Atomic & Molecular Orbitals.
Electrons surrounding atoms are concentrated into regions of space called atomic orbitals. The Heisenberg uncertainty principle states that it is impossible to know both the location and the momentum of an atomic particle, but it is possible to describe the probability that the electron will be found within a given region of space. The boundaries of an atomic orbital are commonly drawn to the region of 90% probability; there is a 90% probability that at any given time, the electron will be within the specified boundary. The electronic configuration of carbon is 1s2 2s2 2sp3. Atomic orbitals with s-character have spherical symmetry.
The wave properties of electrons make the description of the 2s orbital slightly more complex than the corresponding 1s orbital, in that, within the 2s sphere there is a region in which the amplitude of the electron standing wave falls to zero, that is, there is zero probability of finding the electron in this node region.
The electron densities along the x, y and z axes of the 2p orbitals are clearly shown in the figure; the nodes are the points at the origin and at these points, there is zero probability of finding the electron. The sharing of electrons in a covalent bond occurs by overlap of the individual atomic orbitals. Head-on overlap between energetically compatible orbitals generates sigma bonds, while sideways overlap (typically from adjacent p orbitals) generates pi bonds. The nature of the bonding in hydrogen (H2) can be described using Molecular Orbital Theory. As the two 1s atomic orbitals approach each other and begin to overlap, there is a decrease in the net energy of the system because the electrons in each atom tend to become attracted to the positive nucleus of the other atom, as well as its own nucleus.
