For **d-orbitals** or d-subshell, Ɩ = 2, there are five values of m namely -2, -1, 0, 1, 2. It means d- orbitals can have five orientations. They have an even more complex angular distribution than the p orbitals. These are represented by d_{xy}, d_{yz}, d_{zx}, d_{x2-y2 }and d_{z2}; for example, 3d_{xy}, 3d_{yz}, 3d_{zx}, 3d_{x2-y2} and 3d_{z2}. The d_{xy}, d_{yz,} and d_{zx} orbitals have the same shape i.e., cloverleaf shape but they lie in XY, YZ, and ZX planes respectively. Hence, we can say that there are five d-orbitals. These different orbitals essentially have different orientations. The d_{z2} orbital is symmetrical about Z-axis and has a dumbbell shape with a doughnut-shaped electron cloud in the centre. The d_{x2-y2} orbital is also clover leaf-shaped but its leaves are directed along the X and Y-axis.

The reason for the presence of four lobes in any d orbital lies in the fact that the d – orbitals have two nodes, and hence two changes in algebraic sign of ψ, which lead to four lobes. d orbitals have two angular nodes (two angles at which the probability of electron is always zero).

**Fig: Shapes of d-orbitals**

There is a set of five d orbitals (with complicated shapes and names) as well as the 3s and 3p orbitals (3px, 3py, 3pz). At the third level, there is a total of nine orbitals altogether. The magnetic orbital quantum number for d orbitals is given as (-2,-1,0, 1,2). Out of these five d orbitals, shapes of the first four d-orbitals are related to each other, which is different from the dz2 orbital whereas the energy of all five d orbitals is the same.