State whether the statement is true or false.`BF_(3)` is stronger Lewis acid than `BCl_(3)`.

Answer

To determine whether the statement "BF₃ is a stronger Lewis acid than BCl₃" is true or false, we need to analyze the properties of both compounds in terms of their ability to accept electron pairs.

### Step-by-Step Solution:

1. **Definition of Lewis Acid**:
- A Lewis acid is defined as a compound that can accept a pair of electrons. Therefore, the strength of a Lewis acid is related to its ability to accept electrons.

2. **Structure of BF₃ and BCl₃**:
- Both BF₃ (Boron Trifluoride) and BCl₃ (Boron Trichloride) have a trigonal planar geometry due to sp² hybridization of the boron atom. In both compounds, boron is electron-deficient and can accept electron pairs.

3. **Electronegativity Consideration**:
- Fluorine is more electronegative than chlorine. This means that in BF₃, the fluorine atoms pull electron density away from the boron atom more effectively than the chlorine atoms do in BCl₃.

4. **Positive Charge Density**:
- The greater electronegativity of fluorine results in a higher positive charge density on the boron atom in BF₃ compared to the boron atom in BCl₃. A higher positive charge density indicates that BF₃ is more electron-deficient and thus a stronger Lewis acid.

5. **Comparison of Lewis Acidity**:
- Since BF₃ has a higher positive charge density on boron due to the strong electronegativity of fluorine, it is more capable of accepting electron pairs compared to BCl₃. Therefore, BF₃ is indeed a stronger Lewis acid than BCl₃.

6. **Conclusion**:
- Based on the analysis, the statement "BF₃ is a stronger Lewis acid than BCl₃" is **True**.