Name: electron-pair reactions in lewis theory
Uploaded: 2023-04-30T18:28:34.000Z
Duration: 2 min 16 s
Channel: JoVE

In the periodic table, all elements in group 13 have three valence electrons. Consequently, they form trivalent compounds in which the central atom has a sextet of electrons and a vacant orbital.
These electron-deficient compounds achieve their octets by receiving electrons from other species in a chemical reaction.
For example, aluminum chloride reacts with ammonia by accepting the lone pair from the nitrogen atom.
This makes aluminum chloride &#8212; the electron-pair acceptor &#8212; a Lewis acid, and ammonia &#8212; the electron-pair donor &#8212; a Lewis base.
As the two oppositely charged species attract each other, the electron pair is transferred from the base to the acid. This completes the octet of the Lewis acid, resulting in a Lewis acid&#8212;base adduct.
Now, consider another Lewis acid&#8212;base reaction between the electron-deficient boron trifluoride and ammonia.
The electrostatic potential map of boron trifluoride indicates a significant positive charge centered on the boron atom.
In the case of ammonia, the nitrogen atom carries a negative charge on account of its nonbonding electron pair.
The opposite charges cause boron trifluoride and ammonia to react. During the reaction, the lone pair of ammo

lewis acids and bases

Lewis Acids and Bases

This lesson delves into Lewis acids and bases in the context of the octet rule for electron-deficient compounds. Here, the concept is discussed, emphasizing the group 13 elements like boron or aluminium. Since group 13 elements possess three valence electrons, they form trivalent compounds with a sextet of electrons and a vacant orbital for the central atom. Consequently, these electron-deficient compounds accept electrons from other species to complete their octet in a chemical reaction. They are referred to as Lewis acids per the 'generalized theory of acids and bases' proposed by Gilbert N. Lewis.
Lewis's theory dealt with compounds that were not under the purview of Br&#248;nsted's definition. He proposed that the electron-deficient compounds act as a Lewis acid where their valence-shell octets are completed in a chemical reaction. Hence, a Lewis acid is the species that accepts a pair of electrons to form a new bond.
In contrast, a Lewis base is defined as the species that donates an electron pair. This is elucidated using the specific example of aluminum chloride reacting with ammonia to form a Lewis acid-base adduct. Here, the electron pair is transferred between the oppositely charged species to satisfy the octet. The Lewis acids and bases concept is further reiterated by the reaction between electron-deficient boron trifluoride and an electron-rich ammonia, as shown in Figure 1.
<img alt="Lewis acid-base reaction diagram; BF3 + NH3 forming adduct with electron pair donation." src="/files/ftp_upload/11734/11734_PT_Figure%201.svg" style="text-align: center; height: 90px;" data-altupdatedat="1761830439000" data-alt="Figure1" />
Figure 1. The reaction between Boron trifluoride and ammonia
Here, a significant charge is developed on the species. As the boron center has an empty orbital that could accept an electron, it localizes a positive charge. In contrast, the nitrogen center in ammonia accumulates a negative charge due to the presence of a lone pair of electrons. Hence when they interact, the lone pair of electrons in the valence shell of nitrogen is transferred to the boron atom in BF3, indicated by the curved arrow above. Thus, the formal charges on boron and nitrogen are balanced, and as a result, the Lewis acid-base adduct possesses no net charge.
The Lewis theory provides an addendum to the Br&#248;nsted theory that only uses proton transfer to define acid-base reactions by incorporating the transfer of a lone pair. Therefore, while all Br&#248;nsted&#8211;Lowry acids are protic acids, Lewis acids can be protic or aprotic. This is delineated using the example of hydrochloric acid. HCl is an acid, according to the Br&#248;nsted-Lowry definition, given its ability to donate a proton. It is also a Lewis acid since its hydrogen atom loses the shared electrons to chlorine while simultaneously accepting the pair of electrons from ammonia.

This lesson delves into Lewis acids and bases in the context of the octet rule for electron-deficient compounds. Here, the concept is discussed, ...

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Organic Chemistry

Unit 1

Acids and Bases

KEY TERMS AND CONCEPTS

electron-pair reactions in lewis theory

Lewis theory explains acid-base reactions through electron pairs, especially for electron-deficient compounds that do not fit the Brønsted definition. It is often used to describe group 13 elements such as boron and aluminium. These elements have three valence electrons, so their trivalent compounds can have a sextet of electrons and an empty orbital at the central atom.

Because of that electron shortage, these compounds can accept electrons from another species to complete their octet. Gilbert N. Lewis described this behavior in his generalized acid-base theory. In this model, a Lewis acid is the species that accepts a pair of electrons to form a new bond. A Lewis base is the species that donates an electron pair.

The reaction between aluminium chloride and ammonia shows this idea clearly. The electron pair moves from ammonia to the electron-deficient compound, and an acid-base adduct forms. The same pattern appears in the reaction between boron trifluoride and ammonia.

In BF3, the boron center has an empty orbital that can accept electrons, so it carries a positive charge. In ammonia, the nitrogen center has a lone pair of electrons and carries a negative charge. When they interact, the lone pair on nitrogen is transferred to boron, as shown by the curved arrow, and the formal charges are balanced so the adduct has no net charge.

Lewis theory also expands the Brønsted model, which defines acid-base reactions by proton transfer only. This means Lewis acids can be protic or aprotic. Hydrochloric acid is one example, because it donates a proton in the Brønsted-Lowry sense and also acts as a Lewis acid when its hydrogen accepts the pair of electrons from ammonia.

Electron-Pair Reactions in Lewis Theory

br&#248;nsted-lowry acids and bases

Proton Transfer in Brønsted-Lowry Reactions

Brønsted-Lowry acids and bases are defined by proton transfer in a chemical reaction. In this model, a Brønsted acid donates a proton and becomes its conjugate base. A Brønsted base accepts a proton and becomes its conjugate acid. This theory was proposed in 1923 by Johannes Brønsted and Thomas Lowry. It helps explain acid-base reactions by tracking how protons move between species. The electron shifts that drive these transfers are shown with curved arrows. A reaction between acetic acid and...

Lewis theory explains acid-base reactions through electron pairs, especially for electron-deficient compounds that do not fit the Brønsted definition. It is often used to describe group 13 elements such as boron and aluminium. These elements have three valence electrons, so their trivalent compounds can have a sextet of electrons and an empty orbital at the central atom. Because of that electron shortage, these compounds can accept electrons from another species to complete their octet.

acid and bases: ka, pka, and relative strengths

acid-and-bases-ka-pka-and-relative-strengths

Comparing Acid and Base Strength with pKa

Acid and base strength can be compared using Ka and pKa. These values help show how much an acid dissociates in water and how strong a base is through its conjugate acid. Acid strength depends on dissociation into a conjugate base and hydronium ion in water. A strong acid dissociates almost completely, so the hydronium ion concentration is very high. A weak acid dissociates only partly, so more acid remains undissociated and the equilibrium stays farther back. For dilute weak acids, the...

position of equilibrium in acid-base reactions

position-of-equilibrium-in-acid-base-reactions

Predicting Acid-Base Equilibrium Direction

Acid-base equilibrium direction depends on acid strength and pKa values. A lower pKa means a stronger acid, while a higher pKa means a weaker acid. A negative pKa shows a stronger acid, and a positive pKa shows a weaker acid. In the reaction between ammonia and an ethoxide ion, ethanol has a pKa of 15.9 and ammonia has a pKa of 38. That means ethanol is the stronger acid in this pair. A strong acid forms a weak conjugate base, and a weak acid forms a strong conjugate base, so the ethoxide ion...

molecular structure and acidity

What Makes One Acid Stronger?

Acid strength depends on how stable the conjugate base is after the acid loses a proton. When deprotonation forms a more stable anion, the acid is stronger. If the anion is unstable, the acid is weaker. Chemists compare several structure-based factors to judge that stability. One factor is size. When acids come from elements in the same column of the periodic table, atomic size is compared. A larger atom can spread out a negative charge over a bigger space. That makes the conjugate base more...

leveling effect and non-aqueous acid-base solutions

leveling-effect-and-non-aqueous-acid-base-solutions

Solvent Effects on Strong Acids and Bases

Solvent effects can change how strong acids and bases react in a chemical system. The leveling effect is one important example. It shows that a solvent can take part in the reaction before the acid and base react with each other. For a generic acid-base reaction, a generic acid (HA) reacts with a generic base (B−) to form the conjugate base (A−) and the conjugate acid (HB). In a solvent called HX, the solvent may also react if it is stronger than one of the reactants. If HA is weaker than the...

solvating effects

How Solvation Changes Acidity

Solvation changes acidity by affecting how stable a conjugate base is. The strength of an acid is commonly compared with its pKa value. When an ion dissolves, solvent molecules surround it and help stabilize it through solvation. Solvation of ions can happen in three main ways. In donor interaction, the solvent donates unshared electron pairs to the ion. The solvent acts as a Lewis base, and the ion acts as a Lewis acid. In charge-dipole interaction, polar solvent molecules align their partial...

key terms and concepts

Explore Lewis theory through electron-pair transfer, electron-deficient compounds, and examples like BF3, AlCl3, ammonia, and hydrochloric acid.

Electron-Pair Reactions in Lewis Theory | JoVE K12

organic chemistry

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Electron-Pair Reactions in Lewis Theory

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