The number of angular and radial nodes for 4d orbital is respectively
- A.2 and 1
- B.1 and 2
- C.3 and 1
- D.4 and 0
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Explanation
The correct answer is A, 2 and 1. For a 4d orbital, n = 4 and l = 2, which gives l = 2 angular nodes and n − l − 1 = 1 radial node.
A node is a region where the probability of finding the electron is zero. Angular nodes are planes or cones through the nucleus, and their number equals the azimuthal quantum number l. Radial nodes are spherical shells, and their number is n − l − 1. The total number of nodes is always n − 1. For 4d: l = 2 for any d orbital, so 2 angular nodes; radial nodes = 4 − 2 − 1 = 1; total = 3, which equals n − 1. In the same way 3d has 2 angular and 0 radial nodes, and 2s has 0 angular and 1 radial node.
B is wrong: it reverses the two values.
C is wrong: 3 angular nodes belong to an f orbital (l = 3), not a d orbital.
D is wrong: 4 angular nodes would need l = 4, a g orbital, and the total of 4 exceeds n − 1 = 3.
Exam tip: angular nodes = l, radial nodes = n − l − 1, total nodes = n − 1.