Question

Draw a Lewis structure for each of the following molecules, and then use the VSEPR theory to predict the molecular geometry of each: a. \(\mathrm{SCl}_{2}\) b. \(\mathrm{PI}_{3}\) c. \(\mathrm{Cl}_{2} \mathrm{O}\) d. \(\mathrm{NH}_{2} \mathrm{Cl}\) e. \(\mathrm{SiCl}_{3} \mathrm{Br}\) f. \(\mathrm{ONCl}\)

Short answer

For \(\text{SCl}_{2}\), the geometry is bent. For \(\text{PI}_{3}\), the geometry is trigonal pyramidal. For \(\text{Cl}_{2}\text{O}\), it's bent. \(\text{NH}_{2}\text{Cl}\) is trigonal pyramidal. \(\text{SiCl}_{3}\text{Br}\) is tetrahedral and \(\text{ONCl}\) is linear.

Step-by-step solution

- Determine the total number of valence electrons

Sum the number of valence electrons for all atoms in the molecule. For example, in \(\text{SCl}_{2}\), sulfur has 6 valence electrons and each chlorine has 7, so the total is 6 + 7 + 7 = 20 electrons.

- Draw the skeletal structure

Determine the central atom (usually the least electronegative, except hydrogen) and arrange the surrounding atoms around it. For \(\text{SCl}_{2}\), sulfur is the central atom with two chlorine atoms around it.

- Distribute the remaining electrons

Subtract the bonding electrons (2 per bond) from the total valence electrons and distribute the remaining electrons to satisfy the octet rule where possible. For \(\text{SCl}_{2}\), 2 bonds = 4 electrons, so 20 - 4 = 16 electrons need to be distributed.

- Verify and place lone pairs

Check that each atom satisfies the octet rule (8 electrons) by adding lone pairs. Adjust to reflect any need for double bonds if octets aren't satisfied. For \(\text{SCl}_{2}\), place three lone pairs on each chlorine and two lone pairs on sulfur.

- Use VSEPR theory to predict molecular geometry

Based on the number of bonding pairs and lone pairs around the central atom, use VSEPR theory to predict the geometry. For example, \(\text{SCl}_{2}\) has 2 bonds and 2 lone pairs leading to a bent shape (like \(\text{H}_{2}\text{O}\)).

- Repeat for each molecule

Apply the same steps for \(\text{PI}_{3}\), \(\text{Cl}_{2}\text{O}\), \(\text{NH}_{2}\text{Cl}\), \(\text{SiCl}_{3}\text{Br}\), \(\text{ONCl}\).

Key concepts

Molecular geometry

Understanding the molecular geometry of a molecule helps us predict its shape and properties. Molecular geometry is determined by the arrangement of atoms around the central atom, which is influenced by the number of bonding pairs and lone pairs of electrons.
VSEPR theory (Valence Shell Electron Pair Repulsion) is a key tool used to predict the geometry. According to VSEPR theory, electron pairs repel each other and arrange themselves to minimize this repulsion, leading to specific geometric shapes.
For example, in \(\text{SCl}_{2}\), sulfur has two bonding pairs with chlorine atoms and two lone pairs. This arrangement leads to a bent shape, similar to \(\text{H}_{2}\text{O}\).

Valence electrons

Valence electrons are the outermost electrons of an atom and are crucial in forming bonds. Each element has a specific number of valence electrons that can be determined from its group number in the periodic table.
For instance:

• Sulfur (group 16) has 6 valence electrons.
• Chlorine (group 17) has 7 valence electrons.

To draw the Lewis structure, sum the total number of valence electrons for all atoms in the molecule. For \(\text{SCl}_{2}\), sulfur contributes 6 electrons, and each chlorine contributes 7, so the total is 20 electrons.

Octet rule

The octet rule states that atoms tend to form bonds to have eight electrons in their valence shell, achieving a stable electron configuration similar to noble gases.
When drawing Lewis structures, ensure that most atoms (except hydrogen, which needs only 2) have an octet. For instance, in \(\text{SCl}_{2}\), after forming bonds, each hydrogen should have 8 electrons in its valence shell.
There are some exceptions to the octet rule like molecules with odd numbers of electrons or central atoms from the third period and beyond that can hold more than eight electrons.

Bonding pairs

Bonding pairs are pairs of electrons shared between atoms to form covalent bonds. Each pair of shared electrons represents a single bond.
For example, in the molecule \(\text{SCl}_{2}\), sulfur forms two single bonds with chlorine atoms, sharing one pair of electrons with each chlorine.
These shared pairs of electrons form bonding pairs, holding the atoms together within the molecule and contributing to its overall stability.

Lone pairs

Lone pairs are pairs of electrons that are not involved in bonding and remain on a single atom. They play a significant role in determining the shape and properties of the molecule.
In the Lewis structure of \(\text{SCl}_{2}\), sulfur has two lone pairs of electrons in addition to the two bonding pairs. Each chlorine atom also has three lone pairs of electrons.
Lone pairs repel bonding pairs, causing the molecule to adopt a specific shape (e.g., the bent shape of \(\text{SCl}_{2}\)). Understanding lone pairs is crucial for predicting molecular geometry through VSEPR theory.