Which of the following combination of orbitals can not form bond. (If x axis in internuclear axis)

Options

A. `s+p_(z)`
B. `s+s`
C. `p_(z)+p_(x)`
D. `d_(xy)+p_(y)`

Explanation

To determine which combinations of orbitals cannot form a bond, we will analyze each option based on the type of overlap that occurs between the orbitals. The key types of overlaps are head-on overlap (which can form sigma bonds) and sidewise overlap (which can form pi bonds).

### Step-by-Step Solution:

1. **Identify the Combinations**: We have four combinations of orbitals to analyze:
- Option 1: s + pz
- Option 2: s + s
- Option 3: pz + px
- Option 4: dxy + py

2. **Analyze Option 1: s + pz**:
- The s orbital is spherical and can overlap with the pz orbital (which is oriented along the z-axis) only in a head-on manner. However, the pz orbital cannot effectively overlap with the s orbital in a sidewise manner.
- **Conclusion**: This combination cannot form a bond.

3. **Analyze Option 2: s + s**:
- Two s orbitals can overlap in a head-on manner. This overlap is effective and allows for bonding.
- **Conclusion**: This combination can form a bond.

4. **Analyze Option 3: pz + px**:
- The pz and px orbitals are oriented along different axes (z and x respectively). They cannot overlap effectively in a head-on manner, nor can they form a sidewise overlap that would allow for bonding.
- **Conclusion**: This combination cannot form a bond.

5. **Analyze Option 4: dxy + py**:
- The dxy orbital has a specific orientation in the xy-plane, while the py orbital is oriented along the y-axis. These orbitals can overlap effectively, allowing for bonding.
- **Conclusion**: This combination can form a bond.

### Final Answer:
The combinations of orbitals that cannot form a bond are:
- Option 1: s + pz
- Option 3: pz + px

Thus, the answer is **1 and 3**.