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Chapter 16: Problem 78

In terms of orbitals and electron arrangements, what must be present for a molecule or an ion to act as a Lewis acid (use \(\mathrm{H}^{+}\) and \(\mathrm{BF}_{3}\) as examples)? What must be present for a molecule or ion to act as a Lewis base (use \(\mathrm{OH}^{-}\) and \(\mathrm{NH}_{3}\) as examples)?

Short Answer

Expert verified
Lewis acids accept electron pairs (e.g., \(\mathrm{H}^{+}, \mathrm{BF}_{3}\)), while Lewis bases donate them (e.g., \(\mathrm{OH}^{-}, \mathrm{NH}_{3}\)).

Step by step solution

01

Understanding Lewis Acids

Lewis Acids are molecules or ions that can accept a pair of electrons to form a coordinate covalent bond. This generally occurs because they have an empty orbital or are electron-deficient. Examples include \(\mathrm{H}^{+}\), which has no electrons and an empty orbital, allowing it to accept electron pairs. Additionally, \(\mathrm{BF}_{3}\) acts as a Lewis acid because the Boron atom has an empty p-orbital, facilitating its acceptance of an electron pair.
02

Identifying Lewis Bases

Lewis Bases are molecules or ions that donate an electron pair to form a coordinate covalent bond. A key characteristic is having at least one lone pair of electrons. For instance, \(\mathrm{OH}^{-}\) has lone pairs on the oxygen atom, making it capable of donating electron pairs. Similarly, \(\mathrm{NH}_{3}\) has a lone pair on the nitrogen atom, which can be donated to a Lewis acid.
03

Contrasting Electron Configurations

For a Lewis acid like \(\mathrm{H}^{+}\), the complete lack of electrons leaves an open path for accepting electrons. In contrast, \(\mathrm{BF}_{3}\) has filled orbitals around Boron but still an electron-deficient central atom seeking more electrons. For Lewis bases like \(\mathrm{OH}^{-}\) and \(\mathrm{NH}_{3}\), the presence of lone pairs, or non-bonding electrons, is crucial as they are donated to Lewis acids.
04

Relating Examples to Definitions

By evaluating \(\mathrm{H}^{+}\) and \(\mathrm{BF}_{3}\), we see that Lewis acids all have vacant orbitals or an acceptor nature due to their electron deficiency. Similarly, \(\mathrm{OH}^{-}\) and \(\mathrm{NH}_{3}\) illustrate that Lewis bases possess donor tendencies derived from lone electron pairs.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Electron Configuration
Electron configuration is the arrangement of electrons within the orbitals of an atom or molecule. This concept is key to understanding why certain substances act as Lewis acids or bases.

To act as a Lewis acid, a molecule or ion must possess empty orbitals. These empty orbitals are often the result of being electron-deficient. For instance,
  • the hydrogen ion (\(\mathrm{H}^{+}\) ) lacks electrons altogether, hence, it has an empty electron configuration that makes it eager to accept electrons from a donor.
  • Similarly, boron trifluoride (\(\mathrm{BF}_{3}\)) contains an empty p-orbital which allows it to accept electron pairs. Although it has a complete octet when considering covalent bonds, it remains electron-deficient because the boron atom can expand its octet by accepting electrons.
For Lewis bases, electron configuration must include lone pairs, which are pairs of valence electrons that are not involved in bonding.

Examples include:
  • Hydroxide ion (\(\mathrm{OH}^{-}\) ), with its lone pairs on the oxygen, allowing it to donate an electron pair to Lewis acids.
  • Ammonia (\(\mathrm{NH}_{3}\)) has a lone pair on nitrogen, making it a potential electron donor.
Understanding electron configurations helps clarify why some species can act as either acceptors or donors of electrons.
Lewis Base
A Lewis base is defined by its ability to donate an electron pair. This ability hinges on the presence of lone pairs of electrons鈥攗nshared pairs that are not involved in bonding interactions.

When a molecule or ion donates an electron pair, it forms a new bond with a Lewis acid, which is a molecule or ion eager to accept electrons. This donation results in the formation of a coordinate covalent bond.

Classic examples include:
  • The hydroxide ion (\(\mathrm{OH}^{-}\)), which has two lone pairs available for donation. These non-binding electrons on the oxygen atom can easily be shared with a Lewis acid, demonstrating its basic nature.
  • Ammonia (\(\mathrm{NH}_{3}\)), with a lone pair on the nitrogen. This pair is not shared with any other atom and can easily be offered to form a new bond with a Lewis acid.
The willingness and readiness of these molecules and ions to share their lone pair afford them their character as Lewis bases. It is the availability of these lone electron pairs that fundamentally defines the strength and reactivity of a Lewis base.
Coordinate Covalent Bond
A coordinate covalent bond is a specific type of covalent bond where both electrons in the shared pair come from one atom, typically a Lewis base. This is different from a typical covalent bond where each atom contributes one electron to the bond.

In the interaction between a Lewis acid and a Lewis base, the Lewis base donates its lone pair of electrons to the electron-deficient Lewis acid to form a coordinate covalent bond.

Two components are necessary for this bond to form:
  • A Lewis base that has a lone pair of electrons ready for donation.
  • A Lewis acid with an empty orbital or the capacity to accept an electron pair.
For instance,
  • The bond formed between ammonia (\(\mathrm{NH}_{3}\)) and a hydrogen ion (\(\mathrm{H}^{+}\)) results in the formation of the ammonium ion (\(\mathrm{NH}_{4}^{+}\)). All bonding electrons originated from the nitrogen atom in ammonia, characterizing it as a coordinate bond.
  • Similarly, when the boron trifluoride (\(\mathrm{BF}_{3}\)) accepts an electron pair from hydroxide ion (\(\mathrm{OH}^{-}\)), a coordinate covalent bond results.
This form of bonding is essential for various chemical reactions and complexes, showcasing the versatility and adaptability of atoms to bond under different circumstances.

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Most popular questions from this chapter

How does the strength of an oxoacid depend on the electronegativity and oxidation number of the central atom?

List the factors on which the \(K_{\mathrm{a}}\) of a weak acid depends.

The \(\mathrm{pH}\) of a \(0.0642 \mathrm{M}\) solution of a monoprotic acid is 3.86. Is this a strong acid?

The ion-product constant for water is \(1.0 \times 10^{-14}\) at \(25^{\circ} \mathrm{C}\) and \(3.8 \times 10^{-14}\) at \(40^{\circ} \mathrm{C}\). Is the process $$\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{OH}^{-}(a q)$$ endothermic or exothermic?

What does the ionization constant tell us about the strength of an acid?
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