Problem 81 - Chemistry The Molecular Nature of Matter

Question

Ethene, also called ethylene, is a gas used to ripen tomatoes artificially. It has the Lewis structure: What would you expect the \(\mathrm{H}-\mathrm{C}-\mathrm{H}\) and \(\mathrm{H}-\mathrm{C}=\mathrm{C}\) bond angles to be in this molecule? (Caution: Don't be fooled by the way the structure is drawn here.)

Step-by-Step Solution

Answer

The \text{H}-\text{C}-\text{H} bond angles are approximately 120 degrees, as are the \text{H}-\text{C}=\text{C} bond angles due to the trigonal planar geometry around each carbon atom.

Step 1: Identify the Geometry

Identify the geometry of the molecule using the Lewis structure and VSEPR theory. For ethene, each carbon atom is bonded to three other atoms (two hydrogens and one carbon) and has no lone pairs. This arrangement corresponds to a trigonal planar geometry around each carbon atom.

Step 2: Determine Bond Angles

Based on the trigonal planar geometry around each carbon atom, the expected bond angles are approximately 120 degrees. This is because in a trigonal planar shape, the atoms are spaced evenly around the central atom, creating three 120-degree angles.

Step 3: Correct Misconceptions

Do not be misled by the structural depiction that might suggest otherwise; all atoms lie in the same plane due to the double bond’s rigidity, and the bond angles are not significantly affected by how the Lewis structure is drawn.

Key Concepts

Lewis Structures

A Lewis structure is a diagram that represents the valence electrons of atoms within a molecule. These diagrams help us visualize how atoms bond and what molecular structure they create. In constructing a Lewis structure, you will notice that electron pairs either form a bond between two atoms or exist as lone pairs around a single atom. Let's understand this with an example: to draw the Lewis structure for ethene (C2H4), we start by calculating total valence electrons, which are 12 (4 from each carbon and 1 from each hydrogen). We then arrange these electrons to denote bonds between atoms; each line represents a pair of shared electrons.
Remember, carbon atoms form 4 bonds (either single, double, or triple) and hydrogen forms 1 bond. Ethene has a double bond between the two carbon atoms, and each carbon forms two single bonds with hydrogen.

Trigonal Planar Geometry

When discussing molecular geometry, trigonal planar is a term you'll often encounter, especially with molecules where a central atom forms three bonds and has no lone pairs. This configuration spreads the bonded atoms as evenly as possible in a plane about the central atom, leading to a flat, triangular shape. Think of it like a tripod, where all 'legs' are equally spaced to maintain balance. In ethene, each carbon acts as a central atom, forming a trigonal planar setup, which is critical for the molecule's overall structure and stability.

Bond Angles

The bond angles in a molecule profoundly impact its shape and reactivity. In ethene, the ideal bond angles for trigonal planar geometry are 120 degrees. This angle ensures that the electron pairs are as far apart as possible, minimizing repulsion as per VSEPR (Valence Shell Electron Pair Repulsion) theory. For the H–C–H and H–C=C bond angles in ethene, you would expect these to also be roughly 120 degrees. However, always consider the three-dimensional arrangement rather than the simplistic flat drawings often depicted in textbooks, which can sometimes be misleading.

Ethylene Molecular Structure

Ethylene (C2H4) is a simple hydrocarbon with a very important role in the agricultural industry because it influences the ripening of fruit. Its molecular structure consists of a double bond between two carbon atoms, which enforces a plane—imagine all atoms in ethylene lying flat on a table, with no atom sticking up or down. This rigidity is essential for the molecule's function but also impacts its physical and chemical properties, such as its bond angles. The double bond involves more than just shared electron pairs; it also includes 'pi' bonds, which result from the sideways overlap of p orbitals and restrict the molecule to its planar configuration.