4.5.3. Overlapping of Atomic Orbitals

When orbitals of two atoms come close to form bond, their overlap may be positive, negative or zero depending upon the sign (phase) and direction of orientation of amplitude of orbital wave function in space. Positive and negative sign on boundary surface diagrams show the sign (phase) of orbital wave function and are not related to charge.

The concept of overlap refers to how atomic orbitals interact when two atoms approach each other to form a bond. Depending on how these orbitals align, they can either reinforce each other (positive overlap) or cancel each other out (negative overlap). When orbitals overlap in a way that they both have a similar phase (the same sign), it results in a stronger bond, a phenomenon known as positive overlap.

Imagine two people trying to hug each other. If they approach each other and align their arms in a way that they connect well, they create a stronger hug (positive overlap), just like orbitals with positive overlap creating a strong bond. If one person crosses their arms while the other reaches out, they won't connect as well, just like orbitals in negative overlap.

Orbitals forming bond should have same sign (phase) and orientation in space. This is called positive overlap.

For a bond to form effectively, the orbitals involved must have not only the same phase but also the right orientation. This means that the way they point and align in space matters because it allows for maximum interaction. If two orbitals do not align well, their ability to form a stable bond is compromised.

Think of it like a game of catch: if one person throws the ball to another, and they are facing each other correctly, the catch is more likely to succeed (positive overlap). However, if one person is turned away, or if the throw is poorly aimed, the likelihood of making a connection is reduced.

The criterion of overlap, as the main factor for the formation of covalent bonds applies uniformly to the homonuclear/heteronuclear diatomic molecules and polyatomic molecules. We know that the shapes of CH4, NH3, and H2O molecules are tetrahedral, pyramidal and bent respectively. It would be therefore interesting to use VB theory to find out if these geometrical shapes can be explained in terms of the orbital overlaps.

The concept of overlap is crucial not just for diatomic molecules but extends to polyatomic compounds as well. The shapes of molecules like methane (CH4), ammonia (NH3), and water (H2O) can also be understood through the lens of overlap. The distinct shapes arise from how the atomic orbitals of the constituent atoms overlap, thereby influencing the angles and orientations between bonds.

Consider arranging furniture in a room. The way you position each piece (like chairs and tables) affects how the room feels. If everything is well-placed, the room has a good flow (like CH4’s tetrahedral shape). If things are not positioned correctly, movement may feel restricted (like the angles in NH3 and H2O). This illustrates how orbital overlap determines molecular geometry.

Let us first consider the CH4 (methane) molecule. The electronic configuration of carbon in its ground state is [He]2s2 2p2 which in the excited state becomes [He] 2s1 2px1 2py1 2pz1. The energy required for this excitation is compensated by the release of energy due to overlap between the orbitals of carbon and hydrogen.

The formation of methane serves as a prime example of how orbital overlap leads to molecule formation. When carbon atoms are excited, they rearrange their electrons into a state where 4 orbitals (from the two s and two p orbitals) can overlap effectively with the 1s orbitals of four hydrogen atoms. This overlap is crucial as it allows for the formation of four strong C-H bonds, creating the molecule's characteristic tetrahedral geometry.

Think of preparing a salad: you have different ingredients (orbitals). If you mix and toss them properly (overlap), you create a delicious dish (a stable molecule). Just as proper mixing ensures all flavors blend well, effective overlapping of orbitals ensures strong bonding in molecules.

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