YK-2026-Luento_03

the simplest atoms and how their orbitals can combine. When atoms approach each other, their atomic orbitals can overlap, leading to the formation of new molecular orbitals. These molecular orbitals can be either bonding, where electron density is concentrated between the nuclei, or antibonding, where electron density is reduced between the nuclei. The balance between these bonding and antibonding interactions dictates whether a stable chemical bond will form. This fundamental principle is what we refer to as the linear combination of atomic orbitals, or LCAO, which allows us to mathematically describe the shape and energy of these molecular orbitals. in the case hydrogen atoms each contribute a 1s atomic orbital. When these orbitals combine, they form two molecular orbitals: a lower-energy bonding orbital and a higher-energy antibonding orbital. The electrons will preferentially occupy the bonding orbital, leading to a net stabilization and the formation of a covalent bond. The energy released during this bond formation, approximately 430 kJ/mol, is a direct measure of the bond's strength. This release of energy signifies a more stable arrangement, demonstrating why chemical bonds are so prevalent in nature. In more complex scenarios,the overlap can lead to different types of bonds, specifically sigma (σ) and pi (π) bonds. Sigma bonds are formed by head-on overlap of atomic orbitals, resulting in electron density concentrated along the internuclear axis. Pi bonds, on the other hand, arise from the sideways overlap of parallel p-orbitals, leading to electron density above and below the internuclear axis. These different orbital overlaps are critical in determining molecular geometry and reactivity. An important concept in characterizing chemical bonds is the bond order, which is defined as half the difference between the number of electrons in bonding molecular orbitals and antibonding molecular orbitals. For example, in H₂, with two electrons in a bonding orbital and none in an antibonding orbital, the bond order is one, indicating a single covalent bond. In contrast, for a helium molecule, He₂, if it were to form, we would have two electrons in a bonding orbital and two in an antibonding orbital, resulting in a bond order of zero, meaning no stable bond is formed. This explains why helium exists as a monatomic gas rather than a diatomic molecule. Nitrogen atoms contribute 2s and 2p atomic orbitals, which combine to form a series of bonding and antibonding sigma and pi molecular orbitals. The filling of these orbitals according to Hund's rule and the Pauli exclusion principle dictates the molecule's electronic configuration and properties. In N₂, we find a bond order of three, corresponding to a very strong triple bond, which explains nitrogen's high stability and inertness. This strong triple bond consists of one sigma bond and two pi bonds, a characteristic that plays a significant role in nitrogen's chemistry.