An electron has a spin quantum number `+1(1)/(2)` and a magnetic quantum number -1. It cannot be present in

Options

A. d-orbital
B. f-orbital
C. p-orbital
D. s-orbital

Explanation

To solve the problem, we need to analyze the given quantum numbers of the electron: the spin quantum number \( s = +\frac{1}{2} \) and the magnetic quantum number \( m = -1 \). We will determine which orbital this electron cannot occupy based on these quantum numbers.

### Step-by-Step Solution:

1. **Understanding Quantum Numbers**:
- The spin quantum number \( s \) indicates the intrinsic spin of the electron. For electrons, \( s \) can be either \( +\frac{1}{2} \) or \( -\frac{1}{2} \).
- The magnetic quantum number \( m \) indicates the orientation of the orbital and can take values from \( -l \) to \( +l \), where \( l \) is the azimuthal quantum number.

2. **Identifying the Orbital Types**:
- **s orbital**: For an s orbital, \( l = 0 \). Therefore, the only possible value for \( m \) is 0.
- **p orbital**: For a p orbital, \( l = 1 \). The possible values for \( m \) are -1, 0, and +1.
- **d orbital**: For a d orbital, \( l = 2 \). The possible values for \( m \) are -2, -1, 0, +1, and +2.
- **f orbital**: For an f orbital, \( l = 3 \). The possible values for \( m \) are -3, -2, -1, 0, +1, +2, and +3.

3. **Analyzing the Given Quantum Numbers**:
- The electron has \( m = -1 \).
- Since \( m \) can take the value of -1 in both p and d orbitals, it is permissible for these orbitals.
- However, for the s orbital, since \( l = 0 \), the only possible value for \( m \) is 0. Therefore, \( m = -1 \) cannot be accommodated in an s orbital.

4. **Conclusion**:
- The electron with the given quantum numbers cannot be present in an s orbital. Thus, the correct answer is that the electron cannot be present in the **s orbital**.

### Final Answer:
The electron cannot be present in the **s orbital**.