How To Tell If Something Is Polar – Earlier, we saw how to evaluate basicamines by looking at the pKa of its conjugate acid (pKaH). (The higher the pK
H, the foundation is strong). But we still have key concepts that help us, for example, why pyridine (Pk
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H = 11) or why the nitrogen of nitrile is much less than that of amine. This is our topic today.
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Fortunately, even if you’re new to aminos, they don’t have to be. You’ve probably had trouble evaluating the acidity of certain molecules—for example, in the 5 Major Factors Affecting Acidity.
From this unit, you will recall that any factor that makes the conjugate base of a molecule more stable increases its acidity. [Remember Le Chatelier? If you make production more stable, you balance to the right].
What all “acidifying agents” have in common is that they stabilize the negative charge by inductive effects, displacement by resonance, or by bringing the charge closer to the nucleus.
Since acidity and basicity are two sides of the same coin, the main factors that affect acidity also affect the solubility of amines. [Note 1] Substantial evaluation therefore requires consideration of the same concepts, but works in the opposite direction.
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In general, the more unstable an electron pair is, the more basic it is. Therefore, using the same principles as mentioned above, the basis can be increased by removing the inductive effect, removing localization by resonance, or moving the charge further away from the core.
Let’s examine the key factors in turn and use them to derive some key trends for amine bases.
This is probably the easiest factor to assess. If “basicity” can be roughly translated as “electron pair instability”, and instability increases with charge density, then basicity should increase with increasing negative charge.
A simple way to put it: the conjugate base of an amine will always be a stronger base than the amine.
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This relationship of lower charge density leading to lower bases also applies to single pairs, which can change to larger Pi systems through resonance.
“The stronger the acid, the weaker the conjugate base”, is like saying that the conjugate base of phenol is a weaker base than the conjugate base of cyclohexanol.
Why We have already seen that this is true because the conjugate base of phenol can be stabilized by resonance while the conjugate base of cyclohexanol cannot.
The oxygen in phenol is part of a large pi system and can be distributed to the aromatic ring through electron density resonance. (Remember: lower charge density = higher stability).
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The bottom line here is that, all things being equal, a conjugated amine will be less basic than a non-conjugated amine.
Therefore, we expect that electron-withdrawing groups on amines will also decrease their play. And they do! Control morpholine (Pk
However, how do you explain why amides are significantly less basic than amines? Is it resonant? Does it have an induction effect? Is it both?
It seems worthwhile to devote a section to how nitrogenous bases are affected by interactions with other functional groups in the Pi system.
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In particular, the play of nitrogen decreases when it acts as a foot donor and increases when it acts as a foot acceptor.
The first element is an electron-withdrawing oxygen that can remove some electron density from nitrogen. However, there is an important resonance form in which the nitrogen lone pair forms a new p bond with carbon (we call it a “pi-donation”), which causes the movement of a pair of electrons from the C-O bond. pi to oxygen (we refer to this reaction as the “pi acceptor”).
See the resonance form on the right. Nitrogen no longer has a lone pair and therefore cannot act as a base.
If you’ve covered electrophilic aromatic substitution, these functional groups should look familiar. You may recognize that “leg adopters” are all over the category
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H = 9.2). Examining the DMAP resonance forms is enlightening. In the key resonance form, the nitrogen in the ring is negatively charged.
With a pi donor (see above) it becomes less basic but the pyridine nitrogen becomes more basic because it is a pi acceptor here.
Another example of how a nitrogenous base can be extended by the attachment of pi-donors is found in guanidine. In guanidine, there are two pi-donating NHs
Those of you who have studied some biochemistry will recall that arginine is the most basic of the 20 essential amino acids (pK
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You may recall that our explanation for this result was that the hybrid orbitals of alkynes have 50% s character and
Orbitals, the lone pair formed by a conjugate base, feel the positive charge of the nucleus more than the lone pair in an sp.
Hybrid orbital (25% s character). This is similar to why a lone pair is more stable on a more electronegative atom like fluorine than on a less electronegative atom like carbon.
Compared to alkynes, we expect the lone pairs in sp-hybridized nitriles to be the most stable and thus the least basic. We therefore expect that the lone pair in sp
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H = -10) shows that nitriles are indeed very weak bases. Similarly, we can find the lower base of pyridine (pK
(By the way, this is not a resonance! The lone pair in pyridine is in the plane of the ring, and therefore does not belong to the p orbitals).
It turns out that the nitrogen in pyrrole is unusually nonbasic. In fact, even when pyrrole is exposed to acid, it reacts at carbon (C-2), not nitrogen. Pyridine [Pk
Conjugate acid is nootropic. Removal of the lone pair on the nitrogen by protonation destroys the lone pair bonds with the other p orbitals of the ring and the molecule is non-aromatic.
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This might ring a bell. You may recall that cyclopentadiene is an unusually strong acid for a hydrocarbon. We can express this as “the unusually weak conjugate base of cyclopentadiene”.
The key is to pay attention to situations where the formation of new N-H bonds can lead to off-flavors. [Note 2]
Well, I said there are 5 factors that affect acidity, but I’ve explained six here. This is because “pi-donating” and “pi-accepting” behavior is not usually covered much in Org 1 (when acidity is introduced), but when most students encounter amines, they are familiar with these concepts in the field of electrophilic aromatic substitution. .
Today’s key lesson is: “The stronger the acid, the weaker the conjugate base” and “the weaker the acid, the stronger the conjugate base”, any factor that affects acidity is also a factor that affects basicity.
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If you understand factors that stabilize negative charge (and thus make atoms less basic), then by definition you also understand factors that destabilize negative charge (and thus make atoms more basic).
How to act in a situation where many factors play a role? We must resort to experimental measurement (pK
In the next post we’ll look at a specific example of this – and, I might add, a common exam problem. 🙂
Note 1. Since we are only discussing the basicity of nitrogen here, the atomic components of electronegativity and polarizability are not discussed.
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N(-) which is more basic than HO(-) which is more basic than F(-). The less electronegative the element, the more stable the lone pair and therefore the more basic it is.
Another useful trend is that as you go down a column of the periodic table, the base decreases. This is because the valence orbitals increase in size as you move down a column of the periodic table. Therefore, the electrons are “spread out” over a larger volume, resulting in a lower charge density. Our common way of describing this is to use the term “polarizability”.
Note 2. Here is an interesting example where nitrogen is unusually basic because it is aromatic. Tristan Lambert’s research group has developed a family of imine “superbases”.
Note that the significant resonance form of the conjugate acid is the substituted version of the aromatic cycloproponium cation. It helps balance the conjugate acid. pK Although every effort is made to follow the rules of citation style, there may be some inconsistencies. Please refer to the appropriate style manual or other resources if you have any questions.
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Polarity, in chemical bonds, the distribution of electric charge on the atoms connected by the bond. In particular, when bonding between similar atoms, such as H
, are electrically homogeneous in the sense that both hydrogen atoms are electrically neutral, bonds between atoms of different elements are electrically unbalanced. For example, in hydrogen chloride, the hydrogen atom is slightly positively charged while the chlorine atom is slightly negatively charged. The slight electric charge on different atoms is called partial charge and the presence of partial charge is the phenomenon of polar bonding.
The polarity of the bond comes from the relative electronegativity of the elements. Electronegativity is the ability of an atom of an element to attract electrons to itself
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