1. Why does the distance between two nuclei in a covalent bond vary?
2. How does a molecular orbital differ from an atomic orbital?
For #2, check this site.
http://en.wikipedia.org/wiki/Orbital_hybridisation
The bond distance between two atoms varies for at least two reasons:
a. The size of the two atoms.
b. The number of electrons shared.
1. The distance between two nuclei in a covalent bond can vary due to several factors:
a. Atomic size: The size of atoms is determined by the number of electrons and the electron configuration. As you move across a period on the periodic table, the atomic size decreases because the effective nuclear charge increases. Thus, when two atoms come together to form a covalent bond, the distance between their nuclei can vary based on their atomic sizes.
b. Bond strength: The strength of a covalent bond depends on the number of shared electrons and the bond length. Generally, shorter bond lengths imply stronger bonds. When two atoms come closer together, the bond strength increases, resulting in a decrease in the bond length.
c. Electronegativity: Electronegativity is the ability of an atom to attract electron density towards itself in a chemical bond. When two atoms with different electronegativities bond, there will be a partial positive charge on the less electronegative atom and a partial negative charge on the more electronegative atom. This creates an attractive force that can affect the bond length.
d. Bond angle: In some molecules, the atoms are bonded in a way that forms an angle between the three atoms. The bond angle can affect the distance between the nuclei. For example, in a bent molecule such as water (H2O), the bond angle of approximately 104.5 degrees causes the nuclei to be closer together compared to a linear molecule like carbon dioxide (CO2).
2. A molecular orbital (MO) and an atomic orbital (AO) are different in terms of their nature and how they are formed:
a. Nature: An atomic orbital exists in an atom and is associated with a single atom's electronic structure. It describes the probability of finding an electron around the nucleus of that atom. In contrast, a molecular orbital exists in a molecule and describes the probability of finding electrons within the entire molecule, not just around a specific atom. MOs are formed by the combination (overlap) of AOs from different atoms.
b. Formation: Atomic orbitals are formed by the combination of wavefunctions that satisfy the Schrödinger equation for the electron in an isolated atom. They have specific shapes (s, p, d, etc.) and energy levels. Molecular orbitals are formed when atomic orbitals overlap, either constructively (forming a bonding orbital) or destructively (forming an antibonding orbital) along the bonding axis. The combination of AOs results in the formation of new MOs with adjusted energies.
c. Energy: Atomic orbitals have discrete energy levels based on their primary quantum numbers. They are relatively localized around the atomic nucleus. In contrast, molecular orbitals have a continuum of energy levels. They extend throughout the entire molecule and can accommodate electrons from multiple atoms.
d. Electron occupancy: Atomic orbitals are occupied by electrons in accordance with the Pauli exclusion principle and the Aufbau principle. Each atomic orbital can hold a maximum of two electrons with opposite spins. Molecular orbitals, on the other hand, can be occupied by electrons from different atoms, forming bonding or antibonding combinations. The number of electrons in the molecular orbitals depends on the number of atomic orbitals involved in the combination.
Overall, the formation and properties of molecular orbitals are more complex than atomic orbitals as they describe the electronic behavior of a molecule as a whole rather than individual atoms.