Question

Draw a Lewis structure for each of the following molecules or ions. Describe the electron-pair geometry and molecular geometry around the central atom. (a) \(\mathrm{NH}_{2} \mathrm{Cl}\) (b) \(\mathrm{Cl}_{2} \mathrm{O}(\mathrm{O}\) is the central atom) (c) SCN (d) HOF

Short answer

(a) Trigonal pyramidal; (b) Bent; (c) Linear; (d) Bent.

Step-by-step solution

Understanding Lewis Structures

A Lewis structure is a diagrammatic method to represent the valence electrons in a molecule. It shows how the atoms are bonded to each other and whether there are lone pairs of electrons.

Drawing \\( \mathrm{NH}_2 \mathrm{Cl} \\\) Lewis Structure

For \( \mathrm{NH}_2 \mathrm{Cl}\), nitrogen (\( \mathrm{N} \)) is the central atom. It forms single bonds with two hydrogen atoms and one chlorine atom. There is one lone pair on the nitrogen. So, the Lewis structure is written as H-N-H with a single bond to Cl.

Geometries for \\( \mathrm{NH}_2 \mathrm{Cl} \\\)

The electron-pair geometry around N is tetrahedral since it includes three bonds and one lone pair. The molecular geometry is trigonal pyramidal due to the lone pair on nitrogen.

Drawing \\( \mathrm{Cl}_2 \mathrm{O} \\\) Lewis Structure

For \( \mathrm{Cl}_2 \mathrm{O}\), oxygen (\( \mathrm{O} \)) is the central atom. It forms single bonds with two chlorine atoms. There are two lone pairs on the oxygen atom. The Lewis structure is Cl-O-Cl with lone pairs on oxygen.

Geometries for \\( \mathrm{Cl}_2 \mathrm{O} \\\)

The electron-pair geometry around O is tetrahedral because it involves two bonds and two lone pairs. The molecular geometry is bent or V-shaped due to the lone pairs on oxygen.

Drawing \\( \mathrm{SCN}^- \\\) Lewis Structure

The thiocyanate ion (\( \mathrm{SCN}^- \)) has nitrogen as the central atom in the linear structure S=C=N^-. There are double bonds between the sulfur-carbon and carbon-nitrogen bonds with lone pairs on S and N, and the ion carries a negative charge.

Geometries for \\( \mathrm{SCN}^- \\\)

Both electron-pair and molecular geometries for the SCN^- ion are linear since all atoms are arranged in a straight line with no lone pair affecting the shape.

Drawing \\( \mathrm{HOF} \\\) Lewis Structure

For \( \mathrm{HOF}\), oxygen (\( \mathrm{O} \)) is the central atom. It forms single bonds with hydrogen and fluorine. There are two lone pairs on the oxygen. Hence, the structure is H-O-F with lone pairs on oxygen.

Geometries for \\( \mathrm{HOF} \\\)

The electron-pair geometry around O is tetrahedral, involving two bonds and two lone pairs. The molecular geometry is bent or V-shaped due to lone pairs on the oxygen atom.

Key concepts

Electron-pair Geometry

Electron-pair geometry refers to the arrangement of electron groups around the central atom in a molecule. These groups include both bonded atoms and lone pairs of electrons. Understanding electron-pair geometry helps predict the molecular shape and properties of a compound.
For example, in the molecule \( \mathrm{NH}_2 \mathrm{Cl} \), the nitrogen atom is surrounded by three bonded hydrogen and chlorine atoms, as well as one lone pair. This configuration results in a tetrahedral electron-pair geometry. Even though the lone pair is not visible in molecular shape, it is crucial in influencing the overall geometry.
In the case of \( \mathrm{Cl}_2 \mathrm{O} \), with oxygen as the central atom, the electron-pair geometry is also tetrahedral. There are two chlorine atoms bonded to oxygen and two lone pairs on the oxygen. Like \( \mathrm{NH}_2 \mathrm{Cl} \), the lone pairs affect the shape and angle at which the atoms appear.
Recognizing the electron-pair geometry provides insight into potential interactions with other molecules. Compounds with similar geometries may exhibit similar chemical behavior.

Molecular Geometry

Molecular geometry considers only the atoms within the molecule, ignoring lone pairs even though they influence the final shape. Unlike electron-pair geometry, molecular geometry solely reflects the positions of atoms rather than all electron groups.
For \( \mathrm{NH}_2 \mathrm{Cl} \), ignoring the lone pair on nitrogen, the molecular geometry is trigonal pyramidal. The lone pair pushes the three bonded atoms downward, creating a pyramid-like shape. The bond angles are slightly less than the ideal tetrahedral angles due to lone pair repulsion.
In \( \mathrm{Cl}_2 \mathrm{O} \) and \( \mathrm{HOF} \), the molecular geometries are bent or V-shaped. The presence of lone pairs on the central oxygen atom causes the bonded chlorine atoms and hydrogen in the respective molecules to angle downward, forming a non-linear resting contour. The lone pairs compress the ideal bond angles, further influencing the shape.
To summarize, examining molecular geometry gives us essential information on how molecules interact spatially, which is crucial for determining reactivity and how molecules fit together in larger structures.

Valence Electrons

Valence electrons are the outermost electrons of an atom and participate in forming bonds with other atoms. When constructing Lewis Structures, these electrons are represented as dots surrounding atomic symbols and are fundamental in understanding molecular connections.
Consider \( \mathrm{NH}_2 \mathrm{Cl} \) once more. Nitrogen, with five valence electrons, forms bonds using three of them, and holds onto a lone pair with the remaining two electrons. Similarly, chlorine contributes one valence electron to bond with nitrogen. Hydrogen, with one valence electron, uses it to bond with nitrogen.
In \( \mathrm{Cl}_2 \mathrm{O} \), each chlorine atom uses one valence electron for bonding with oxygen. Oxygen, having six valence electrons, engages two to create bonds and retains its four remaining electrons as two lone pairs.
Understanding valence electrons is crucial because they determine how atoms interact to form stable molecules. Counting and arranging these electrons help predict the most favorable molecular interactions and formations.