Use hybrid orbitals to draw the following molecules.
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Verified step by step guidance

1. Step 1: Understand the concept of hybrid orbitals. Hybrid orbitals are formed by the mixing of atomic orbitals (such as s, p, and sometimes d orbitals) to create new orbitals that are better suited for bonding in molecules. Common types include sp, sp², and sp³ hybridization.
2. Step 2: Determine the central atom in the molecule and its steric number. The steric number is calculated as the sum of the number of bonded atoms and lone pairs around the central atom. This will help identify the type of hybridization.
3. Step 3: Assign the hybridization based on the steric number. For example: steric number 2 corresponds to sp hybridization, steric number 3 corresponds to sp² hybridization, and steric number 4 corresponds to sp³ hybridization.
4. Step 4: Draw the molecule using the hybrid orbitals. Represent the central atom with its hybrid orbitals and show the bonds formed with other atoms. Use appropriate bond angles based on the hybridization (e.g., 180° for sp, 120° for sp², and 109.5° for sp³).
5. Step 5: Include lone pairs and pi bonds if applicable. Lone pairs occupy hybrid orbitals, while pi bonds are formed by unhybridized p orbitals overlapping sideways. Ensure the drawing reflects the correct geometry and bonding.

Key Concepts

Hybridization

Hybridization is the process of combining atomic orbitals to form new hybrid orbitals that can accommodate the bonding requirements of a molecule. This concept is crucial for understanding molecular geometry and bonding properties, as it explains how atoms can form equivalent bonds in molecules. For example, in methane (CH4), the carbon atom undergoes sp3 hybridization, resulting in four equivalent bonds with hydrogen atoms.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. It is determined by the types of hybrid orbitals involved and the number of electron pairs around the central atom. Understanding molecular geometry is essential for predicting the shape and reactivity of molecules, as seen in the tetrahedral shape of methane due to sp3 hybridization.

Bonding and Lone Pairs

In molecular structures, bonding pairs of electrons are shared between atoms, while lone pairs are non-bonding electrons localized on a single atom. The presence of lone pairs affects the geometry and angles between bonds, as they occupy space and repel bonding pairs. For instance, in water (H2O), the two lone pairs on oxygen lead to a bent molecular shape, deviating from the ideal tetrahedral angle.