Question

What do you understand by (i) electron-deficient. (ii) electron-precise, and (iii) electron-rich compounds of hydrogen? Provide justification with suitable examples.

Short answer

Electron-deficient have incomplete octets, electron-precise have exact electrons for full valence, and electron-rich have extra electrons possibly forming lone pairs.

Step-by-step solution

Define Electron-Deficient Compounds

Electron-deficient compounds are those which do not have a complete set of electrons to form the stable octet structure around their atoms. This often leads to these compounds forming bonds by sharing electrons, such as in boranes, like diborane (B2H6), where each boron makes use of additional hydrogen bridges to achieve a stable structure.

Define Electron-Precise Compounds

Electron-precise compounds are those that have just enough electrons to form all the covalent bonds necessary for maintaining the standard valency of each atom. In these compounds, there are no surplus or deficit of electrons. An example can be methane (CH4), where each hydrogen shares one electron with carbon, leading to a full valence shell for carbon and hydrogen.

Define Electron-Rich Compounds

Electron-rich compounds contain more electrons than needed for their standard valency bonds, often resulting in the presence of lone pairs or double bonds. Ammonia (NH3) is a classic example, where the nitrogen has a lone pair of electrons, which can participate in hydrogen bonding or act as a Lewis base.

Key concepts

Electron-Precise Compounds

Electron-precise compounds are substances where each atom has the exact number of electrons to form covalent bonds that satisfy its standard valency. In simpler terms, every atom in these compounds uses its valence electrons to make the necessary connections for a stable molecular structure, with no surplus or deficit.

For instance, consider methane (\( ext{CH}_4\)). Here, carbon uses its four valence electrons to form four covalent bonds with hydrogen atoms, each hydrogen providing one electron. Together, they create a stable, electron-precise environment. No excess electrons or lone pairs are left, making it a very balanced form of molecular structure.

Electron-Rich Compounds

Electron-rich compounds have more electrons than needed for forming covalent bonds according to their standard valency. This surplus often results in the presence of lone pairs or enables the possibility of forming multiple bonds such as double or triple bonds.

A classic example is ammonia (\( ext{NH}_3\)). In this molecule, nitrogen forms three covalent bonds with hydrogen and retains a pair of non-bonding electrons, known as a lone pair. This lone pair gives the compound its electron-rich nature and provides flexibility in chemical reactivity, such as acting as a Lewis base or engaging in hydrogen bonding.

Covalent Bonds

Covalent bonds are the primary force holding atoms together in many compounds. They occur when two atoms share one or more pairs of valence electrons, allowing each atom to attain a stable electron configuration. By sharing electrons, atoms can fulfill their valency requirements.

An important aspect of covalent bonding is that the shared electrons "belong" to both bonded atoms, resulting in a mutual dependence for maintaining the bond's stability. This sharing is often equal in bonds between identical atoms (like in diatomic gases), while bonds between different atoms may have unequal sharing, leading to polar covalent bonds.

Valency

Valency is a concept that refers to the ability of an atom to form bonds with other atoms. It describes the number of electrons an atom can lose, gain, or share in order to achieve a full outer shell, resulting in a stable configuration.

For example, nitrogen typically has a valency of three, meaning it can form three bonds with other atoms, as seen in ammonia (\( ext{NH}_3\)). In contrast, carbon has a valency of four, allowing it to form four bonds, which is evident in methane (\( ext{CH}_4\)). Valency dictates the structural possibilities of molecules and directly influences chemical reactivity and bonding patterns.

Lone Pairs

Lone pairs consist of valence electrons that are not involved in covalent bonding. These non-bonding pairs of electrons reside with a single atom and can significantly influence the geometry and reactivity of a molecule.

In ammonia (\( ext{NH}_3\)), nitrogen's lone pair not only affects the molecule's shape by creating a trigonal pyramidal structure, but it also is influential in chemical interactions. For example, the lone pair can readily bond with a hydrogen ion (\( ext{H}^+\)), transforming ammonia into an ammonium ion (\( ext{NH}_4^+\)). Lone pairs can thus play a critical role in determining a molecule's properties and behavior in chemical reactions.