Question

Match Column - I with Column - II.

Column - I	Column - II
1. The orbital with 4 total nodes	a. $3s$
2. The orbital with one nodal plane	b. $5s$
3. The orbital with 2 nodal regions	c. $5d$
4. The orbital with no nodal plane	d. $4p$

Answer 1

Total number of nodal planes in an orbital is equal to azimuthal quantum number$\left(l\right)$.

The number of spherical or radial nodes in an orbital=$(n-l-1)$.

1. The orbital with 4 total nodes is $5s$ as $n=1$ and $l=0$ therefore number of nodes=$(5-0-1)=4$ i.e option b.

2. The orbital with one nodal plane is $4p$ as number of nodal plane is equal to $l=1$ i.e option d.

3. The orbital with two nodal region is $5d$ as number of nodal region is equal to $l=2$ i.e option c.

The orbital with no nodal plane is $3s$ as for$3s$ $l=0$ i.e option a.

Therefore, correct sequence is $1-b,2-d,3-c,4-a$