To determine which of the following is a relatively strong Lewis acid, we need to analyze the properties of the boron halides, particularly BF3, BCl3, BBr3, and BI3. Here's a step-by-step solution:
### Step 1: Understand Lewis Acids
A Lewis acid is defined as a substance that can accept an electron pair. This typically occurs when the substance has an incomplete octet or is electron-deficient.
### Step 2: Analyze Boron Halides
Boron (B) has three valence electrons. When it forms compounds with halogens (F, Cl, Br, I), it can form BX3 (where X is a halogen). In these compounds, boron does not complete its octet, making it electron-deficient.
### Step 3: Evaluate the Electron Deficiency
- **BF3**: Boron is bonded to three fluorine atoms. Fluorine is highly electronegative, which means it pulls the bonding electrons towards itself. This increases the electron deficiency on boron, making BF3 a strong Lewis acid.
- **BCl3**: Chlorine is less electronegative than fluorine, so the electron deficiency on boron is less pronounced than in BF3.
- **BBr3**: Bromine is even less electronegative, leading to a further decrease in the electron deficiency on boron.
- **BI3**: Iodine is the least electronegative of the halogens, resulting in the least electron deficiency on boron compared to the other halides.
### Step 4: Compare the Lewis Acidity
As we move down the group from fluorine to iodine, the electronegativity decreases, which leads to a decrease in the electron-withdrawing effect on boron. Therefore, the order of Lewis acidity among these compounds is:
- BF3 > BCl3 > BBr3 > BI3
### Conclusion
Thus, BF3 is the strongest Lewis acid among the boron halides due to the highest electron deficiency on boron caused by the strong electronegativity of fluorine.
### Final Answer
The relatively strong Lewis acid among the options provided is **BF3**.
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