Question

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

Short answer

The Lewis structure for the given vile-smelling orange liquid with an empirical formula of SCl and a formal charge of zero on all elements in the compound is: Cl || [S]--↑ Here, the central sulfur atom has an octet with 1 lone pair and a double bond with the chlorine atom, while the chlorine atom also has an octet with 3 lone pairs and a double bond with the sulfur atom. This compound is known as sulfur monochloride (SCl).

Step-by-step solution

Determine the valence electrons

First, we need to determine the number of valence electrons for sulfur (S) and chlorine (Cl). Sulfur is in group 16, so it has 6 valence electrons. Chlorine is in group 17, so it has 7 valence electrons. Now that we have the number of valence electrons, let's proceed with the next step.

Calculate the total number of electrons to be used in the Lewis structure

Since the empirical formula of the compound is SCl, we add the valence electrons of sulfur and chlorine Total valence electrons = 6 (from S) + 7 (from Cl) = 13 valence electrons

Place the least electronegative atom at the center and complete the octets

Sulfur is less electronegative than chlorine, so we put it at the center and surround it with the chlorine atom: Cl | S However, with 13 electrons, we cannot form a complete octet for both S and Cl. So, we need to use a double bond and an odd number of electrons on one of the atoms.

Draw the Lewis structure with complete octets

To fulfill the formal charge of zero on all elements in the compound, we assign an extra electron to the central sulfur atom. The Lewis structure looks like this: Cl || [S]--↑ In this Lewis structure, the central sulfur atom has an octet with 1 lone pair and a double bond with the chlorine atom. The chlorine atom also has an octet with 3 lone pairs and a double bond with the sulfur atom. This is the Lewis structure for the given compound, which is known as sulfur monochloride (SCl).

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom that are involved in forming chemical bonds. They are critical in determining how atoms will interact and bind with one another. For instance, in the exercise where we explore sulfur (S) and chlorine (Cl), we find that sulfur, located in group 16 of the periodic table, has 6 valence electrons, and chlorine, in group 17, has 7 valence electrons.

The number of valence electrons helps predict the reactivity of an element. Atoms typically seek to reach a stable configuration, often analogous to the noble gases with 8 valence electrons, which is referred to as the 'octet rule'. However, there are exceptions to this rule, especially for elements that cannot physically accommodate 8 electrons due to their position on the periodic table.

It's also important to note that the arrangement and pairing of these electrons in the Lewis structure, as seen in our given exercise, are fundamental for predicting the geometry and reactivity of the molecule formed.

Empirical Formula

An empirical formula represents the simplest whole-number ratio of elements in a compound. It does not necessarily reflect the exact number of atoms in a molecule but instead gives the ratio of the types of atoms present. In the context of the provided exercise, the empirical formula SCl tells us that for every one sulfur atom present, there is one chlorine atom.

It is essential to understand the empirical formula is not the same as the molecular formula, which specifies the exact number of atoms of each element in a molecule. However, the empirical formula can be used as a stepping stone to deducing the molecular formula if the molecular weight of the compound is known.

Since the empirical formula provides no information about the structure or geometry of the compound, it is necessary to combine this information with other chemical properties and rules, like the octet rule and formal charges, to determine the structure of the molecule accurately.

Formal Charge

The formal charge is a concept used to determine the charge distribution within a molecule. It is an accounting tool that helps chemists keep track of the electrons in a molecule's Lewis structure. The formal charge can be calculated for each atom in a molecule by using the formula:
\[\begin{equation}\text{Formal Charge} = (\text{Valence Electrons}) - (\text{Non-bonding Electrons}) - \frac{1}{2}(\text{Bonding Electrons})\end{equation}\]

In a stable molecule, the sum of the formal charges on all atoms should be equal to the overall charge of the molecule. For the vile-smelling orange liquid in the exercise, sulfur monochloride, ensuring a formal charge of zero on both the sulphur and chlorine is crucial to drawing the correct Lewis structure.

By calculating formal charges, chemists can predict the most likely structure for a molecule. A correct Lewis structure will have formal charges that match the actual charge distribution as closely as possible, which helps with predicting the stability and reactivity of the molecule.