As an expert in the field of organic chemistry, I can provide a detailed explanation as to why monochloroacetic acid is more acidic than its non-chlorinated counterpart, acetic acid.

Firstly, it is important to understand the concept of acidity in organic compounds. Acidity in organic chemistry is generally measured by the pKa value of the compound, which is the negative logarithm of the acid dissociation constant (Ka). The lower the pKa value, the stronger the acid. Monochloroacetic acid has a lower pKa value than acetic acid, indicating that it is a stronger acid.

The increased acidity of monochloroacetic acid can be attributed to the presence of the chlorine atom. Chlorine is a halogen and is known to be a strong electron-withdrawing group (EWG). When chlorine is attached to the carbon chain of an organic molecule, it exerts a significant inductive effect. This inductive effect is due to the electronegativity of chlorine, which is higher than that of carbon. As a result, chlorine pulls electron density away from the carbon chain, including the carbon atom that is bonded to the hydroxyl group (-OH) in the case of monochloroacetic acid.

The inductive effect leads to a decrease in electron density around the oxygen atom of the hydroxyl group. This makes it easier for the oxygen atom to lose a proton (H+), thus increasing the tendency of the molecule to act as an acid. In other words, the chlorine atom facilitates the ionization of the hydroxyl group, making monochloroacetic acid a stronger acid than acetic acid.

Furthermore, the presence of the chlorine atom also affects the stability of the conjugate base that is formed after the acid donates a proton. The conjugate base of monochloroacetic acid is the chloroacetate ion. The negative charge on this ion is delocalized through resonance with the chlorine atom. Resonance structures can be drawn where the negative charge is shared between the oxygen and chlorine atoms. This delocalization of the negative charge over a larger area stabilizes the conjugate base, making it less likely to recombine with the proton to reform the acid. Consequently, the acid dissociation is more favorable, and the acid is stronger.

It is also worth noting that the chlorine atom attached to the beta carbon of monochloroacetic acid can participate in a neighboring group effect. This means that the chlorine atom can stabilize the negative charge on the alpha carbon through a mechanism known as anchimeric assistance. However, this effect is not as significant in the case of monochloroacetic acid as it is in other compounds with a beta-chlorine atom.

In summary, the increased acidity of monochloroacetic acid compared to acetic acid is primarily due to the electron-withdrawing inductive effect of the chlorine atom, which facilitates the ionization of the hydroxyl group and stabilizes the conjugate base through resonance. This results in a lower pKa value and a stronger acidic character for monochloroacetic acid.

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Chlorine being a strong electron withdrawing group, pulls the negative charge towards itself by inductive effect so that the negative charge density on the oxygen atom is reduced, hence stabilizing the conjugate base of chloroacetic acid. ... Chloroacetic acid has a chlorine atom attached to its beta carbon.read more >>