Question

Write Lewis structures and predict whether each of the following is polar or nonpolar a. \(\mathrm{HOCN}\) (exists as \(\mathrm{HO}-\mathrm{CN} )\) b. \(\mathrm{COS}\) c. \(\mathrm{XeF}_{2}\) d. \(\mathrm{CF}_{2} \mathrm{Cl}_{2}\) e. \(\mathrm{SeF}_{6}\) f. \(\mathrm{H}_{2} \mathrm{CO}(\mathrm{C} \text { is the central atom })\)

Short answer

The molecules' polarities based on their Lewis structures and molecular geometries are as follows: a. \(\mathrm{HOCN}\): Polar b. \(\mathrm{COS}\): Polar c. \(\mathrm{XeF}_{2}\): Nonpolar d. \(\mathrm{CF}_{2}\mathrm{Cl}_{2}\): Nonpolar e. \(\mathrm{SeF}_{6}\): Nonpolar f. \(\mathrm{H}_{2}\mathrm{CO}\): Polar

Step-by-step solution

1. Draw the Lewis Structures for Each Molecule:

To draw the Lewis structures, we need to determine the number of valence electrons in each molecule, place the least electronegative atom in the center (except for hydrogen, which always goes on the outside), create bonds, and fill in the remaining valence electrons. a. \(\mathrm{HOCN}\): - Total valence electrons: \(1 + 6 + 5 + 4 = 16\) - Central atom: O - Lewis Structure: H-O-C≡N b. \(\mathrm{COS}\): - Total valence electrons: \(4 + 6 + 6 = 16\) - Central atom: C - Lewis Structure: O=C=S c. \(\mathrm{XeF}_{2}\): - Total valence electrons: \(8 + 2 \times 7 = 22\) - Central atom: Xe - Lewis Structure: F-Xe-F, Xe has 3 lone pairs around it. d. \(\mathrm{CF}_2\mathrm{Cl}_2\): - Total valence electrons: \(4 + 2 \times 7 + 2 \times 7 = 28\) - Central atom: C - Lewis Structure: F-C-Cl, 90-degree angles, with F and Cl opposite each other and two lone pairs on C. e. \(\mathrm{SeF}_{6}\): - Total valence electrons: \(6 + 6 \times 7 = 48\) - Central atom: Se - Lewis Structure: Se is surrounded by six F atoms which form a hexagonal configuration. No lone pairs on Se. f. \(\mathrm{H}_{2}\mathrm{CO}\): - Total valence electrons: \(2 \times 1 + 4 + 6 = 12\) - Central atom: C - Lewis Structure: H-C-O; C is bonded to each H and to O with a double bond.

2. Determine the Molecular Geometry of Each Molecule:

a. \(\mathrm{HOCN}\): Linear (O in the center) b. \(\mathrm{COS}\): Linear c. \(\mathrm{XeF}_{2}\): Linear d. \(\mathrm{CF}_{2}\mathrm{Cl}_{2}\): Tetrahedral e. \(\mathrm{SeF}_{6}\): Octahedral f. \(\mathrm{H}_{2}\mathrm{CO}\): Trigonal planar around carbon

3. Assess the Polarity of Each Molecule:

a. \(\mathrm{HOCN}\): Polar due to the difference in electronegativity between H, O, C, and N, and a linear shape, which does not cancel out the dipole moment. b. \(\mathrm{COS}\): Polar due to the difference in electronegativity between O, C, and S, and a linear shape, which does not cancel out the dipole moment. c. \(\mathrm{XeF}_{2}\): Nonpolar due to the linear shape, which helps to cancel out dipole moments between F atoms. d. \(\mathrm{CF}_{2}\mathrm{Cl}_{2}\): Nonpolar due to the symmetrical tetrahedral shape, which cancels out any polar bonds. e. \(\mathrm{SeF}_{6}\): Nonpolar due to the symmetrical octahedral shape, which cancels out any polar bonds. f. \(\mathrm{H}_{2}\mathrm{CO}\): Polar due to the difference in electronegativity between H, C, and O, and a trigonal planar shape, which does not cancel out the dipole moment.

Key concepts

Molecular Geometry

Understanding molecular geometry is crucial for predicting the shape and properties of molecules. It refers to the three-dimensional arrangement of atoms within a molecule. By knowing the molecular geometry, you can predict how a molecule will interact with others, its reactivity, and its physical properties.

• Linear Geometry: A molecule is linear if its atoms are arranged in a straight line, as seen in HOCN, COS, and XeF₂. This shape often results in bond angles of 180 degrees, making molecules linear in appearance.
• Tetrahedral Geometry: CF₂Cl₂ demonstrates a tetrahedral shape, which means it has four atoms symmetrically surrounding the central atom, leading to bond angles of about 109.5 degrees.
• Octahedral Geometry: SeF₆ has an octahedral shape, indicating six atoms symmetrically distributed around a central atom, creating bond angles of 90 degrees.
• Trigonal Planar Geometry: In H₂CO, the trigonal planar shape is observed, with atoms arranged in a plane around the central atom, resulting in 120-degree bond angles.

Knowing these geometries helps in predicting the molecule's properties, including how it might interact with electric fields, which is important in chemistry.

Polarity of Molecules

Polarity in molecules is determined by differences in electronegativity and the symmetry of the molecule. A molecule is considered polar if it has regions of partial positive and negative charge due to an unequal distribution of electrons.

• Polar Molecules: Examples include HOCN, COS, and H₂CO. The lack of symmetry and significant differences in electronegativity lead to partial charges that do not cancel each other out, resulting in overall polarity.
• Nonpolar Molecules: Examples include XeF₂, CF₂Cl₂, and SeF₆. Despite some polar bonds, their symmetrical shape allows for the dipole moments to cancel out, leading to nonpolar characteristics.

This concept is crucial in determining how a molecule interacts with water and other solvents, with polar molecules generally being more soluble in polar solvents.

Valence Electrons

Valence electrons are the outermost electrons involved in chemical bonding. Understanding the number of valence electrons is essential for drawing Lewis structures and predicting molecular interactions. Each element in the periodic table has a specific number of valence electrons:

• HOCN: 16 valence electrons are used here, coming from hydrogen (1), oxygen (6), carbon (4), and nitrogen (5).
• COS: Utilizes 16 valence electrons from carbon (4), oxygen (6), and sulfur (6).
• XeF₂: Involves 22 valence electrons, with xenon contributing 8 and fluorine 7 each.
• CF₂Cl₂: Involves 28 valence electrons from carbon (4), fluorine (7 each), and chlorine (7 each).
• SeF₆: Takes 48 from selenium (6) and the surrounding fluorine atoms (7 each).
• H₂CO: Combines 12 from the hydrogen (1 each), carbon (4), and oxygen (6).

By counting the valence electrons correctly, one can draw accurate Lewis structures and predict molecular bonding and reactivity.

Molecular Shape

Molecular shape is important for determining how molecules fit together and interact. This shape results from the electron pair geometry and affects the molecule's properties and reactivity.

• Linear: Molecules like HOCN, COS, and XeF₂ exhibit linear shapes, which are often associated with certain chemical characteristics, such as being less complex for potential reactions.
• Tetrahedral: CF₂Cl₂'s tetrahedral shape creates a symmetrical distribution of charge which affects its interaction with electric fields and solvents.
• Octahedral: The octahedral shape of SeF₆ results in maximum spatial separation of ligands, minimizing repulsion between electron pairs.
• Trigonal Planar: In H₂CO, this shape implies readiness for certain reactions due to its flat, triangular arrangement.

Understanding molecular shapes aids in predicting the behavior of molecules in reactions, including their stereochemistry and potential energy levels.

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. Differences in electronegativity between elements can cause bonds to be polar covalent.

• Polar Bonds: Are formed when there is a significant difference in electronegativity, such as between hydrogen and oxygen in HOCN, or carbon and oxygen in H₂CO.
• Nonpolar Bonds: Occur when atoms share electrons equally or when the dipole moments cancel out in symmetrical molecules like SeF₆.

The concept of electronegativity is essential for predicting polarity and for understanding the distribution of electrons within molecules. It is a fundamental property that influences molecule behavior, including bonding and reactivity.