're pre-med. make sure you understand it now because it will make organic and studying for your entrance exam that much easier later. if your teacher isn't making sense to you, try going online and finding something that will. i didn't search around much but here are a couple pages that might help from a quick google search. you can probably find something better if you look around for what makes sense to you though.
http://www.mhhe.com/physsci/chemistry/carey5e/Ch02/ch2-2-3.html (click the arrows at the bottom to keep going)
http://en.wikipedia.org/wiki/Orbital_hybridisation

 

[joetrisman]

Thank god I don't have to know any of this stupid **** in the real world.

I totally agree, but at the same time, I know a LOT more people that would have been applying to be vets/docs/pharms if they had been able to pass orgo. I'm really surprised that biochem isn't a prereq to weed out students for most programs.

 

[MiataMike]

I'm really surprised that biochem isn't a prereq to weed out students for most programs.

It is for some schools. I'm taking it in my 1st semester.

 

[joetrisman]

It is for some schools. I'm taking it in my 1st semester.

Yeah, I only knew that one school east had it like UNC or something as a prereq.

 

[tungsten87]

Yeah, I only knew that one school east had it like UNC or something as a prereq.

UTenn requires two semesters of biochem as well.

 

[qmoto]

Here is the quiz I was mentioning. Perhaps this will give you an idea.

 

[Passion4Sci]

Wow, you kept stuff from Gen Chem. :O

 

[Project2015]

The latter for the oxygen in a carbonyl. P orbitals don't normally make lone pairs unless they are part of a conjugated pi system or being used to bond in metal-ligand complexes. You need to get farther into MO theory to deal with pi back bonding in metal-ligand complexes though.

Yeah I vaguely remember M-L complexes. I've heard that some people don't use hybridized orbitals for lone pairs, thus the first answer. In that case, the geometry doesn't make sense since the electrons are not in an optimal configuration (i.e., minimal repulsion).

 

[NaOH]

Wow, you kept stuff from Gen Chem. :O

I thought everyone did? 😕

<-(has boxes of old notes that do nothing but take up closet space)

 

[Passion4Sci]

I thought everyone did? 😕

<-(has boxes of old notes that do nothing but take up closet space)

I definitely do not.

After the quarter/semester, right into the recycling bin.

 

[b1234]

Yeah I vaguely remember M-L complexes. I've heard that some people don't use hybridized orbitals for lone pairs, thus the first answer. In that case, the geometry doesn't make sense since the electrons are not in an optimal configuration (i.e., minimal repulsion).

If you have a lone pair that is part of a conjugated pi system they go in a non hybridized p orbital. In all other cases I can think of they go in orbitals with at least some s character. It makes more sense that way too. If you think about the shapes of the orbitals, electron-nucleus attractions, and repulsions between the nuclei of different atoms. Lone pair electrons will actually fill orbitals that are more s-character than bonding electrons - most people don't need to know this though.

Here is the quiz I was mentioning. Perhaps this will give you an idea.

Notice no check mark by where you wrote 2p orbital for the sigma bond. Either that particular portion wasn't checked - or whoever was grading it just missed it. Sigma bonds have some s character. If you analyze your answer a bit more, you will find that an sp2 orbital and a p orbital probably will not overlap at all or at least have minimal overlap.

For the most part I will agree with the rest of your answer, but I wonder why you put 6(3)+2=20-18=+2 if you're trying to find the number of valence electrons, I'd use 6+4+5+1(3)=18
Putting a bond angle between a lone pair of electrons and an atom isn't really correct - that was probably just an oversight - it's fairly obvious what you meant by that.

 

[Project2015]

Yeah you can't have bond angle between a bond and nothing. 😛

In any case, it should be > 109.5 due to the increased repulsion of the closer-to-the-nucleus lone pair.

Ahhh.. the chemistry is starting to come back.

 

[b1234]

Yeah you can't have bond angle between a bond and nothing. 😛

In any case, it should be > 109.5 due to the increased repulsion of the closer-to-the-nucleus lone pair.

Ahhh.. the chemistry is starting to come back.

Yeah but at the level of that class I doubt they got into the slightly less or greater than bond angles so for that class it's probably straight, or 109.5 or 120.

 

[Praziquantel86]

I forgot how awful all that was...

Good luck to you.

 

[ArkansasRanger]

Well, for now it's over. The test is Tuesday. We got three days of covalent stuff and tomorrow's lecture is our second day of ionic stuff. Then we'll be done with that unit. Up next week: gas. Oddly enough tomorrow's lab is about gas stuff when we have yet to cover any of it. 😕

 

[MiataMike]

Haha yes, my labs frequently covered material that only the next lecture would catch up to. Lab manual should have all you need to know, though.