Acids, Bases and Salts — Core Principles

Core Principles

Acids, bases, and salts are fundamental chemical classifications. Acids are substances that typically donate protons ($H^+$) or accept electron pairs, often characterized by a sour taste and a pH less than 7.

Bases are substances that accept protons or donate electron pairs, usually feeling slippery and having a pH greater than 7. Salts are ionic compounds formed from the neutralization reaction between an acid and a base, consisting of a cation from the base and an anion from the acid.

Key theories defining these include Arrhenius (based on $H^+$ and $OH^-$ in water), Brønsted-Lowry (proton donors/acceptors), and Lewis (electron-pair acceptors/donors). The pH scale quantifies acidity/basicity, with $pH = -log[H^+]$.

The strength of an acid or base is determined by its extent of dissociation, quantified by $K_a$ or $K_b$. Salts can undergo hydrolysis in water, leading to acidic, basic, or neutral solutions depending on the strengths of their parent acid and base.

Buffer solutions, composed of a weak acid/base and its conjugate, resist pH changes, playing vital roles in biological and chemical systems.

Important Differences

vs Acid-Base Theories

Aspect	This Topic	Acid-Base Theories
Definition of Acid	Arrhenius Theory: Produces $H^+$ (or $H_3O^+$) in water.	Brønsted-Lowry Theory: Proton ($H^+$) donor.
Definition of Base	Arrhenius Theory: Produces $OH^-$ in water.	Brønsted-Lowry Theory: Proton ($H^+$) acceptor.
Solvent Requirement	Arrhenius Theory: Restricted to aqueous solutions.	Brønsted-Lowry Theory: Not restricted to aqueous solutions.
Examples of Acids	$HCl$, $H_2SO_4$	$HCl$, $H_2SO_4$, $NH_4^+$
Examples of Bases	$NaOH$, $KOH$	$NaOH$, $KOH$, $NH_3$, $CO_3^{2-}$
Scope	Narrowest, only for substances producing $H^+$ or $OH^-$.	Broader, includes reactions without $OH^-$ formation, introduces conjugate pairs.

The three major acid-base theories—Arrhenius, Brønsted-Lowry, and Lewis—represent an evolution in chemical understanding, each expanding the scope of what constitutes an acid or a base. Arrhenius is the most restrictive, focusing on $H^+$ and $OH^-$ ions in water. Brønsted-Lowry broadens this to proton transfer, introducing the concept of conjugate pairs and allowing for non-aqueous reactions. Lewis theory is the most encompassing, defining acids as electron-pair acceptors and bases as electron-pair donors, which explains a wider range of reactions, including those without proton transfer, like the formation of coordination compounds. For NEET, understanding all three is crucial for classifying diverse chemical species.