Lewis Acid and Bases

Lewis definition of acids and bases

In a Lewis acid-base reaction, the base will "attack" the acid and donate an electron pair. This causes a covalent bond to be formed between the two.This general reaction is shown below:

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Fig. 1:The Lewis base "attacks" the acid, forming a bond

Now let's talk about this reaction on an orbital level. The base's lone pair of electrons (non-bonded electrons), are in its highest occupied molecular orbital (called HOMO). The base is taking these electrons and interacting with the acid's lowest unoccupied molecular orbital (called LUMO). When they interact, they form a bond at a lower energy level, as shown below:

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Lewis acid and bases examples

Now that we know what the general reaction looks like, let's look at some examples:

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In the top example, the fluoride ion (F^-) is our base, which attacks the acid compound boron trifluoride (BF₃). After the reaction, a new B-F bond is formed (shown in red) to make the new compound boron tetrafluoride (BF₄^-).

In the bottom example, ammonia (NH₃) is our base and reacts with the acidic proton/hydrogen ion (H⁺). This forms a new N-H bond and the new compound ammonium (NH₄⁺)

In general, if a species has a negative charge, it will be a base. However, if a species has a positive charge, it will be an acid.

Coordination complexes

The Lewis theory of acids and bases is important, since it is able to explain the formation of coordination compounds, while other theories cannot.

Coordination complexes are formed from repeated acid base reactions, as shown below:

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Here, the cyanide ion (CN^-) attacks the positively charged zinc ion (Zn²⁺), which forms a bond between them. This reaction happens 4 times, for a total of 4 Zn-CN bonds. This is considered a complex ion since it is a charged coordination compound.

In general, the Lewis base will be your ligand(s), while your Lewis acid is the metal atom/ion.

Strength of Lewis acids and bases

Strength of Lewis acids

Lewis acids are electrophiles, so their strength is based on electrophilicity.

Electrophilicity is based on (mainly) two things:

1. Charge.
2. Electronegativity.

Let's look at these two concepts separately:

Charge

Electrophiles tend to have positive charges. A positive charge indicates a lack of electrons/electron density, therefore it wants strongly wants some electrons. Because of this, positively charged species are going to be more electrophilic than their neutral counterparts.

Generally speaking:

$$A^{+}>AX$$

Where A is a positively charged species and X is a negatively charges species

For example:

$$M{g}^{2+}>MgC{l}_2$$

Electronegativity

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The elements the closer to the top right (fluorine:F) of the periodic table are more electronegative, therefore, these elements are stronger electrophiles

For example, here is the trend in electophilicity for group 2:

$$\text{most electrophilic}B{e}^{2+}>M{g}^{2+}>C{a}^{2+}>S{r}^{2+}\text{less electrophilic}$$

Strength of Lewis bases

Lewis bases are nucleophiles, so their strength is based on nucleophilicity.

Nucleophilicity is based on four things

1. Charge
2. Electronegativity
3. Resonance/Charge localization
4. Steric hindrance

Let's break this down piece by piece

Charge

Nucleophiles tend to have negative charges. This is because negative charges indicate electron density (i.e. an excess of electrons). Negative charges tell us, "I have extra electrons and need to get rid of them!"

So in general:

$$A^{-}>AH$$

Where A^- is the negative species (conjugate base) of the acid AH.

For example, HS^- is more nucleophilic than H₂S.

Electronegativity

Since electronegativity is the tendency to attract electrons, nucleophiles are stronger when they are less nucleophilic. Basically, highly electronegative species "like" having a more negative charge/high electron density, so they don't want to "give up" their electrons as much.

For example, here is what this trend looks like for our halides (group 17):

$$\text{more nucleophilic}{I}^{-}>B{r}^{-}>C{l}^{-}>F^{-}\text{less nucleophilic}$$

Resonance/Charge localization

Species are more nucleophilic when the charge is localized, rather than delocalized, as shown below:

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When a charge is delocalized, as shown on the right, the charge is essentially being weakened by being spread out. Since the charge is "weaker", it is less reactive, and therefore less nucleophilic

Steric hinderance

"Sterics" relate to the spatial arrangement of atoms, so "steric hinderance" means "the way the atoms are arranged means they are in the way"

For example:

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Basically, when a nucleophile is "bulky" it makes the reaction slower since it's bulkiness is "in the way". Because of this, bulky=less nucleophilic

Lewis acid and bases practice problems

Now that we've covered a lot about Lewis acids and bases, let's work on some practice problems:

Now let's try a trickier one:

Using what you know about Lewis acid-base reactions, show the product of this reaction and label the Lewis acid and Lewis base

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Fig.8: Example problem

At first glance, it might be hard to determine what the carbon molecule is. As a hint, I drew the partial charge above the carbon atom. This molecule is a nucleophile (Lewis acid), because of carbon's partial positive charge.

The water molecule has two lone pairs, so it has electrons it can donate. This means it is the Lewis base.

Now for the product. Carbon can only have four bonds, so how can the base bond to it? Well, the electrons in one of the carbon-oxygen bonds are donated back to the oxygen, giving it a negative charge. Now that carbon has three bonds again, it can bond to the base and be completely neutral, as shown below:

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Fig.9: Worked out example