Question

Write Lewis structures for (a) \(\mathrm{SCl}_{2}\) (b) \(\mathrm{Cl}_{3}^{+}\) (c) \(\mathrm{SOCl}_{2}\) (d) \(\mathrm{ClOClO}_{3}(\) contains \(\mathrm{Cl}-\mathrm{O}-\mathrm{Cl}\) bond \()\)

Short answer

Draw Lewis structures respecting octet rules for each compound, considering bond formation and electron distribution.

Step-by-step solution

Understand the Basics of Lewis Structures

Lewis structures are diagrams that represent chemical bonds between atoms using dots for electrons and lines for shared or bonding pairs. To write them, count total valence electrons, distribute electrons to complete octets, starting with bonds around central atoms.

Write Lewis Structure for \( \mathrm{SCl}_{2} \)

Determine valence electrons: Sulfur has 6, and each Chlorine has 7, totaling 20 valence electrons (\(6 + 7*2\)). Place sulfur in the center with two chlorine atoms attached with single bonds (2 electrons each, 4 total). Distribute remaining 16 electrons to complete chlorine octets first, then sulfur's.

Write Lewis Structure for \( \mathrm{Cl}_{3}^{+} \)

Determine valence electrons: Each Chlorine has 7, totaling 21, minus 1 for the positive charge yields 20 valence electrons. Place one chlorine in the center with single bonds to both terminal chlorines (4 electrons for bonds). Distribute remaining 16 electrons to complete octets, adjusting for central atom deficit if necessary.

Write Lewis Structure for \( \mathrm{SOCl}_{2} \)

Determine valence electrons: Sulfur has 6, Oxygen has 6, and each Chlorine has 7, totaling 26 electrons. Place sulfur in the center bonded to oxygen and both chlorines, using single bonds for the chlorines (4 electrons) and a double bond for the oxygen (4 electrons). Distribute remaining 18 electrons to complete octets.

Write Lewis Structure for \( \mathrm{ClOClO}_{3} \)

Determine the structure based on the given Cl-O-Cl bond. Calculate total valence electrons: \( \mathrm{Cl} = 7 \), \( \mathrm{O} = 6 \) each, totaling 38 electrons. Use single bonds for Cl-O-Cl, then attach \( \mathrm{O}_3 \) with appropriate multiple bonds to account for octet fulfillment for all atoms.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons in an atom and play a crucial role in forming chemical bonds. These electrons are involved in bonding as they can be shared or exchanged between atoms. To determine the number of valence electrons in an atom, one can refer to the periodic table. For example, sulfur has 6 valence electrons and chlorine has 7. In a molecule like \(\mathrm{SCl}_{2}\), you first tally up the valence electrons: sulfur's 6 plus the two chlorines' 7 each, making a total of 20 valence electrons.
Understanding the valence electrons is the first step in drawing Lewis structures, as these electrons are depicted as dots around the elemental symbol to represent potential bonds. Knowing how these electrons are shared allows the prediction of molecule shapes and interactions, important for understanding chemical reactions.
This understanding aids in visualizing how atoms connect and the number of bonds formed during compound formation.

Chemical Bonds

Chemical bonds hold atoms together to form molecules. These bonds are formed by the sharing or transferring of valence electrons between atoms. There are primarily two types of bonds: covalent and ionic, but within Lewis structures, we mainly focus on covalent bonds.
Each line in a Lewis structure represents a pair of shared electrons, known as a bonding pair. For instance, in \(\mathrm{SCl}_{2}\), each sulfur-chlorine bond is depicted by a line connecting the atoms, representing the two electrons they share.
It's essential to determine which atoms will connect and how many bonds can form between them, which often depends on the number of unpaired electrons an atom has. By doing this, one can figure out how to satisfy the octet rule, creating stable molecules.

Octet Rule

The octet rule is a fundamental chemical principle that states most atoms aim to have eight electrons in their valence shell to achieve stability, similar to noble gases. This rule guides the distribution of electrons in a Lewis structure.
For sulfur and chlorine in \(\mathrm{SCl}_{2}\), chlorine atoms must complete their octet by sharing electrons with sulfur. Once sulfur forms two bonds with chlorine, the total shared and unshared electrons around it ideally need to equal 8 to satisfy the octet rule.
However, exceptions exist, such as molecules with sulfur and phosphorus, where these atoms can have more than eight electrons, forming expanded octets, mainly when bonded to highly electronegative atoms like oxygen.

Central Atom

In Lewis structures, the central atom is typically the one that can form the most bonds. Often, it is the least electronegative element or the one present in the smallest quantity within the molecule.
For the molecule \(\mathrm{SCl}_{2}\), sulfur serves as the central atom because it can form single bonds with both chlorines. The central atom is crucial because it usually provides a foundation around which other atoms attach themselves.
Identifying the central atom effectively can simplify the process of distributing valence electrons to satisfy all atoms' needs, especially when considering those that require more or fewer electrons for stability. For complex molecules, careful consideration of bond angles and molecule geometry around the central atom is needed.

Single and Multiple Bonds

Single and multiple bonds play a critical role in complete understanding of a molecule's structure. A single bond, represented by a single line in a Lewis structure, implies two shared electrons between two atoms. This is the simplest bond form seen in \(\mathrm{SCl}_{2}\), where sulfur and chlorine form single bonds.
Multiple bonds, such as double and triple bonds, involve two or three pairs of shared electrons, respectively, and provide more structural stability to particular molecules. In \(\mathrm{SOCl}_{2}\), the bond between sulfur and oxygen is depicted as a double bond, reflecting four shared electrons.
Such bonds influence the molecule's properties, including its reactivity and rigidity. Identifying the type of bond in Lewis structures helps predict how a molecule will interact chemically and physically in different environments.