Do electron-withdrawing groups make carboxylic acids weaker?
Short answer: -I groups like Cl stabilise the carboxylate anion, so the acid gets stronger.
The trap
The wrong idea
“Electron-withdrawing groups make carboxylic acids weaker.”
It sounds reasonable, and a lot of JEE aspirants carry it into the exam. Here's a one-question check, then the explanation that makes it click.
Quick check
Try it before you peek
Which is the stronger acid?
- CH₃COOH
- ClCH₂COOH
- Same strength
Show the answer
B. ClCH₂COOH
-I groups like Cl stabilise the carboxylate anion, so the acid gets stronger.
The click
It is the opposite: electron-withdrawing groups make carboxylic acids stronger. Acid strength depends on how stable the conjugate base is once the H⁺ has left. A more stable carboxylate anion (RCOO⁻) means the acid gives up its proton more readily.
Chlorine is more electronegative than carbon, so in ClCH₂COOH it pulls electron density through the σ bonds (a −I effect). This spreads out the negative charge on the carboxylate, stabilising ClCH₂COO⁻ compared with CH₃COO⁻. As a result, chloroacetic acid (pKa ≈ 2.9) is roughly 80 times stronger than acetic acid (pKa ≈ 4.76). A lower pKa means a stronger acid.
The effect grows with the number and strength of the withdrawing groups: Cl₃CCOOH (pKa ≈ 0.7) > Cl₂CHCOOH > ClCH₂COOH > CH₃COOH. A more electronegative atom has a bigger effect, so FCH₂COOH (pKa ≈ 2.6) is stronger than ClCH₂COOH. The inductive effect fades with distance, so 2-chlorobutanoic acid is a stronger acid than 4-chlorobutanoic acid.
Electron-donating groups work the other way: alkyl groups (+I) destabilise the carboxylate and weaken the acid, which is why formic acid, HCOOH (pKa ≈ 3.75), is stronger than acetic acid. Rule of thumb: whatever stabilises the anion strengthens the acid.
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