Question

Describe the molecular geometry of \(\mathrm{H}_{2} \mathrm{O}\) suggested by each of the following methods: (a) Lewis theory; (b) valence-bond method using simple atomic orbitals; (c) VSEPR theory; (d) valence-bond method using hybridized atomic orbitals.

Short answer

The molecular geometry of \(H_2O\) is bent. This conclusion is reached using diverse methods: Lewis theory, valence-bond method with simple atomic orbitals, VSEPR theory, and the valence-bond method with hybridized atomic orbitals.

Step-by-step solution

Lewis Theory Method

In the Lewis structure, the central atom is Oxygen (O), which has six valence electrons. Each Hydrogen (H) contributes one electron, resulting in a total of eight electrons. The best structure is one in which two of the oxygen's electrons are shared in covalent bonds with the hydrogen atoms, leaving two lone pairs on the oxygen, resulting a bent geometry.

Valence-Bond Method

Valence bond theory describes the electronic structure of molecules in terms of overlaps of atomic orbitals. For \(H_2O\), the 1s orbital of each hydrogen atom overlaps with the 2p orbital of the oxygen atom, creating sigma bonds. The extra electron pairs on the oxygen atom (2p orbital) provide a region of negative charge which repels the two hydrogen atoms, causing a bent structure.

VSEPR Theory Method

VSEPR (Valence Shell Electron Pair Repulsion) model theory states that electron pairs around a central atom orient themselves so that they can be as far apart as possible. In \(H_2O\), there are two bonding pairs (from the two O-H bonds) and two lone pairs on the central oxygen atom. According to VSEPR theory, this results in a tetrahedral arrangement but the molecule appears bent due to lone pair - bond pair repulsion.

Valence-Bond Method using Hybridized Atomic Orbitals

In water, the central oxygen atom undergoes sp3 hybridization, which is the combination of one 2s and three 2p orbitals, yielding four sp3 hybrid orbitals oriented at the tetrahedral angle (109.5 degrees). Two of these orbitals create sigma bonds with the 1s orbitals of hydrogen atoms while the other two carry the lone pairs. Therefore, considering the repulsion between the pairs of electrons, the molecule takes up a bent geometry.

Key concepts

Lewis Structure

When we talk about the Lewis Structure of a molecule, we're essentially creating a map of how atoms bond and where electrons reside. For water (\( \text{H}_2 \text{O} \)), oxygen sits at the center with its six valence electrons. Hydrogen atoms, each bringing one electron, form bonds by sharing electrons with oxygen. In this structure:

• Two pairs of electrons from oxygen are involved in bonding (these are the bonding pairs).
• The remaining two pairs are lone pairs, which stay on the oxygen atom.

These bond and lone pairs dictate the molecule's shape. In the case of water, they create a bent or V-shaped structure. The crucial takeaway here is that the Lewis Structure provides a straightforward visual of electron distribution, helping us predict molecular geometry.

Valence Bond Theory

The Valence Bond Theory dives into the overlapping of atomic orbitals to explain bonding. Imagine the orbitals as clouds where electrons reside. For \( \text{H}_2 \text{O} \), here's what occurs:

• Oxygen uses its 2p orbitals to overlap with the 1s orbitals of the two hydrogen atoms, forming sigma bonds.
• These bond formations create a cloud overlap rich with electron density.

What's fascinating is how leftover electrons (the lone pairs) in the 2p orbitals influence the molecule's shape. They repel the hydrogen atoms slightly, forcing them into a bent form. Essentially, the bending results from electrons finding equilibrium between repulsive forces and attractive bond formations, emphasizing how crucial orbital overlaps are for molecular geometry.

VSEPR Theory

The VSEPR (Valence Shell Electron Pair Repulsion) Theory explores how molecules take shape based on electron pair repulsion. For water, this is a game of spacing:

• There are two bonding pairs (from the O-H bonds) and two lone pairs on oxygen.
• These pairs repel each other and arrange themselves as far apart as possible.

This configuration initially suggests a tetrahedral shape due to symmetry. However, because lone pairs take up more space and exert greater repulsion, they push the hydrogen atoms closer together, distorting the shape into what's recognized as the bent form. VSEPR theory thus allows us to understand molecular geometry by considering repulsive forces among electron pairs, a simple yet powerful concept.

Hybridization

Hybridization is our way of understanding how atomic orbitals combine to form new, hybrid orbitals used in bond formation. In \( \text{H}_2 \text{O} \), the oxygen atom undergoes sp3 hybridization:

• Oxygen blends one s orbital with three p orbitals, producing four sp3 hybrid orbitals.
• Of these, two form sigma bonds with hydrogen's 1s orbitals, creating the \( \text{O-H} \) bonds.
• The remaining two hybrid orbitals carry lone pairs.

These hybrid orbitals are oriented at about 109.5 degrees, typical for tetrahedral geometry. Yet, the presence of lone pairs bends the molecule, reducing the bond angle to around 104.5 degrees in water. This understanding of hybridized atomic orbitals shows how adjustments in electron distribution can alter and create specific molecular shapes.