Lewis Acid and Bases

When you think of acids, what do you think of? Maybe you think of acidic foods like pineapple or vinegar, or maybe you think of the more technical definition, where a species donates a proton. Well, that definition of an acid is from the BrØnsted-Lowry theory. Today, we are going to be talking about a different theory: the Lewis acid and base theory. 

Lewis Acid and Bases Lewis Acid and Bases

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Contents
Table of contents

    In this article, we will be learning about Lewis acids and bases: what they are, some examples, and how to determine their strength.

    • This article covers Lewis acids and bases.
    • First, we will define what Lewis acids and bases are.
    • Next, we will look at some examples and learn how the Lewis acid-base theory explains how coordination complexes are formed.
    • After that, we will learn how to determine the strength of Lewis acids and bases.
    • Lastly, we will work on some practice problems.

    Lewis definition of acids and bases

    A Lewis acid is any species that accepts an electron pair. These species are also called electrophiles, since they are "electron loving" or "electron attracting"

    A Lewis base is any species that donates an electron pair. These species are also called nucleophiles, since they "nucleus loving" i.e, they want to be closer to their nucleus by getting rid of electrons

    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: Lewis Acid and Bases General Acid-Base Reaction StudySmarterFig. 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:

    Lewis Acid and Bases Lewis Acid-Base Bonding StudySmarterFig.2: The base's HOMO interacts with the acid's LUMO to form a bond

    Lewis acid and bases examples

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

    Lewis Acid and Bases Lewis Acid-Base reaction examples StudySmarterFig.3: Examples of Lewis acid-base reactions

    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.

    Think about what this charge is telling us. For example, positively charge species lack electrons, which means they are likely acids since they want more electrons.

    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 species where a metal ion is the center atom and other species (called ligands) are bonded to it.

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

    Lewis Acid and Bases Formation of a coordination complex StudySmarterFig.4: Formation of the coordination compound zinc tetracyanide

    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:

    $$Mg^{2+}>MgCl_2$$

    Electronegativity

    Electronegativity is measures the tendency of a species to attract/gain electrons.

    In other words, highly electronegative species want/can handle more electrons.Below is a table showing the electronegativity values for most elements:

    Lewis Acid and Bases Table of electronegativities StudySmarterFig.5: Table of electronegativities

    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:

    $$\text{most electrophilic}\,Be^{2+}>Mg^{2+}>Ca^{2+}>Sr^{2+}\,\text{less electrophilic}$$

    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):

    $$\text{more nucleophilic}\,I^->Br^->Cl^->F^-\,\text{less nucleophilic}$$

    Resonance/Charge localization

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

    Lewis Acid and Bases Resonance effect on nucleophilicity StudySmarterFig.6: Localized charges are more nucleophilic

    "R" is a stand in for any group that contains a carbon-hydrogen component.

    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:

    Lewis Acid and Bases Steric effect on nucleophilicity StudySmarterFig.7: More R-groups mean more "bulkiness"

    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:

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

    $$Cl^- + Ag^+ \rightarrow\,?$$

    Cl- is our Lewis base since it has a negative charge, while Ag+ is our Lewis acid since it has a positive charge. In a Lewis acid-base reaction, the base "attacks" the acid, so they form a bond, meaning the product of the reaction is AgCl.

    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

    Lewis Acid and Bases Lewis acid-base example problem StudySmarterFig.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:

    Lewis Acid and Bases Lewis acid-base example problem solved StudySmarterFig.9: Worked out example

    Lewis Acid and Bases - Key takeaways

    • A Lewis acid is any species that accepts an electron pair. These species are also called electrophiles, since they are "electron loving" or "electron attracting"
    • A Lewis base is any species that donates an electron pair. These species are also called nucleophiles, since they "nucleus loving" i.e, they want to be closer to their nucleus by getting rid of electrons
    • Coordination complexes are species where a metal ion is the center atom and other species (called ligands)are bonded to it
      • The Lewis acid is the metal center, while the Lewis base is the ligand(s)
    • The strength of a Lewis acid is based on:
      1. Charge: Positively charged species are stronger
      2. Electronegativity: More electronegative species are stronger
    • The strength of a Lewis base is based on:
      1. Charge: Negatively charged species are stronger
      2. Electronegativity: Less electronegative species are stronger
      3. Resonance/Charge localization: Localized charge is stronger than delocalized charge
      4. Steric hindrance: Bulky species are weaker

    References

    1. Fig.5-Table of electronegativities (https://upload.wikimedia.org/wikipedia/commons/thumb/4/42/Electronegative.jpg/640px-Electronegative.jpg) by ad blocker on Wikimedia commons licensed by CC BY-SA 3.0 (https://creativecommons.org/licenses/by-sa/3.0/)
    Frequently Asked Questions about Lewis Acid and Bases

    What is a Lewis acid and base, with examples?

    A Lewis acid is any species that accepts an electron pair. These species are also called electrophiles, since they are "electron loving" or "electron attracting" Ex: Mg2+


    A Lewis base is any species that donates an electron pair. These species are also called nucleophiles, since they "nucleus loving" i.e, they want to be closer to their nucleus by getting rid of electrons Ex: OH-



    How do you identify a Lewis acid and base?

    Lewis acids are usually positively charged

    Lewis bases are usually negatively charged and/or have a lone pair of electrons

    Are Lewis bases proton donors?

    No, Lewis bases donate electrons

    What is the difference between a Lewis acid and a Lewis base?

    Lewis acid: gains/needs electrons

    Lewis base: loses/donates electrons

    Which characteristic is associated with Lewis bases?

    a. A positive charge

    b. Double bonds

    c. High electronegativity

    d. Low electronegativity


    The answer is d, since Lewis bases are nucleophiles and have low electronegativity

    Test your knowledge with multiple choice flashcards

    What happens in a Lewis acid-base reaction?

    True or False: In general, Lewis acids have a positive charge

    In a coordination complex, the __ is a Lewis acid

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