To determine where an electron with a spin quantum number of +1/2 and a magnetic quantum number of -1 cannot be present, we need to analyze the quantum numbers involved.
### Step-by-step Solution:
1. **Understanding Quantum Numbers**:
- The four quantum numbers are:
- Principal quantum number (n): Indicates the energy level or shell.
- Azimuthal quantum number (l): Indicates the subshell (s, p, d, f).
- Magnetic quantum number (m): Indicates the orientation of the orbital.
- Spin quantum number (s): Indicates the spin of the electron.
2. **Given Values**:
- Spin quantum number (s) = +1/2
- Magnetic quantum number (m) = -1
3. **Determining Possible Values of l**:
- The magnetic quantum number (m) can take values from -l to +l, including 0.
- This means:
- If l = 0 (s subshell), m can only be 0.
- If l = 1 (p subshell), m can be -1, 0, or +1.
- If l = 2 (d subshell), m can be -2, -1, 0, +1, or +2.
- If l = 3 (f subshell), m can be -3, -2, -1, 0, +1, +2, or +3.
4. **Analyzing the Given Magnetic Quantum Number**:
- Since m = -1, the possible values of l that can accommodate this m value are:
- l = 1 (p subshell)
- l = 2 (d subshell)
- l = 3 (f subshell)
5. **Identifying the Subshells**:
- For the s subshell (l = 0), m cannot be -1.
- For the p subshell (l = 1), m = -1 is possible.
- For the d subshell (l = 2), m = -1 is possible.
- For the f subshell (l = 3), m = -1 is possible.
6. **Conclusion**:
- The electron cannot be present in the s subshell because the only possible value for m when l = 0 is 0. Therefore, the correct answer is that the electron cannot be present in the **s subshell**.
### Final Answer:
The electron cannot be present in the **s subshell**.
