Acids, Bases, and pH

SECTIONS

	1.	Introduction
	2.	Acids and Bases
	3.	The pH Scale
	4.	Strong and Weak Acids
	5.	Acid-Conjugate Base Pairs
	6.	The Henderson- Hasselbalch Equation
	7.	Buffers
	8.	Review

ACIDS AND THEIR CONJUGATE BASES

Remember the Bronsted–Lowry definition of an acid: an acid is a proton donor. Clearly, in our previous example acetic acid donates its proton to H₂O in the forward reaction, resulting in the formation of the acetate ion. However, in the reverse reaction, the acetate ion accepts a proton from H₃O⁺ to revert to CH₃COOH. From the Bronsted–Lowry definition, a base is a proton acceptor.

CH₃COOH	+	H₂O		CH₃COO^–	+	H₃O⁺
acetic acid		water		acetate ion		hydronium ion

Thus, the acetate ion acts as a base in the reverse reaction, and is called the conjugate base of acetic acid. Acetic acid and the acetate ion are known as an acid–base conjugate pair. Interestingly enough, there is a second acid–base conjugate pair in the reaction above. Notice how in the reverse reaction, H₃O⁺ donates a proton (making it an acid), while in the forward reaction, H₂O accepts a proton (making it a base). H₂O and H₃O⁺ are a conjugate acid–base pair, with H₃O⁺ being the conjugate acid and H₂O being the conjugate base.

In a conjugate acid–base pair, the conjugate acid has one more proton than its corresponding conjugate base.