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 and a representation of the surface of the carbon 1s orbital is shown below.

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. Nodes are most easily seen in the description of the 2p atomic orbitals, which are shown below.

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 (s) bonds, while sideways overlap (typically from adjacent p orbitals) generates pi (p) bonds. Examples of sigma and p-bond bond formation between atoms "A" and "B" are shown below.

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 their own nucleus. The more the orbitals overlap, the more the energy decreases, until the nuclei approach so closely that they begin to repel each other. The point at which the repulsive and attractive forces balance defines the bond distance for a given covalent bond.

In molecular orbital theory, the number of atomic orbitals used to make the covalent bond must equal the total number of molecular orbitals in the molecule. In the example cited above, the atomic orbitals combine to form one bonding orbital, containing the two electrons, and one high-energy antibonding orbital which is empty. The molecular orbital description of this simple covalent bonding is shown below. As described above, the bonding orbital is referred to as a s-orbital, while the corresponding antibonding orbital is referred to as a s*-orbital.

In a similar manner, sideways overlap of adjacent p-orbitals forms a covalent p-orbital and a corresponding high-energy p*-antibonding molecular orbital. In general, electrons only populate antibonding orbitals when the molecule is in an excited state, and such orbitals are typically ignored in the discussion of organic reaction mechanisms. A more useful description of bonding is often given by the Valence Shell Electron Pair Repulsion (VSEPR) model, in which electrons are positioned around an atom to minimize electrostatic repulsion. This concept is described in more detail in the following section on orbital hybridization.