Which of the following statements is/are wrong?

Options

A. Van der waals' radius of iodine is more than its covalent radius
B. all isoelectronic ions belong to same period of the peridic table
C. `I.E_(1)` of N is higher than that of O while `I.E_(2)` of O is higher than that of N
D. The electron affinity N is almost zero while that of P is 74.3 kJ `mol^(-1)`

Explanation

To determine which of the statements is/are wrong, we will analyze each statement one by one.

### Step 1: Analyze the first statement
**Statement:** The van der Waals radius of iodine is more than the covalent radius.

- **Explanation:** The van der Waals radius (VDR) refers to the distance between two non-bonded atoms at their closest approach, while the covalent radius refers to the distance between the nuclei of two bonded atoms. Since bonded atoms are closer together than non-bonded atoms, the van der Waals radius is indeed greater than the covalent radius. Therefore, this statement is **correct**.

### Step 2: Analyze the second statement
**Statement:** Isoelectronic ions belong to the same period of the periodic table.

- **Explanation:** Isoelectronic species are atoms or ions that have the same number of electrons. For example, N³⁻, O²⁻, and F⁻ are isoelectronic anions and belong to the same period (period 2). However, cations like Na⁺ and Mg²⁺, which are also isoelectronic with these anions, belong to the next period (period 3). Hence, not all isoelectronic ions belong to the same period. This statement is **incorrect**.

### Step 3: Analyze the third statement
**Statement:** The ionization energy of N is higher than that of O while the second ionization energy of O is higher than that of N.

- **Explanation:** Nitrogen (N) has an electronic configuration of 2s² 2p³, which is half-filled and stable. Therefore, it requires more energy to remove an electron compared to oxygen (O), which has the configuration 2s² 2p⁴. Thus, the first ionization energy of N is indeed higher than that of O. For the second ionization energy, after losing one electron, N becomes N⁺ (2s² 2p²) and O becomes O⁺ (2s² 2p³). Removing an electron from the half-filled configuration of O⁺ requires more energy than removing from N⁺. Hence, this statement is **correct**.

### Step 4: Analyze the fourth statement
**Statement:** The electron affinity of N is almost 0 while that of phosphorus is 74.3.

- **Explanation:** Nitrogen has a half-filled p subshell (2p³), making it relatively stable and less likely to gain an electron, resulting in an electron affinity close to zero. Phosphorus, on the other hand, has a higher tendency to gain an electron due to its larger size and lower repulsion compared to nitrogen, resulting in a positive electron affinity value of 74.3 kJ/mol. Therefore, this statement is **correct**.

### Conclusion
The only incorrect statement is the second one regarding isoelectronic ions.

### Final Answer
The incorrect statement is **option 2**.

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