Radial nodes in an atomic orbital are areas where the probability of finding an electron is zero. The number of radial nodes in an orbital is given by the formula:

$ \text{number of radial nodes} = n - l - 1 $

where $n$ is the principal quantum number and $l$ is the azimuthal quantum number. The azimuthal quantum number ($l$) can have values from 0 to $n-1$, and it determines the shape of the orbital (s, p, d, f, etc.). For an s orbital, $l=0$; for a p orbital, $l=1$; for a d orbital, $l=2$; and so on.

Let's calculate the number of radial nodes for each given orbital:

1. 7s: $n=7$, $l=0$, so the number of radial nodes is $7 - 0 - 1 = 6$, not 5.
2. 7p: $n=7$, $l=1$, so the number of radial nodes is $7 - 1 - 1 = 5$.
3. 6s: $n=6$, $l=0$, so the number of radial nodes is $6 - 0 - 1 = 5$.
4. 8p: $n=8$, $l=1$, so the number of radial nodes is $8 - 1 - 1 = 6$, not 5.
5. 8d: $n=8$, $l=2$, so the number of radial nodes is $8 - 2 - 1 = 5$.

Therefore, the orbitals with 5 radial nodes are 7p, 6s, and 8d, so there are 3 such orbitals.