Multiple choice

Match the column – I (number of radial nodes) with column – II (orbital) and select the correct option.

 
Column – I
Column – II
1. Three A. 3p
2. Two B. 4s
3. One C. 3s
4. Zero D. 2p

  1. 1 - B, 2 - D, 3 – A, 4 - C

  2. 1 - B, 2 - C, 3 – A, 4 - D

  3. 1 - A, 2 - C, 3 – B, 4 - D

  4. 1 - B, 2 - C, 3 – D, 4 - A

  5. 1 - C, 2 - B, 3 – D, 4 - A
Reveal answer Fill a bubble to check yourself
B Correct answer
Explanation

For a given orbital with principal quantum number, n and azimuthal quantum number, l Number of radial nodes = (n - l - 1). For 3s orbital: n = 2 and l = 0. Therefore, number of radial nodes = 3 - 0 - 1 = 2. For 4s orbital: n = 4 and l = 0. Therefore, number of radial nodes = 4 - 0 - 1 = 3. For 2p orbital: n = 2 and l = 1Therefore, number of radial nodes = 2 - 1 - 1 = 0. For 3p orbital: n = 3 and l = 1. Therefore, number of radial nodes = 3 - 1 - 1 = 1.

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