Question

When molten sulfur reacts with chlorine gas, a vile-smelling orange liquid forms that has an empirical formula of \(\mathrm{SCl}\). The structure of this compound has a formal charge of zero on all elements in the compound. Draw the Lewis structure for the vile-smelling orange liquid.

Short answer

The Lewis structure for the vile-smelling orange liquid with an empirical formula of \(\mathrm{SCl}\) and a formal charge of zero on all elements in the compound is: Cl :: S == Cl ::

Step-by-step solution

Determine the number of valence electrons for each element

For Sulfur (S), it is in Group 16 (or Group VI) of the periodic table, having 6 valence electrons. Chlorine (Cl) is in Group 17 (or Group VII) of the periodic table, having 7 valence electrons. In the compound, there is 1 sulfur atom and 1 chlorine atom, so the total number of valence electrons is \(6+7=13\).

Place the least electronegative atom in the center

Since sulfur is less electronegative than chlorine, we place sulfur in the center and connect it with chlorine using a single bond, which accounts for 2 of the 13 valence electrons. S - Cl

Distribute the remaining valence electrons

We have 11 valence electrons left. We will first complete the octet for chlorine, using 6 more electrons (3 lone pairs). S - Cl : : Now, we have 5 valence electrons remaining. We will place these electrons on the sulfur atom. S - Cl : : : : This gives Sulfur a total of 7 electrons. However, now we still lack one electron to complete the Sulfur octet and form a neutral molecule.

Modify the structure to achieve a formal charge of zero for both elements

To correct the lack of an electron for the sulfur atom and achieve a formal charge of zero for both elements, we will move one of the lone pairs on chlorine to create a double bond with sulfur. S = Cl : : : : Now, both sulfur and chlorine have a formal charge of zero and their octets are complete (Sulfur having 6+2=8 electrons and Chlorine having 6+2=8 electrons).

Draw the Lewis structure

The Lewis structure for the compound (vile-smelling orange liquid) \(\mathrm{SCl}\) is: Cl :: S == Cl :: Now, we have successfully drawn the Lewis structure for the vile-smelling orange liquid with a formal charge of zero for all the elements in the compound.

Key concepts

Understanding Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom. They play a pivotal role in chemical bonding because they can be gained, lost, or shared to form chemical bonds. The number of valence electrons that an atom has can usually be determined by its group number on the periodic table. For example, sulfur (S), which is in Group 16, possesses six valence electrons, while chlorine (Cl) in Group 17 has seven valence electrons.

To depict how atoms bond, Lewis structures are used. These diagrams represent valence electrons as dots around the elemental symbols. A good starting point is to calculate the total number of valence electrons in the compound. This is critical when drawing the Lewis structure because it allows a chemist to determine how the electrons are distributed among the atoms, forming a molecule.

Exercise Improvement Advice: Valence Electrons

To fully understand the Lewis structure of a molecule, always start by calculating the total number of valence electrons. This not only aides in creating an accurate Lewis structure but also ensures that you carefully account for all the outer electrons that can participate in bonding.

Calculating Formal Charge

The formal charge is a concept that helps us predict the distribution of electrons in molecular structures. It's calculated by taking the number of valence electrons in an isolated atom, subtracting the number of electrons assigned to the atom in the molecule (counting each bond as one electron and non-bonding as two). A formal charge of zero indicates that the atom has a neutral charge within the molecule.

For accurate representation, distributing the electrons appropriately to achieve a formal charge of zero wherever possible is important — this is often a sign of the most stable structure. In the exercise, after initial placement of electrons and bonds, one of the chlorine's lone pairs was moved to form a double bond with sulfur to achieve that stability.

Exercise Improvement Advice: Formal Charge

Always check the formal charge to ensure you have the most stable Lewis structure. Double-checking formal charges can lead to revisions in the structure, like creating double or triple bonds where necessary to attain a neutral charge or the lowest possible formal charges on each atom.

Applying the Octet Rule

The octet rule is a chemical rule of thumb stating that atoms are most stable when they have eight electrons in their valence shell. This rule guides the placement of electrons around the atoms to fulfill the atoms' 'desire' for a complete octet, leading to greater stability. When drawing the Lewis structure for compounds, the aim is to give most of the atoms (particularly the main-group elements) a full octet.

In our exercise, the octets for Sulfur (S) and Chlorine (Cl) are completed by sharing electrons. Initially, Sulfur had only six electrons after forming a single bond, and a double bond was created to complete its octet. Doing so honored the octet rule and ensured a stable molecule while also satisfying the condition of having a formal charge of zero.

Exercise Improvement Advice: Octet Rule

Remember to satisfy the octet rule whenever possible, which may involve forming multiple bonds, as seen in the example problem. Ensuring that each atom has a complete octet (except for exceptions to the rule like hydrogen and beryllium) can sometimes resolve issues with formal charges and offer the most stable structure for a molecule.