Many important compounds in the chemical industry are derivatives of ethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{4}\right)\). Two of them are acrylonitrile and methyl methacrylate. Complete the Lewis structures, showing all lone pairs. Give approximate values for bond angles \(a\) through \(f .\) Give the hybridization of all carbon atoms. In acrylonitrile, how many of the atoms in the molecule must lie in the same plane? How many \(\sigma\) bonds and how many \(\pi\) bonds are there in methyl methacrylate and acrylonitrile?

Understanding the basics of Lewis structures is crucial for predicting the properties and behavior of molecules. A Lewis structure is a diagram that represents the bonding between atoms of a molecule and the lone pairs of electrons that may exist. In these structures, each dot represents an electron, and a line represents a bond between two atoms.

When drawing Lewis structures, balance is key. You should ensure that each atom, except for hydrogen, follows the octet rule, having eight electrons to fill its outer shell. Hydrogen is an exception as it can be stable with just two electrons. For the molecules in our exercise, the electrons are arranged to reflect the bonding pattern and molecular shape. By observing these drawings, one can predict the connectivity of the atoms and the distribution of electrons which acts as the link to understanding the molecule's geometry and reactivity.

The bond angles of a molecule can greatly affect its shape and physical properties. These angles are determined by the spatial arrangement of bonds around a central atom which is influenced by the electrons' repulsion. In molecules where the central atom has sp² hybridization, such as acrylonitrile, the bond angles are approximately 120°. This corresponds to the trigonal planar arrangement of molecules.

Contrarily, a tetrahedral molecule, where the central atom has an sp³ hybridization, usually has bond angles close to 109.5°. These angles can provide insight into the spatial distribution of a molecule's atoms. Estimating bond angles allows students to better understand molecular shapes and the forces at play within the structure.

The concept of hybridization helps explain the observed shapes and bond angles in molecules. Hybridization occurs when atomic orbitals mix to form new, identical hybrid orbitals. This mixing directly influences molecular bonding and geometry.

For instance, in acrylonitrile and methyl methacrylate, the type of hybridization stipulates the bonding properties: an sp² hybridized atom will have one unhybridized p-orbital available for π bonding which is essential for double bonds. Similarly, an sp³ hybridized atom will form single bonds only. Recognizing hybridization patterns enables a student to predict not just shapes, but also the potential for chemical reactions.

Breaking down bonds into sigma (σ) and pi (π) types offers deeper insight into how atoms in molecules connect. Sigma bonds are the single bonds that form between atoms. They allow for free rotation around the bond axis because they involve head-on overlapping of orbitals. In contrast, pi bonds are formed by the side-to-side overlap of atomic orbitals, typically p orbitals, which occur in double and triple bonds.

Pi bonds restrict the rotation about the bond axis, which contributes to the rigidity of the molecule's structure. By counting the sigma and pi bonds, as done in acrylonitrile and methyl methacrylate, students can better grasp the molecular complexity and the effects on the molecule’s behavior.

The spatial arrangement of atoms in a molecule, or its molecular geometry, is determined by the electronic and steric interactions. The VSEPR (Valence Shell Electron Pair Repulsion) theory is a helpful tool for predicting molecular shapes based on the repulsion between electron pairs in the valence shell of the central atom.

Geometries like linear, trigonal planar, tetrahedral, and others are all derived from the specific electron pair repulsions. In both acrylonitrile and methyl methacrylate, the geometry affects the compound's polarity, reactivity, and physical properties. Grasping the principles of molecular geometry is vital for students to anticipate the behavior of different molecular structures in real-life chemical contexts.