Question

Boron halides behave as Lewis acids because they (a) Have only six electron in the valence shell (b) Have a lone pair of electrons on the \(\mathrm{B}\) atom (c) Are covalent compounds (d) Are ionic compounds

Short answer

(a) Boron halides behave as Lewis acids because they have only six electrons in the valence shell.

Step-by-step solution

Understanding the Behavior of Boron Halides

Boron halides are chemical compounds composed of boron (B) and a halogen (like fluorine or chlorine). To understand why they behave as Lewis acids, we need to analyze the electron configuration of the boron atom.

Examining Boron's Electron Configuration

Boron has an atomic number of 5, with an electron configuration of 1s² 2s² 2p¹. In its compounds, like boron halides (BX₃), boron forms three covalent bonds with halogens, sharing one electron with each. This results in boron having only six electrons in its valence shell, missing two to complete an octet.

Applying Lewis Acid Definition

Lewis acids are defined as chemical species that can accept an electron pair. Due to the electron deficiency (only having six electrons in its valence shell), boron in boron halides can accept electron pairs from other species, qualifying them as Lewis acids.

Evaluating the Options

Now that we understand the electron deficiency of boron in these compounds, we can evaluate the options: (a) Correct - Boron halides behave as Lewis acids because they have only six electrons in the valence shell, making them electron-deficient. (b) Incorrect - Boron does not have a lone pair of electrons in these compounds. (c) Irrelevant for Lewis acidity - Though they are covalent, this does not directly relate to their behavior as Lewis acids. (d) Incorrect - Boron halides are not ionic; they are covalent compounds.

Key concepts

Boron Halides

Boron halides are fascinating compounds formed by the combination of boron (B) with a halogen element, such as fluorine (F), chlorine (Cl), or bromine (Br). These compounds are known for their role as Lewis acids, which means they can accept a pair of electrons.
For instance, in boron trifluoride (\(\text{BF}_3\)), boron forms covalent bonds with three fluorine atoms. Each fluorine atom contributes a pair of electrons to form a bond. However, boron ends up with only six electrons in its outer shell, two fewer than the complete octet typically found in stable molecules.
This lack of electrons makes boron halides versatile in chemical reactions. They can interact with electron-rich species, accepting electron pairs to achieve stability. This unique property is key to their characterization as Lewis acids, allowing them to be crucial components in various industrial and laboratory applications.

Electron Deficient

The term "electron deficient" describes a situation where an atom has fewer than the usual eight electrons in its valence shell. This typically applies to certain elements that readily form compounds that can accept electrons from others.
In boron halides, the central boron atom is electron deficient. It forms three covalent bonds but has only six electrons around it, two short of the stable octet. This deficiency is why boron halides are potent Lewis acids—they are naturally predisposed to accept electron pairs to fill their outer shell.
Electron deficiency is a driving force behind the reactivity of these compounds, drawing them to engage with other molecules to achieve electron pair donation. This eagerness to accept electrons is why boron halides play a crucial role in numerous chemical synthesis processes, acting as catalysts or reactants.

Covalent Bonds

Covalent bonds are chemical bonds formed by the sharing of electron pairs between atoms. In boron halides, covalent bonds are created when boron shares its electrons with halogen atoms to form a stable compound.
These bonds are essential as they bind the atoms together to form a molecule that, despite being electron deficient, is stable enough to engage with other chemicals. For example, in boron trifluoride (\(\text{BF}_3\)), each bond involves the sharing of electrons between boron and fluorine atoms.
Although the bond formation fills up part of boron's valence shell, it leaves the molecule open to further interactions due to its electron-deficient nature. Thus, the covalent bonds not only secure the structural integrity of boron halides but also influence their behavior as Lewis acids. They allow these compounds to interact flexibly with electron donors, making them integral to many chemical reactions.