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 vile-smelling orange liquid.

Short answer

The final Lewis structure for the vile-smelling orange liquid with the empirical formula SCl and a formal charge of zero on all elements is: S = Cl : :

Step-by-step solution

Determine the Number of Valence Electrons

Sulfur (S) is in group 16 and has 6 valence electrons, while Chlorine (Cl) is in group 17 and has 7 valence electrons. As the empirical formula is SCl, we have one atom of Sulfur and one atom of Chlorine in our molecule. The total number of valence electrons for the compound is: Total Valence Electrons = (S_valence_electrons) + (Cl_valence_electrons) = 6 + 7 = 13

Create the Initial Structure

We start by placing the least electronegative atom, sulfur (S), in the center and put the other atom (chlorine) next to it. We place a single bond between sulfur and chlorine, representing two shared electrons. S - Cl

Distribute Remaining Valence Electrons

We have already used two valence electrons for the single bond between S and Cl. Hence, we need to distribute the remaining 11 electrons (13 - 2 = 11). Start with the outer atom (Cl) by adding electron pairs until it achieves an octet. S - Cl : : : Sulfur (S) also needs to achieving its octet. However, adding an electron pair to sulfur will lead to more than eight electrons in its outer shell. To get around this, we can create a double bond with chlorine.

Create the Double Bond

To achieve an octet for both atoms without violating the octet rule for sulfur, we can move one electron pair from chlorine to form a double bond with sulfur. This will give eight electrons to both sulfur and chlorine. S = Cl : :

Check the Formal Charges and Complete the Lewis Structure

Now, we check the formal charges of each atom. The formal charge for sulfur (S) and chlorine (Cl) is: Formal_charge_S = [Valence_electrons_S] - [Non_bonding_electrons_S] - [0.5*Bonding_electrons_S] = 6 - 0 - (0.5*4) = 0 Formal_charge_Cl = [Valence_electrons_Cl] - [Non_bonding_electrons_Cl] - [0.5*Bonding_electrons_Cl] = 7 - 4 - (0.5*4) = 0 As both formal charges equal to zero, we have achieved the desired Lewis structure with a formal charge of zero on all elements. Final Lewis Structure: S = Cl : :

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in chemical bonding. They determine how an atom interacts with others. To find the number of valence electrons, look to the periodic table. Atoms belonging to the same group typically have the same number of valence electrons. For example:

• Sulfur (S) is in group 16, which means it has 6 valence electrons.
• Chlorine (Cl) is in group 17, so it has 7 valence electrons.

By adding up the valence electrons of all atoms in a molecule, you can determine the molecule's total valence electron count. For SCl, this adds up to 13 because we sum the valences of S and Cl: 6 + 7 = 13. This figure is key to correctly drawing the Lewis structure, as it determines the bonding and arrangement of surrounding electrons.

Octet Rule

The octet rule is a guiding principle in chemistry that atoms tend to form bonds until they are surrounded by eight valence electrons, achieving a stable configuration akin to that of noble gases. This rule is fundamental when constructing Lewis structures. However, it's important to note that not all elements follow this rule strictly due to size or other atomic properties. For instance:

• The octet rule suggests that chlorine predominantly forms single bonds as it seeks to gain an electron to complete its octet.
• Sulfur, being in the third period of the periodic table, can expand beyond its octet when forming compounds, especially with elements from the same period or lower.

In SCl, we attempt to satisfy the octet rule for both atoms. Initially, we form a single bond between sulfur and chlorine. However, chlorine requires additional electrons to complete its octet. By adjusting the bonds, particularly through double bonds, we ensure both atoms achieve their preferred electron configurations.

Formal Charge

The concept of formal charge helps identify the most stable Lewis structure by indicating the distribution of charges across a molecule. Calculating the formal charge involves the following steps:

• Identify the number of valence electrons in the free atom.
• Subtract the non-bonding (lone pair) electrons from this figure.
• Subtract half of the bonding (shared) electrons used in molecular bonds.

For instance, with sulfur in \[ ext{SCl}_1 \] super molecule:

• The formal charge for sulfur (S) is zero, as calculated by: 6 (valence) - 0 (lone pair) - 4/2 (bonding) = 0.
• Similarly for chlorine (Cl): 7 - 4 (non-bonding) - 4/2 (bonding) = 0.

Having a formal charge of zero for all atoms in the molecule points to a balanced and stable structure, which is a primary goal when deriving a Lewis structure.

Chemical Bonding

Chemical bonding is the force that holds atoms together within a compound, enabling the formation of molecules and compounds through electron sharing and exchange. There are various types of chemical bonds:

• **Covalent bonds:** These involve the sharing of valence electrons between atoms to attain stability, akin to what is seen in most organic compounds or between nonmetals such as sulfur and chlorine.
• **Ionic bonds:** Characterized by the transfer of electrons from one atom to another, typically between a metal and a non-metal, resulting in a bond due to opposing charges.
• **Metallic bonds:** Found in metals, where electrons are shared among many atoms, creating a 'sea of electrons' effect.

In \[ ext{SCl}_1 \] super molecule, a covalent bond forms as sulfur and chlorine share their electrons to satisfy their respective valence needs. This sharing permits SCl to exist as a stable compound, with a double bond ensuring both atoms satisfy their valence electron requirements effectively.