Below is the molecular orbital diagram for a diatomic molecule. Which molecule is this and how many bonds form between the atoms? 2p molecule #bonds 1. F₂.2 2. F₂, 1 3. Cl₂, 1 4. Cl₂, 2 5. 02,3 ## #20 # 2p u 2p JJ 2s ##+ 2p -1/₂5
Formal Charges
Formal charges have an important role in organic chemistry since this concept helps us to know whether an atom in a molecule is neutral/bears a positive or negative charge. Even if some molecules are neutral, the atoms within that molecule need not be neutral atoms.
Polarity Of Water
In simple chemical terms, polarity refers to the separation of charges in a chemical species leading into formation of two polar ends which are positively charged end and negatively charged end. Polarity in any molecule occurs due to the differences in the electronegativities of the bonded atoms. Water, as we all know has two hydrogen atoms bonded to an oxygen atom. As oxygen is more electronegative than hydrogen thus, there exists polarity in the bonds which is why water is known as a polar solvent.
Valence Bond Theory Vbt
Valence bond theory (VBT) in simple terms explains how individual atomic orbitals with an unpaired electron each, come close to each other and overlap to form a molecular orbital giving a covalent bond. It gives a quantum mechanical approach to the formation of covalent bonds with the help of wavefunctions using attractive and repulsive energies when two atoms are brought from infinity to their internuclear distance.
#### Explanation of the Diagram
The diagram illustrates the molecular orbital (MO) theory for a diatomic molecule, specifically focusing on the 2s and 2p atomic orbitals. Here’s a breakdown of the components:
1. **2s Orbital Interaction:**
- Each atom contributes a 2s atomic orbital.
- These 2s orbitals combine to form a lower-energy bonding molecular orbital (\(\sigma_{2s}\)) and a higher-energy antibonding molecular orbital (\(\sigma^*_{2s}\)).
- The diagram shows a pair of electrons in \(\sigma_{2s}\) and another pair in \(\sigma^*_{2s}\) orbitals.
2. **2p Orbital Interaction:**
- Each atom also contributes three 2p atomic orbitals.
- These 2p orbitals form bonding (\(\sigma_{2p_z}\) and \(\pi_{2p}\)) and antibonding (\(\sigma^*_{2p_z}\) and \(\pi^*_{2p}\)) molecular orbitals.
- Four electrons occupy the \(\pi_{2p}\) bonding orbitals.
- Four electrons occupy the \(\pi^*_{2p}\) antibonding orbitals.
3. **Molecular Orbital Energy Levels:**
- **Bonding Orbitals:** These occupy a lower energy level and contribute to bond formation.
- **Antibonding Orbitals:** These occupy a higher energy level and can destabilize the molecule if occupied.
#### Bonding Analysis
To determine the number of bonds, we use the formula:
\[ \text{Bond Order} = \frac{1}{2} (\text{Electrons in bonding orbitals} - \text{Electrons in antibonding orbitals}) \]
From the diagram:
- Bonding electrons: \(2 (\sigma_{2s}) + 4 (\pi_{2p}) = 6\)
- Antibonding electrons: \(2 (\sigma^*_{2s}) + 4 (\pi^*_{2p}) = 6\)
\[ \text{Bond Order} =](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Ff41550e9-2d86-4872-bfd8-a28f70a40714%2F3f4563c7-2f3b-4d40-ab85-9102023ae2a5%2F5u2d5pgi_processed.jpeg&w=3840&q=75)
Trending now
This is a popular solution!
Step by step
Solved in 3 steps
