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Magazine Chapter 4: Problem 32
For each equation, label the Lewis acid and the Lewis base. In addition, show all unshared pairs of electrons on the reacting atoms, and use curved arrows to show the flow of electrons in each reaction. (a) \(\mathrm{F}^{-}+\mathrm{BF}_{3} \longrightarrow \mathrm{BF}_{4}^{-}\) (b)
Short Answer
Expert verified
Answer:
The Lewis acid in this reaction is \(\mathrm{BF}_{3}\), and the Lewis base is \(\mathrm{F}^{-}\).
Step by step solution
01
Identify the Lewis Acids and Bases
To identify the Lewis acids and bases, we have to look for electron acceptors (Lewis acids) and donors (Lewis bases). In the given reaction, \(\mathrm{F}^{-}\) donates an electron pair, making it the Lewis base, while \(\mathrm{BF}_{3}\) accepts an electron pair, making it the Lewis acid. So we have:
Lewis Acid: \(\mathrm{BF}_{3}\)
Lewis Base: \(\mathrm{F}^{-}\)
02
Show Unshared Pairs of Electrons
Now let's show all unshared pairs of electrons on the reacting atoms:
$$\mathrm{F}^{-}:\quad \quad \quad \quad \quad \mathrm{BF}_{3}$$
On the fluoride ion, there are three lone pairs of electrons and an additional negative charge.
03
Use Curved Arrows to Show Electron Flow
Now, we will use curved arrows to show the flow of electrons in the reaction. The electron pair from the fluoride ion (Lewis base) is donated to the boron atom in \(\mathrm{BF}_{3}\) (Lewis acid), resulting in the formation of the tetrafluoroborate ion:
$$\mathrm{F}^{-}\longrightarrow\mathrm{BF}_{3} \longrightarrow \mathrm{BF}_{4}^{-}$$
with a curved arrow from the fluoride ion to the boron atom:
$$\small{\underset{\oplus}{\mathrm{B}}(\underset{\ominus}{\mathrm{F}})(\underset{\ominus}{\mathrm{F}})(\underset{\ominus}{\mathrm{F}})}\quad\longleftarrow\quad \small{\mathrm{F}}^{-}$$
04
Final Complete Reaction
Now we can show the complete reaction with the Lewis acid, Lewis base, electron pairs, and the electron flow:
$$\small{\mathrm{F}}^{-}\quad\longleftarrow\quad\small{\underset{\oplus}{\mathrm{B}}(\underset{\ominus}{\mathrm{F}})(\underset{\ominus}{\mathrm{F}})(\underset{\ominus}{\mathrm{F}})}\longrightarrow \mathrm{BF}_{4}^{-}$$
This completes the step-by-step solution for the given exercise:
(a) \(\mathrm{F}^{-}+\mathrm{BF}_{3} \longrightarrow \mathrm{BF}_{4}^{-}\)
Lewis Acid: \(\mathrm{BF}_{3}\)
Lewis Base: \(\mathrm{F}^{-}\)
Unshared pairs of electrons on reacting atoms
$$\small{\mathrm{F}}^{-}\quad\longleftarrow\quad\small{\underset{\oplus}{\mathrm{B}}(\underset{\ominus}{\mathrm{F}})(\underset{\ominus}{\mathrm{F}})(\underset{\ominus}{\mathrm{F}})}$$
$$\mathrm{F}^{-}+\mathrm{BF}_{3} \longrightarrow \mathrm{BF}_{4}^{-}$$
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Lewis acid-base reaction
Understanding the intricacies of Lewis acid-base reactions provides clarity to many chemical processes. These reactions are fundamental to chemistry and involve the transfer of electron pairs between molecules.
A Lewis acid is a molecule or ion that is an electron pair acceptor. On the other hand, a Lewis base is an electron pair donor. Lewis acid-base reactions can be seen as a kind of 'dance' where bases donate an electron pair to acids, leading to compound formation or molecular transformation.
For instance, in the reaction where fluoride ion (\texttt{F}\(^{-}\)) combines with boron trifluoride (\texttt{BF}\(_{3}\)), the fluoride ion, which has electron pairs that are not shared with other atoms, acts as the Lewis base. Conversely, \texttt{BF}\(_{3}\), which is electron deficient because boron has only three valence electrons, acts as the Lewis acid. The pair bonds, forming the tetrafluoroborate ion (\texttt{BF}\(_{4}^{-}\)), showcasing a classic Lewis acid-base interaction. This interaction highlights the importance of identifying which species in a reaction can donate an electron pair (the base) and which can accept it (the acid).
A Lewis acid is a molecule or ion that is an electron pair acceptor. On the other hand, a Lewis base is an electron pair donor. Lewis acid-base reactions can be seen as a kind of 'dance' where bases donate an electron pair to acids, leading to compound formation or molecular transformation.
For instance, in the reaction where fluoride ion (\texttt{F}\(^{-}\)) combines with boron trifluoride (\texttt{BF}\(_{3}\)), the fluoride ion, which has electron pairs that are not shared with other atoms, acts as the Lewis base. Conversely, \texttt{BF}\(_{3}\), which is electron deficient because boron has only three valence electrons, acts as the Lewis acid. The pair bonds, forming the tetrafluoroborate ion (\texttt{BF}\(_{4}^{-}\)), showcasing a classic Lewis acid-base interaction. This interaction highlights the importance of identifying which species in a reaction can donate an electron pair (the base) and which can accept it (the acid).
Curved arrow notation in organic chemistry
To effectively communicate the electron flow during chemical reactions, chemists use curved arrow notation. This notation is pivotal in visually representing the movement of electron pairs in a reaction mechanism.
A curved arrow, starting from the electron pair donor (Lewis base), points toward the electron pair acceptor (Lewis acid). This helps illustrate how bonds are broken and formed during the reaction. It's essential to draw these arrows correctly; the tail of the arrow represents the origin of the electron pair movement, while the head signifies the destination.
In our fluoride and boron trifluoride example, the lone pairs on the fluoride ion (\texttt{F}\(^{-}\)) are represented with a pair of dots, and the arrow is drawn from one of these pairs to the boron atom in \texttt{BF}\(_{3}\). By following the curved arrows, one can easily trace the path electrons take to form the new \texttt{BF}\(_{4}^{-}\) ion. This visual tool is invaluable for students and chemists to understand and predict the outcome of chemical reactions.
A curved arrow, starting from the electron pair donor (Lewis base), points toward the electron pair acceptor (Lewis acid). This helps illustrate how bonds are broken and formed during the reaction. It's essential to draw these arrows correctly; the tail of the arrow represents the origin of the electron pair movement, while the head signifies the destination.
In our fluoride and boron trifluoride example, the lone pairs on the fluoride ion (\texttt{F}\(^{-}\)) are represented with a pair of dots, and the arrow is drawn from one of these pairs to the boron atom in \texttt{BF}\(_{3}\). By following the curved arrows, one can easily trace the path electrons take to form the new \texttt{BF}\(_{4}^{-}\) ion. This visual tool is invaluable for students and chemists to understand and predict the outcome of chemical reactions.
Electron pair donors and acceptors
In the context of Lewis acid-base chemistry, molecules or ions can be classified as electron pair donors and electron pair acceptors. A deeper grasp of this concept aids in predicting reactivity and understanding molecular interactions.
An electron pair donor, also known as a Lewis base, typically possesses lone pairs of electrons that are available for bonding. These electron-rich species are keen to share their bounty and can form a coordinate covalent bond with an electron acceptor. Lone pairs are shown as pairs of dots in chemical diagrams, representing the available electrons for donation.
Electron pair acceptors, or Lewis acids, are electron-poor species such as ions or molecules with incomplete valence shells. They are 'hungry' for electrons and readily accept them from donors to form stable chemical bonds. In our example, \texttt{BF}\(_{3}\) is eagerly accepting an electron pair from \texttt{F}\(^{-}\), resulting in a bond that completes its valence shell, demonstrating their roles as electron acceptor and donor, respectively.
An electron pair donor, also known as a Lewis base, typically possesses lone pairs of electrons that are available for bonding. These electron-rich species are keen to share their bounty and can form a coordinate covalent bond with an electron acceptor. Lone pairs are shown as pairs of dots in chemical diagrams, representing the available electrons for donation.
Electron pair acceptors, or Lewis acids, are electron-poor species such as ions or molecules with incomplete valence shells. They are 'hungry' for electrons and readily accept them from donors to form stable chemical bonds. In our example, \texttt{BF}\(_{3}\) is eagerly accepting an electron pair from \texttt{F}\(^{-}\), resulting in a bond that completes its valence shell, demonstrating their roles as electron acceptor and donor, respectively.
