Question

Among the following compound which one has maximum number of lone pairs of electrons on central atom (a) \(\left[\mathrm{ClO}_{3}\right]\) (b) \(\mathrm{XeF}_{4}\) (c) \(\mathrm{SF}_{4}\) (d) \(\left[I_{3}\right]^{-}\)

Short answer

\(\left[I_{3}\right]^{-}\) has the maximum number of lone pairs (3 lone pairs).

Step-by-step solution

Identify Central Atoms

Each compound has a central atom: Cl in \(\left[\mathrm{ClO}_{3}\right]\), Xe in \(\mathrm{XeF}_{4}\), S in \(\mathrm{SF}_{4}\), and I in \(\left[I_{3}\right]^{-}\). We will determine the number of lone pairs for each central atom.

Determine Lone Pairs for \(\left[\mathrm{ClO}_{3}\right]\)

For \(\left[\mathrm{ClO}_{3}\right]\), Cl starts with 7 valence electrons. It forms three single bonds with O atoms and, since it is negatively charged, it gains an electron, adding up to four bonds and allocating any remaining electrons as lone pairs. However, after bonding, Cl doesn't have any electrons left for lone pairs.

Determine Lone Pairs for \(\mathrm{XeF}_{4}\)

Xe starts with 8 valence electrons. Forming four single bonds with four F atoms uses 4 electrons. The remaining 4 electrons form 2 lone pairs around Xe.

Determine Lone Pairs for \(\mathrm{SF}_{4}\)

S starts with 6 valence electrons. Forming four single bonds with F atoms uses 4 electrons, so 2 electrons remain as 1 lone pair on S.

Determine Lone Pairs for \(\left[I_{3}\right]^{-}\)

Iodine starts with 7 valence electrons. In \(\left[I_{3}\right]^{-}\), it forms two single bonds with other I atoms. Including the negative charge, one more electron is added. After forming bonds, Iodine has 3 lone pairs.

Conclusion

Comparing the lone pairs, \(\left[I_{3}\right]^{-}\) with 3 lone pairs has the maximum number.

Key concepts

Valence Electrons

In atoms, valence electrons are the outermost electrons that play a key role in forming chemical bonds. The number of valence electrons determines how an atom will bond with others. These are typically found in the outermost shell of an atom and can be determined by examining the element’s group number in the periodic table for main group elements.
For instance:

• Chlorine (Cl) has 7 valence electrons.
• Xenon (Xe) has 8 valence electrons.
• Sulfur (S) has 6 valence electrons.
• Iodine (I) has 7 valence electrons.

The manipulation of these electrons, such as sharing or transferring them, is what leads to bond formation, a process we'll look into in more detail in the next section.

Bond Formation

Bond formation involves the coming together of atoms to complete their outer electron shells, usually seeking stability similar to noble gases. This is achieved through either sharing or transferring electrons, depending on the types of atoms involved.
Consider the following examples:

• In XeF₄, xenon forms covalent bonds by sharing electrons with four fluorine atoms.
• Similarly, SF₄ involves sulfur sharing its valence electrons with fluorine atoms to form four covalent bonds.

Notice how bond formation doesn’t always use up all valence electrons, which means some electrons might still be left as lone pairs.

Central Atom

The central atom in a molecule is typically the atom with the highest valence or the least electronegative, allowing it to form multiple bonds with its neighboring atoms. Identifying the central atom is critical as it often has lone pairs affecting the molecule's geometry.
Each example compound features a central atom:

• Cl in ClO₃^-
• Xe in XeF₄
• S in SF₄
• I in I₃^-

The position of these atoms, surrounded by other atoms and lone pairs, dictates the compound's structure and properties.

Negative Charge

When dealing with ions, it's important to consider the presence of a negative charge. This often means an atom has gained extra electrons, influencing its electron count and behavior in bonding.
For example, in I₃^-, the additional negative charge signifies an extra electron, which affects how electrons are distributed around the iodine atom, resulting in the presence of three lone pairs instead of just two if it were neutral.
Understanding charges helps predict the number of electrons available for bonding and lone pair formation.

Electron Counting

Electron counting is a systematic way to determine how many electrons are involved in bonding and how many remain as lone pairs. This involves tallying up valence electrons and then distributing them to form bonds (usually 2 electrons per bond) and any remaining as lone pairs.
For instance, in XeF₄:

• Xe starts with 8 valence electrons.
• Four bonds with fluorine use up 4 valence electrons.
• This leaves 4 electrons as 2 lone pairs.

This simple arithmetic is crucial for drawing accurate Lewis structures and understanding molecule shapes.

Chemical Bonding

Chemical bonding occurs when atoms interact to achieve greater stability, often by completing their valence shells. These interactions manifest as different types of bonds: ionic, covalent, and metallic.
The compounds in the exercise primarily exhibit covalent bonding, which involves sharing electrons between atoms. In some cases, such as polyatomic ions ( I₃^-), the molecule as a whole may carry a charge, influencing how bonds and lone pairs are arranged.
Understanding bonding is fundamental to predicting molecular behavior, reactivity, and properties.