Problem 6
Question
Which of the following are Lewis acids? (a) \(\mathrm{PH}_{3}\) and \(\mathrm{BCl}_{3}\) (b) \(\mathrm{AlCl}_{3}\) and \(\mathrm{SiCl}_{4}\) (c) \(\mathrm{PH}_{3}\) and \(\mathrm{SiCl}_{4}\) (d) \(\mathrm{BCl}_{3}\) and \(\mathrm{AlCl}_{3}\)
Step-by-Step Solution
Verified Answer
Options (a), (b), and (d) contain Lewis acids.
1Step 1: Understand the Definition of a Lewis Acid
A Lewis acid is a chemical species that can accept a pair of electrons from a Lewis base to form a coordinate covalent bond. Typically, these are atoms or molecules with empty orbital, such as certain metal cations or electron-deficient compounds.
2Step 2: Analyze Each Compound for Electron Deficiency
We'll examine each compound for its ability to accept electron pairs. A Lewis acid often has an incomplete octet or can expand its valency: - (a) \(\text{PH}_3\) has a full octet (not a Lewis acid), \(\text{BCl}_3\) has an empty p orbital (Lewis acid).- (b) \(\text{AlCl}_3\) is electron-deficient with a vacant p orbital (Lewis acid), \(\text{SiCl}_4\) has a complete octet and cannot accept more electrons (not a Lewis acid).- (c) \(\text{PH}_3\) has a full octet (not a Lewis acid), \(\text{SiCl}_4\) also has a full octet (not a Lewis acid).- (d) \(\text{BCl}_3\) and \(\text{AlCl}_3\) both have vacant orbitals and are electron-deficient (both are Lewis acids).
3Step 3: Identify Options Containing Lewis Acids
Based on the analysis:- Option (a) includes one Lewis acid, \(\text{BCl}_3\).- Option (b) contains one Lewis acid, \(\text{AlCl}_3\).- Option (c) does not contain any Lewis acids.- Option (d) contains two Lewis acids, \(\text{BCl}_3\) and \(\text{AlCl}_3\).
Key Concepts
Electron DeficiencyCoordinate Covalent BondChemical Species with Empty Orbitals
Electron Deficiency
Electron deficiency is a key concept in understanding Lewis acids. A chemical species is considered electron-deficient when it does not have enough electrons to complete its octet, making it eager to accept electrons from other species. This condition often leads to the formation of compounds that can act as Lewis acids.
For example, in the case of boron trichloride \(\text{BCl}_3\), boron is electron-deficient. Boron has only three electrons in its outer shell, and although it shares three pairs with chlorine atoms, it still lacks enough electrons to fulfill the octet rule. This "hunger" for electrons makes \(\text{BCl}_3\) a classic example of an electron-deficient Lewis acid.
Such electron-deficient compounds are typically found in either group 13 elements or transition metals with incomplete d or f orbitals. Notably, these deficiencies are critical because they highlight the potential of these species to interact and form bonds with Lewis bases that supply the necessary electron pairs.
For example, in the case of boron trichloride \(\text{BCl}_3\), boron is electron-deficient. Boron has only three electrons in its outer shell, and although it shares three pairs with chlorine atoms, it still lacks enough electrons to fulfill the octet rule. This "hunger" for electrons makes \(\text{BCl}_3\) a classic example of an electron-deficient Lewis acid.
Such electron-deficient compounds are typically found in either group 13 elements or transition metals with incomplete d or f orbitals. Notably, these deficiencies are critical because they highlight the potential of these species to interact and form bonds with Lewis bases that supply the necessary electron pairs.
Coordinate Covalent Bond
A coordinate covalent bond, also known as a dative bond, is crucial in the formation of a complex between a Lewis acid and a Lewis base. In this type of bond, both electrons in the shared pair originate from the same atom—typically the Lewis base.
Take the reaction between \(\text{BCl}_3\) (a Lewis acid) and ammonia \(\text{NH}_3\) (a Lewis base) as an example. Ammonia has a lone pair of electrons on the nitrogen atom. In the formation of the bond, this lone pair is donated to the empty orbital of the boron atom in \(\text{BCl}_3\).
This process effectively fills the electron-deficient boron, forming a stable complex. The fascinating aspect of coordinate covalent bonds is how they allow electron-deficient compounds to stabilize through interactions with electron-rich species.
Take the reaction between \(\text{BCl}_3\) (a Lewis acid) and ammonia \(\text{NH}_3\) (a Lewis base) as an example. Ammonia has a lone pair of electrons on the nitrogen atom. In the formation of the bond, this lone pair is donated to the empty orbital of the boron atom in \(\text{BCl}_3\).
This process effectively fills the electron-deficient boron, forming a stable complex. The fascinating aspect of coordinate covalent bonds is how they allow electron-deficient compounds to stabilize through interactions with electron-rich species.
Chemical Species with Empty Orbitals
Chemical species that can act as Lewis acids often have empty orbitals available for bonding. These empty orbitals provide sites where electron pairs from Lewis bases can be accepted. This adaptability of orbitals in chemical species is what distinguishes them as potential Lewis acids.
For instance, aluminum chloride \(\text{AlCl}_3\) is a prime example in which the aluminum atom has an empty p-orbital. This vacant p-orbital can accept electron pairs from donor molecules, aiding in the formation of coordinate covalent bonds.
Having empty orbitals makes these species versatile in chemical reactions, allowing them to participate in a variety of interactions that are foundational in fields like coordination chemistry. Hence, recognizing species with empty orbitals is essential for predicting their behavior as Lewis acids in diverse chemical environments.
For instance, aluminum chloride \(\text{AlCl}_3\) is a prime example in which the aluminum atom has an empty p-orbital. This vacant p-orbital can accept electron pairs from donor molecules, aiding in the formation of coordinate covalent bonds.
Having empty orbitals makes these species versatile in chemical reactions, allowing them to participate in a variety of interactions that are foundational in fields like coordination chemistry. Hence, recognizing species with empty orbitals is essential for predicting their behavior as Lewis acids in diverse chemical environments.
Other exercises in this chapter
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