MO Theory: Homonuclear Diatomic Molecules: Videos & Practice Problems

Understanding molecular orbital theory is essential when studying homonuclear diatomic molecules, which consist of two identical atoms, such as N₂ or F₂. The molecular orbital diagrams differ between elements based on their period in the periodic table. For period 1 elements like hydrogen and helium, valence electrons reside in the 1s orbital, leading to the formation of bonding (σ_1s) and antibonding (σ*_1s) molecular orbitals. As we move to period 2 elements, from lithium to nitrogen, the 2s and 2p orbitals come into play, creating a more complex arrangement with bonding (σ_2s, π_2p) and antibonding (σ*_2s, π*_2p) molecular orbitals. However, for oxygen to neon, due to increased electronegativity and decreased atomic size, the order of σ_2p and π_2p orbitals is reversed, impacting the molecular orbital structures. This concept is crucial in predicting the behavior and bonding of molecules.

Homonuclear Diatomic Molecules

A whole nuclear diatomic molecule consists of two identical elements bonded together, such as nitrogen (N₂) or fluorine (F₂). The molecular orbital theory explains how the arrangement of electrons in these molecules changes based on the elements involved. As electronegativity increases and atomic size decreases, the order of molecular orbitals for elements from oxygen to neon is reversed.

In molecular orbital diagrams, we first consider hydrogen (H₂) and helium (He₂). Both hydrogen and helium have their valence electrons in the 1s orbital, which is the first period of the periodic table. The bonding molecular orbital is represented as σ_1s, while the antibonding molecular orbital is denoted as σ_1s*.

When examining diatomic lithium (Li₂) through diatomic nitrogen (N_22s, followed by the antibonding molecular orbital σ_2s* and then the 2p orbitals. For the 2p orbitals, the order of filling is π_2p, σ_2p, π_2p*, and σ_2p*.

As we move from diatomic oxygen (O₂) to diatomic neon (Ne₂), the order of the molecular orbitals changes due to the increased electronegativity and decreased atomic size. In this case, π_2p moves to a higher energy level, while σ_2p moves to a lower energy level. This shift is crucial for understanding the stability and bonding characteristics of these molecules.

Overall, the arrangement of molecular orbitals varies depending on the period of the elements involved, influencing the properties and behaviors of homonuclear diatomic molecules. Understanding these concepts is essential for predicting molecular stability and reactivity in various chemical contexts.

Homonuclear Diatomic Molecules MO Diagram Example

To determine the electron configuration for the difluorine anion, \( \text{F}_2^{-1} \), we start by calculating the total number of valence electrons. Each fluorine atom, being in group 7A, has 7 valence electrons. Therefore, for two fluorine atoms, we have:

\( 7 \, \text{(from F)} \times 2 = 14 \, \text{valence electrons} \)

Since the anion has gained one additional electron, the total becomes:

\( 14 + 1 = 15 \, \text{valence electrons} \)

Next, we construct the molecular orbital (MO) diagram. The electrons fill the molecular orbitals in the following order based on their energy levels. For period 2 elements like fluorine, the filling starts from the \( 2s \) orbital:

1. The \( \sigma_{2s} \) orbital can hold 2 electrons (filled with 2 electrons).
2. The \( \sigma^*_{2s} \) orbital can hold 2 electrons (filled with 2 electrons).
3. The \( \sigma_{2p} \) orbital can hold 2 electrons (filled with 2 electrons).
4. The \( \pi_{2p} \) orbitals can hold a total of 4 electrons (2 in each of the two degenerate \( \pi \) orbitals).
5. The \( \pi^*_{2p} \) orbitals can hold a total of 4 electrons (2 in each of the two degenerate \( \pi^* \) orbitals).
6. Finally, the \( \sigma^*_{2p} \) orbital can hold 1 electron (filled with 1 electron).

Now, we can summarize the electron configuration in the molecular orbital notation:

\( \sigma_{2s}^2 \, \sigma^*_{2s}^2 \, \sigma_{2p}^2 \, \pi_{2p}^4 \, \pi^*_{2p}^4 \, \sigma^*_{2p}^1 \)

This notation indicates the distribution of the 15 electrons across the molecular orbitals. The filled molecular orbital diagram for the difluorine anion reflects the stability and bonding characteristics of the molecule, showcasing how the additional electron influences its electronic structure.

Determine the number of electrons found in the π_2p orbitals for the dioxygen dication, O₂²⁺.

• A. 2
• B. 0
• C. 4
• D. 3

Answer: (C)

Using a MO diagram, write the electron configuration for the P₂ molecule

• A. (σ_2s)²(σ*_2s)²(π_2p)⁴(σ_2p)²
• B. (σ_2s)²(σ*_2s)²(π_2p)²(σ_2p)²(π*_2p)²
• C. (σ_3s)²(σ*_3s)²(π_3p)⁴(σ_3p)²
• D. (σ_3s)²(σ*_3s)²(π_3p)²(σ_3p)²(π*_3p)²

Answer: (C)

Here’s what students ask on this topic:

What is a homonuclear diatomic molecule?

A homonuclear diatomic molecule consists of two identical atoms bonded together. Examples include N₂ (nitrogen gas) and F₂ (fluorine gas). These molecules are characterized by having the same element in both atomic positions, leading to symmetrical molecular structures. Understanding these molecules is crucial in molecular orbital theory, as the bonding and antibonding interactions between the identical atoms' orbitals determine the molecule's stability and properties.

How do molecular orbital diagrams differ for period 1 and period 2 elements?

For period 1 elements like hydrogen (H₂) and helium (He₂), the valence electrons are in the 1s orbital, forming bonding (σ1s) and antibonding (σ*1s) molecular orbitals. In period 2 elements, such as lithium (Li₂) to nitrogen (N₂), the 2s and 2p orbitals are involved, creating a more complex arrangement with bonding (σ2s, π2p) and antibonding (σ*2s, π*2p) molecular orbitals. For oxygen (O₂) to neon (Ne₂), the order of σ2p and π2p orbitals is reversed due to increased electronegativity and decreased atomic size, impacting the molecular orbital structures.

Why is the order of σ2p and π2p orbitals reversed for oxygen to neon?

The order of σ2p and π2p orbitals is reversed for oxygen (O₂) to neon (Ne₂) due to increased electronegativity and decreased atomic size. In these elements, the increased nuclear charge pulls the 2p orbitals closer to the nucleus, lowering their energy. This causes the σ2p orbital to have a lower energy than the π2p orbital, reversing their order compared to elements like nitrogen (N₂). This reversal affects the molecular orbital diagram and the bonding properties of these molecules.

What are bonding and antibonding molecular orbitals?

Bonding molecular orbitals are formed when atomic orbitals combine constructively, resulting in a lower energy state that stabilizes the molecule. Examples include σ1s, σ2s, and π2p orbitals. Antibonding molecular orbitals, denoted with an asterisk (e.g., σ*1s, σ*2s, π*2p), are formed when atomic orbitals combine destructively, resulting in a higher energy state that destabilizes the molecule. Electrons in bonding orbitals contribute to the bond strength, while electrons in antibonding orbitals weaken the bond.

How do you fill molecular orbitals for homonuclear diatomic molecules?

To fill molecular orbitals for homonuclear diatomic molecules, start with the lowest energy orbitals and move upwards. For period 1 elements, fill the σ1s orbital first, followed by the σ*1s orbital. For period 2 elements, fill the σ2s orbital, then the σ*2s orbital, followed by the π2p and σ2p orbitals. For oxygen to neon, the order is reversed, so fill the σ2s, σ*2s, σ2p, and then the π2p orbitals. Always follow the Pauli exclusion principle and Hund's rule when filling these orbitals.