energy, are o3s, 3s", T3p, 03p, T3p* and o3p*; in other words, the o2s to o2,* pattern is repeated. Determine the bond order of a Mg, molecule. 2p bond order: 2p Predict the stability of a Mg, molecule. stable 2s 2s 2s O unstable
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.
![The energy-level diagram shown can be continued to higher energies. The next few orbitals, in order of increasing energy, are \( \sigma_{3s}, \sigma_{3s}^*, \pi_{3p}, \sigma_{3p}, \pi_{3p}^*, \) and \( \sigma_{3p}^* \); in other words, the \( \sigma_{2s} \) to \( \sigma_{2p}^* \) pattern is repeated.
**Determine the bond order of a Mg\(_2\) molecule.**
Bond order: \_\_\_\_\_\_
**Predict the stability of a Mg\(_2\) molecule.**
- [ ] stable
- [ ] unstable
### Diagram Explanation
The diagram illustrates the molecular orbital configuration for a Mg\(_2\) molecule:
- **\( \sigma_{1s} \) and \( \sigma_{1s}^* \)**: Two filled orange boxes indicate completely filled bonding and antibonding orbitals from the 1s atomic orbitals.
- **\( \sigma_{2s} \) and \( \sigma_{2s}^* \)**: Two filled orange boxes depict filled orbitals from the 2s atomic orbitals.
- **\( \pi_{2p} \) and \( \pi_{2p}^* \)**: Green rectangles show that the π bonding orbitals are filled, while the antibonding π* orbitals are not.
- **\( \sigma_{2p} \) and \( \sigma_{2p}^* \)**: One filled orange box and one unfilled indicate a filled bonding orbital but an unfilled antibonding orbital.
To find the bond order, use the formula:
\[ \text{Bond order} = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2} \]
The stability of a molecule typically depends on whether the bond order is greater than zero.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fdf94d9f2-cf96-4b11-aec9-277c75004e04%2F2ff8203b-1c25-498f-9e57-7ed919c7fbfe%2Fzfeoif9_processed.jpeg&w=3840&q=75)
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