To identify the acid, base, conjugate acid, and conjugate base in a reaction, we first refer to the Bronsted-Lowry definitions: an acid is a proton (H⁺) donor, while a base is a proton acceptor. In a typical acid-base reaction, the substance that donates the proton is the acid, and the substance that accepts it is the base. The product formed from the acid after it donates a proton is the conjugate base, while the product formed from the base after it accepts a proton is the conjugate acid.
For example, in the reaction:
HA + B → A⁻ + BH⁺
- HA is the acid because it donates a proton (H⁺).
- B is the base because it accepts the proton (H⁺).
- A⁻ is the conjugate base because it is the remaining part of the acid after the proton is lost.
- BH⁺ is the conjugate acid because it is formed after the base accepts the proton.
Using the periodic table, we can explain the relative acidity and basicity: elements with a higher electronegativity (such as oxygen in HA) are better at stabilizing negative charges, making their conjugate bases (A⁻) more stable. Conversely, elements like nitrogen or alkali metals, which are more basic and less electronegative, tend to readily accept protons due to their less stable electrons.
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