Lewis structure

[hypervalent_iodine]

2 hours ago, Rachel Maddiee said:

Carbon has formed three covalent bonds in the structure so it would have 7 electrons around it (not 8 as shown) Likewise the oxygen fails to follow the rule.

No. This was wrong in your first answer and it’s still wrong here. Both the carbon and the oxygen in the Lewis structure in the question have 8 electrons around them. Both satisfy the octet rule. I’m not sure how else to explain this, so perhaps I will work on something later and you can review the posts in this thread in the meantime. 

2 hours ago, Rachel Maddiee said:

There should be two pairs of electrons (a double bond) between the C and O, and no lone pair on the C.

Yes, but how do you know that? Saying that carbon has to have four bonds isn’t good enough reasoning, because that statement isn’t always true. Same for oxygen. Please go through the post I made in response to MigL. It works through this aspect of the problem.

4 hours ago, Rachel Maddiee said:

Incorrect structure: 8 bonding electrons and 2 lone pairs Number of electrons = 6 + 8 = 14 total electrons 

Correct structure; CH2O There are 4 valence electrons in carbon, 1 each in hydrogen and 6 in oxygen, so there are 12 electrons total

Great! So what does this tell you about the Lewis structure? 

4 hours ago, studiot said:

I suggest you stick to the dots.

Because of the many exceptions the dot structure is taught and then students move on.

So I also suggest you do enough to earn the marks (since you must and HI doesn't need to) but remember better methods follow

My comment was made for two reasons. Based on the OP, Rachel is perfectly comfortable using lines to represent bonds and can tell that each bond has two electrons in it. The second reason is that it is not standard notation and when you get to later topics or even when you get further into drawing Lewis structures, using dot notation to represent bonds is confusing. Many higher institutions will mark you wrong or deduct marks for using this notation - it is a bad habit to fall in to. 

It is truly the bane of our first year teaching staff’s existence. They take this notation from this topic and start drawing bonds as dots when covering, for example, organic chemistry. We have systematic methods of drawing structures for a reason! The other rather irritating part is that it is taught differently and inconsistently in different schools and so there is a lot to have to stamp out. 

[Rachel Maddiee]

It reels me that there are too many electrons in the structure but I’m still struggling with the specific rule.

[hypervalent_iodine]

6 minutes ago, Rachel Maddiee said:

It reels me that there are too many electrons in the structure but I’m still struggling with the specific rule.

That’s right. If you only have 12 valence electrons available to make up the structure, as you do, you can’t then wind up with 14 electrons. So the Lewis structure must be wrong (too many electrons!). To get to 12 electrons, we have to put in a double bond. Once you do that and you count the number of electrons again, you should have the right number. Have a look at the structure in your very first post with the correct Lewis structure and count the electrons. 

[Rachel Maddiee]

So why does the question ask to explain which rule the structure does not obey? Since it satisfies the octet rule.

Hw would I write it?

[hypervalent_iodine]

“The rules of drawing Lewis structures,” includes having the correct number of electrons. Given that, and given everything else we have discussed, how do you think you should write it?

[Rachel Maddiee]

The Lewis structure is not correct. There should be two pairs of electrons (a double bond) between the C and O, and no lone pair on the C. Carbon needs 4 bonds and Oxygen 2. Carbon has 4 valence electrons and oxygen has 6 valence electrons. Carbon shares two electrons with the two Hydrogens, and shares four electrons with the Oxygen, (a double bond.)
Correct structure: 2(1 for each H) + 4(for C) + 6(for O) = 12 valence electrons
Incorrect structure: 8 bonding electrons and 2 lone pairs 
Number of electrons = 6 + 8 = 14 total electrons 

The Lewis structure is incorrect because it has too many electrons (to get to 12 electrons, we have to put in a double bond)

Edited January 7, 2020 by Rachel Maddiee

[hypervalent_iodine]

The only addition I would make is to include a sentence that states that the Lewis structure does obey the octet rule, since both the carbon and oxygen have 8 electrons around them.

[Rachel Maddiee]

Okay, thank you so much!!

[studiot]

2 hours ago, hypervalent_iodine said:

7 hours ago, studiot said:

I suggest you stick to the dots.

Because of the many exceptions the dot structure is taught and then students move on.

So I also suggest you do enough to earn the marks (since you must and HI doesn't need to) but remember better methods follow

My comment was made for two reasons. Based on the OP, Rachel is perfectly comfortable using lines to represent bonds and can tell that each bond has two electrons in it. The second reason is that it is not standard notation and when you get to later topics or even when you get further into drawing Lewis structures, using dot notation to represent bonds is confusing. Many higher institutions will mark you wrong or deduct marks for using this notation - it is a bad habit to fall in to. 

It is truly the bane of our first year teaching staff’s existence. They take this notation from this topic and start drawing bonds as dots when covering, for example, organic chemistry. We have systematic methods of drawing structures for a reason! The other rather irritating part is that it is taught differently and inconsistently in different schools and so there is a lot to have to stamp out. 

 

However GN Lewis did not use lines.

Further the OP scan/photo of the actual question paper shows a figure that does not contain lines, but relies totally on the Lewis dot notation.
The question was "Is that dot diagram correct or incorrect and if it is incorrect draw a correct diagram.?"
I would suggest that since the diagram was incorrect only in the placement of the dots, and Rachel recognised this, a corrected one should include only dots.

I would venture to suggest that most of the problems arising from using 'Lewis Notation' arises from later persons changing the original for their own ends, particular mixing lines and dots is IMHO bad practice.

Here is an extract from the original 1916 paper.

Nowhere in the 25 pages is the line used to represent a bond.

It should be remembered that decade before Schrodinger, Heisenberg and London  developed orbital theory 10 to 15 years later.

It was even 3 years before the Pauli exclusion principle.

The important point about Lewis is that here was the first suggestion of sharing electrons.

 

Edited January 8, 2020 by studiot

[hypervalent_iodine]

27 minutes ago, studiot said:

 

However GN Lewis did not use lines.

I would venture to suggest that most of the problems arising from using 'Lewis Notation' arises from later persons changing the original for their own ends, particular mixing lines and dots is IMHO bad practice.

Here is an extract from the original 1916 paper.

Nowhere in the 25 pages is the line used to represent a bond.

It should be remembered that decade before Schrodinger, Heisenberg and London  developed orbital theory 10 to 15 years later.

It was even 3 years before the Pauli exclusion principle.

The important point about Lewis is that here was the first suggestion of sharing electrons.

 

There are a lot of things we believed or did in chemistry in 1916 that we no longer do and no longer teach, for good reason. “It’s the way Lewis himself did it,” isn’t a good enough reason to teach students notation that can cause later confusion more than 100 years after the original publication came out. 

[studiot]

14 minutes ago, hypervalent_iodine said:

There are a lot of things we believed or did in chemistry in 1916 that we no longer do and no longer teach, for good reason. “It’s the way Lewis himself did it,” isn’t a good enough reason to teach students notation that can cause later confusion more than 100 years after the original publication came out. 

So there is even less reason to teach a bastardised version.

My suggestion was to use exactly what was asked and no more and then move on to better things.

[hypervalent_iodine]

12 minutes ago, studiot said:

So there is even less reason to teach a bastardised version.

My suggestion was to use exactly what was asked and no more and then move on to better things.

Not sure what you’re talking about here. My method isn’t a bastardised version; it’s how chemists draw structures when including lone pairs. It’s how Rachel (correctly) drew her structure in her first post. What I was saying was that drawing covalent bonds as electron dots is a bad habit (eg C::O instead of C=O). I didn’t say anything about lone pairs. Without the lone pairs it wouldn’t be a Lewis structure. Sorry if that wasn’t clear in my posts. 

[studiot]

5 minutes ago, hypervalent_iodine said:

Not sure what you’re talking about here.

I wasn't referring to your method in particular.

I must have looked through at least a couple of dozen books and websites, including UK official teaching ones for Lewis Notation and do you know what?

Every one of them is different.

I think that is also the point you are making in your own way.

 

[hypervalent_iodine]

32 minutes ago, studiot said:

Every one of them is different.

I think that is also the point you are making in your own way.

More or less, yes. Introductory chemistry is taught in a lot of different ways, some good, some bad. In Australia, teaching content is not standardised and teachers can essentially pick and choose what they want to teach and how they want to teach it (within reason). This can be good, but it also makes the lives of course coordinators at universities difficult when you have to deal with classes with such a varying base knowledge. I, for example, had never been taught how to calculate pH when I got to university. 

Honestly, I wouldn't have a problem with using Lewis dot notation for bonds if it were only used within that one context for teaching students about constructing Lewis Structures and  if I didn't know how ingrained that notation can become. That being said, it can get very messy when you start doing it for atoms with more than four bonds (PCl₅ for example). I've marked and taught first year pretty extensively for I guess the last decade now, and even after an entire 1-2 semesters of being told not to draw structures like that, we still see a significant number of students drawing out reaction schemes with dots for bonds. The one I actually dislike the most is the use of a line to represent a lone pair - it looks like a negative charge! Yet without fail, every semester, despite never having been taught in first year to show lone pairs like that, there is always a handful of students who do it. 

[studiot]

9 hours ago, hypervalent_iodine said:

More or less, yes. Introductory chemistry is taught in a lot of different ways, some good, some bad. In Australia, teaching content is not standardised and teachers can essentially pick and choose what they want to teach and how they want to teach it (within reason). This can be good, but it also makes the lives of course coordinators at universities difficult when you have to deal with classes with such a varying base knowledge. I, for example, had never been taught how to calculate pH when I got to university. 

Honestly, I wouldn't have a problem with using Lewis dot notation for bonds if it were only used within that one context for teaching students about constructing Lewis Structures and  if I didn't know how ingrained that notation can become. That being said, it can get very messy when you start doing it for atoms with more than four bonds (PCl₅ for example). I've marked and taught first year pretty extensively for I guess the last decade now, and even after an entire 1-2 semesters of being told not to draw structures like that, we still see a significant number of students drawing out reaction schemes with dots for bonds. The one I actually dislike the most is the use of a line to represent a lone pair - it looks like a negative charge! Yet without fail, every semester, despite never having been taught in first year to show lone pairs like that, there is always a handful of students who do it. 

If you ask me the  Lewis notation has gained ground since it was a passing page or page and a half in my textbooks in the 1960s.
Compare that with  Rachel's text book : ref p253 - 260 so at least 7 pages.
This is borne out in the space devoted to this subject in both the more recent textbooks and website I looked at.
Hardly phasing the subject out.
Which in my opinion is what should be happening.

I think a better way forward would be to say

"Here is an older notation you may well come across in reading.
So will present it in outline so if you ever need to use it or read it in the future you will have heard of it.
But we will not examine it as its use is being phased out."

That was how I saw it in 1966 and have never used it or needed it in anger since.

But today teachers have set out Lewis rules and are testing students on them.

So the only view that matters to Rachel is the examiners' and she has to present her work in whatever way they require.

[hypervalent_iodine]

1 hour ago, studiot said:

So the only view that matters to Rachel is the examiners' and she has to present her work in whatever way they require.

Sure, I’m not disagreeing with that. I’m only discussing this further because you made a point of mentioning it and suggested Rachel stick to dots when she hadn’t been using them and her structures were fine. The question used dots, sure, but that doesn’t she’ll be marked down for using the notation she did in her OP. 
 

1 hour ago, studiot said:

If you ask me the  Lewis notation has gained ground since it was a passing page or page and a half in my textbooks in the 1960s.
Compare that with  Rachel's text book : ref p253 - 260 so at least 7 pages.
This is borne out in the space devoted to this subject in both the more recent textbooks and website I looked at.
Hardly phasing the subject out.
Which in my opinion is what should be happening.

I think a better way forward would be to say

"Here is an older notation you may well come across in reading.
So will present it in outline so if you ever need to use it or read it in the future you will have heard of it.
But we will not examine it as its use is being phased out."

 

Agreed! At least here, a lot of the problem (and I’m talking beyond Lewis notation) is born from the mouths of high school chemistry teachers who aren’t really that well equipped to be teaching chemistry, largely because they aren’t required to take that much of it themselves at university. A topic for another time though.