2p atomic orbitals 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 (o)bonds, while sideways overlap (typically from adjacent pl orbitals)generates pi (T) bonds. Examples of sigma and T-bond bond formation between atoms a and"b are shown below B The nature of the bonding in hydrogen (H,)can be described using Molecular Orbital Theory. As the two ls 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. The more the orbitals overlap, the more the energy decreases, until the nuclei approach so closely that they begin to repel each. The point at which the repulsive and attractive forces balance defines the bond distance for a given covalent bond. 1s atomic 1s atomic bonding molecular orbital 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 covalentThe 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. Examples of sigma and -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 its own nucleus. The more the orbitals overlap, the more the energy decreases, until the nuclei approach so closely that they begin to repel each. 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