Predict the molecular structure (including bond angles) for each of the following. a. \(\mathrm{PCl}_{3}\) b. \(\mathrm{SCl}_{2}\) c. \(\mathrm{SiF}_{4}\)

1. Central atom: Phosphorus (P) is the central atom, as it is less electronegative than Chlorine (Cl). 2. Bonding electron pairs: \(3\) P-Cl bonds, making it \(3\) bonding electron pairs. 3. Lone electron pairs: To find the number of lone pairs, first determine the number of valence electrons for P and Cl. P has \(5\) valence electrons, and each Cl has \(7\). Since there are \(3\) Cl atoms, the total number of electrons is \((5 + 3\times 7) = 26\). Subtract \(3\times2\) electrons from the bonds, which leaves \(20\) electrons in \(10\) lone pairs. Two lone pairs are on the central atom P, and the others are on the Cl atoms. 4. With \(3\) bonding pairs and \(1\) lone pair on the central atom, we have a tetrahedral arrangement of electron pairs, a "trigonal pyramidal" molecular geometry, and bond angles of around \(109.5^{\circ}\).

1. Central atom: Sulfur (S) is the central atom, as it is less electronegative than Chlorine (Cl). 2. Bonding electron pairs: \(2\) S-Cl bonds, making it \(2\) bonding electron pairs. 3. Lone electron pairs: To find the number of lone pairs, first determine the number of valence electrons for S and Cl. S has \(6\) valence electrons, and each Cl has \(7\). Since there are \(2\) Cl atoms, the total number of electrons is \((6 + 2\times 7) = 20\). Subtract \(2\times2\) electrons from the bonds, which leaves \(16\) electrons in \(8\) lone pairs. Two lone pairs are on the central atom S, and the others are on the Cl atoms. 4. With \(2\) bonding pairs and \(2\) lone pairs on the central atom, we have a tetrahedral arrangement of electron pairs, a "bent" or "v-shaped" molecular geometry, and bond angles of around \(104.5^{\circ}\).

1. Central atom: Silicon (Si) is the central atom, as it is less electronegative than Fluorine (F). 2. Bonding electron pairs: \(4\) Si-F bonds, making it \(4\) bonding electron pairs. 3. Lone electron pairs: To find the number of lone pairs, first determine the number of valence electrons for Si and F. Si has \(4\) valence electrons, and each F has \(7\). Since there are \(4\) F atoms, the total number of electrons is \((4 + 4\times 7) = 28\). Subtract \(4\times 2\) electrons from the bonds, which leaves \(20\) electrons in \(10\) lone pairs. There are no lone pairs on the central atom Si, the others are on the F atoms. 4. With \(4\) bonding pairs and no lone pairs on the central atom, we have a tetrahedral arrangement of electron pairs, a "tetrahedral" molecular geometry, and bond angles of around \(109.5^{\circ}\).