7 Factors that stabilize negative charge in organic chemistry

Like I wrote about last time, it’s good – but not enough – to recognize partial charges and to figure out where they interact.

Since reactions involve processes that lead to the gain or loss of charges, understanding the factors that stabilize (or destabilize) charge have a tremendous impact on how likely a reaction is given to occur! Let’s talk about negative charge today.

Let’s talk about a concrete example. For instance if a reaction leads to the formation of a very unstable negative charge, it’s unlikely to occur. But if it leads to the loss of a very unstable negative charge, it’s considerably more likely.

For instance, that’s why one of these reactions of methane is likely and the other is unlikely. That’s going to be explored in more detail in future posts.

So what factors lead to the stabilization of negative charge? Two main things.

negative charge tends to be less stable when it’s concentrated and more stable when it’s dispersed.

Think about that as you look at this list of seven factors that stabilize negative charge.

1. High charge densities are unstable

This one’s fairly straightforward to understand. High charge densities are unstable. So as we move from water to OH(-) to O(2-), we are getting progressively more unstable here.

2. Electronegativity

Electronegativity is a rough measure how effectively the positively charged nucleus of an atom can “pull” electrons toward it. (Opposite charges attract.) Electronegativity increases as we go across the periodic table. So if you compare the anions going from C , N, O to F across the periodic table, the stability of the negative charge will increase.

3. Polarizability

Down the periodic table, it’s a little more helpful to think “dispersal of charge is good!” rather than “opposite charges attract”. Compare fluorine and iodine. The size of the fluorine ion (radius: 119 pm) is much smaller than iodine (radius: 206 pm). However, they both have a charge of negative 1.

Imagine two balls, each weighing one pound. But one is made of iron, and the other is made of rubber. Which ball is going to be smaller? The iron ball (smaller and harder) is like fluorine, and the rubber ball (larger and squishier) is like iodine. And a certain “squishiness” helps to stabilize charge, since it isn’t as concentrated over a small volume. That’s a way of expressing the greater polarizability of iodine.

4. Resonance

Along the same lines, a negative charge that is adjacent to one or more Pi bonds can disperse its negative charge over multiple atoms. We describe this phenomenon as “resonance”. So in the example below, the negatively charged alkane on the left is much less stable than the adjacent negatively charged species, where the negative charge can be dispersed over multiple carbons through resonance.

5. Electron withdrawing groups.

This one falls more into the auspices of “opposite charges attract”. A negative charge that is adjacent to an atom with electron withdrawing groups on it will be stabilized greater than one that is not. In the extreme case of CCl3(-), the resulting ion is many orders of magnitude more stable than H3C(-) itself. (This is the basis of the haloform reaction).

6. Orbitals.

s orbitals are closer to the nucleus than p orbitals are. So electrons that are in s orbitals will be closer to the nucleus than electrons in p orbitals – and therefore, lower energy (“opposite charges attract”). For this reason, electrons that are in sp orbitals are lower energy than sp2, which is lower energy than sp3, since they have greater s character (33% for sp2) than sp3 (25%). This makes the anions more stable.

7. Aromaticity.

This is a special case, covered in detail in organic chemistry 2. Certain molecules possess a special stability – called aromaticity – that is enormously stabilizing, kind of like qualifying for an huge tax break from the government. Certain negatively charged molecules – such as the cyclopentadienyl anion, pictured below – are aromatic, and therefore possess much greater stability than they would have otherwise.

So how the heck do we keep track of all of this?

Seven factors?!!! So how do we know which is most important? That’s a great question! These trends can interact with each other in unpredictable ways, and it’s hard to judge which is most important.

Thankfully, there’s a concept you’ve probably already met for figuring out the stability of these species, which can be readily measured. It’s called basicity. These factors determine how stable a base will be!

The basicity of a species tells you about how stable its lone pair of electrons are.

How do we find a good measure of basicity? Simple. It’s in the pKa table, a collection of measurements that’s been compared to the table of hand strengths in poker.

Bottom line:

Two factors to watch out for: opposite charges attract, and dispersal of charge.

unstable anions will tend to be at the intial tails of arrows (form bonds).

stable anions will tend to be at the final heads of arrows (likely to be leaving groups)

This might be a stupid question but i’m a little confused as to why fluorine is the most electronegative and most stable in figure 2, but in figure 3 fluorine is the least polarizable and least stable? How would you know when to consider it electronegative or polarizable?

It would be difficult to know from first principles whether or not electronegativity or polarizability is most significant in determining the stability of anions. Measurement of acidities, however, has given us the ability to evaluate this on a universal scale (acidity). We can observe the empirical trend (electronegativity is most important as we go across a row of the periodic trend, polarizability is most important as we go down), but explaining exactly why this is is a deep and difficult question to answer.

Hi,
Please help with options. School project – to charge perspex negatively to extract polystyrene from sand. We used perspex sheet and rubbed it with wool roller to load perspex negatively, but the charge fluctuates and is not lifting polystyrene constantly. Suggestions as to how to increase to negative charge to be stronger?

Hello James,
I wanted to thank you for this awesome project of yours here, I really wish lecturers would pay more attention to intuition and proper teaching as you do here instead of just showing mechanisms.
I am struggling to understand a couple of concepts here and would love to hear what you think about them:
1. You wrote that high charge densities are unstable and followed it up by writing that electronegative atoms are more stable, which appears to me as contradictory. Isn’t a more electronegative atom/ion necessarily creates a high charge density area due its strong electron attraction?
2. Following your reasoning sp orbitals are more stable than sp2 than sp3, as they have more s character and are therefore closer to the nucleus and lower in energy, but we also get more pi bonds (that are higher in energy) for fewer hybridizations in such molecules, while in more hybridized molecules these pi bonds will have at least some s character that will lower their energy. Aren’t these reasonings contradictory as well?
Thanks for your time!

1. “High charge densities are unstable” is a good rule of thumb, although it’s not easy to look at this quantitatively because some of the variables are different. It’s more of an intuitive principle. For example, compare the stabilities of H2O vs HO(-) vs O(2-) , or NH3 vs NH2(-) vs NH(2-) vs N(3-).
Given that, if one is going to have a negative charge, negative charge will be stabilized *more* by a more electronegative element. This trend is most evident going across the periodic table – for example F(-) versus HO(-) versus NH2(-) versus CH3(-). Fluorine is the most electronegative element in this series and F(-) is also the most stable; carbon is the least electronegative element and CH3(-) is the least stable.
Going down the periodic table it is not as straightforward because valence electrons are held in orbitals more distant from the nucleus and polarizability is more important.

2. Pi bonds don’t have any s character. Pi bonds are made solely from the overlap of p orbitals. Just to make sure this is clear, in a molecule like ethene where the two carbons are attached by a double bond, only one of those bonds is a pi bond – the other one is a sigma bond between the two sp2 hybridized carbons.

Sorry, I meant sp(n) hybridized orbital (instead of pi), say CN- vs. CH3-, CN- has a low energy sp orbital and two high energy p orbitals vs. CH3- that has three hybridized orbitals that are between the two (energy wise).

The charge is localized in the sp(n) orbitals so the p orbitals in CN(-) are irrelevant to the stability.

One way of looking at it is that an sp orbital has a higher effective electronegativity than an sp2, which has high effective electronegativity than an sp3 orbital. The negative charge is held more tightly to the nucleus.

Thank you, I think I am starting to get it. This is analogous to electronegativity since in both cases (higher sp character / electronegativity) the bond is pulled closer towards the nucleus.
Would it be right to say that the delocalization of the p orbitals between both atoms in CN(-) also plays a part in stabilizing the negative charge?

Idan

Edit: Noticed that the lone pair on the carbon is localized so it can’t be stabilized through delocalization so never mind that last question :).
Thanks again.

About Master Organic Chemistry

After doing a PhD in organic synthesis at McGill and a postdoc at MIT, I applied for faculty positions at universities and it didn’t work out, yada yada yada. So I decided to teach organic chemistry anyway! Master Organic Chemistry is the resource I wish I had when I was learning the subject.