Equilibrium

Many reactions are reversible — they proceed in both directions. When such a reaction runs in a closed container, it reaches a state of equilibrium where the rate of the forward reaction equals the rate of the reverse reaction. Equilibrium is dynamic: molecules keep reacting both ways, but the concentrations of reactants and products no longer change. This dynamic-but-constant nature is the first idea NEET checks.

At equilibrium, the concentrations obey the law of mass action, giving the equilibrium constant. For a reaction $aA + bB \rightleftharpoons cC + dD$, $K_c = \dfrac{[C]^c [D]^d}{[A]^a [B]^b}$ in terms of molar concentrations. For gas-phase reactions an equivalent constant $K_p$ uses partial pressures. The two are related by $K_p = K_c (RT)^{\Delta n_g}$, where $\Delta n_g$ is the change in moles of gas — a formula NEET applies often.

The size of $K$ tells you how far a reaction goes: a large $K$ means products are favoured at equilibrium, a small $K$ means reactants dominate. Pure solids and pure liquids are left out of the expression (their 'concentration' is constant), a point that trips students in heterogeneous equilibria.

To predict the direction of a reaction not yet at equilibrium, compute the reaction quotient $Q$ using the same expression as $K$ but with current concentrations. If $Q < K$ the reaction moves forward (toward products); if $Q > K$ it moves backward; and if $Q = K$ it is already at equilibrium. This $Q$-versus-$K$ comparison is a guaranteed NEET tool.

Once a reaction is at equilibrium, what happens if we disturb it? Le Chatelier's principle gives the answer: if a system at equilibrium is subjected to a change in concentration, pressure or temperature, the equilibrium shifts in the direction that opposes (partly undoes) that change. This single principle predicts the outcome of most NEET equilibrium-disturbance questions.

Concentration: adding a reactant pushes the equilibrium toward products (to consume the added reactant); removing a product also pulls it forward. The reverse changes shift it backward. The value of $K$ stays the same — only the position moves.

Pressure (for gases): increasing the pressure (decreasing volume) shifts the equilibrium toward the side with fewer gas moles, to reduce the pressure. If both sides have equal gas moles, pressure has no effect. Adding an inert gas at constant volume does not shift the equilibrium (partial pressures are unchanged) — a classic NEET trap.

Temperature: this is the only factor that actually changes the value of $K$. Raising the temperature favours the endothermic direction (it absorbs the added heat), while lowering it favours the exothermic direction. A catalyst changes neither $K$ nor the equilibrium position — it only helps the system reach equilibrium faster by speeding both directions equally. These rules are applied directly in the industrial Haber process (high pressure, moderate temperature) and Contact process, common NEET contexts.

Acids and bases are described by three theories of increasing scope. The Arrhenius view: an acid gives $\text{H}^+$ and a base gives $\text{OH}^-$ in water. The broader Bronsted-Lowry view: an acid is a proton donor and a base a proton acceptor, which introduces conjugate acid-base pairs that differ by one $\text{H}^+$ (e.g. $\text{HCl}/\text{Cl}^-$). The most general Lewis view: an acid is an electron-pair acceptor and a base an electron-pair donor. NEET regularly asks you to identify conjugate pairs and Lewis acids/bases.

Acids and bases differ in strength. Strong acids/bases ionise completely in water; weak ones ionise only partly, described by a dissociation constant $K_a$ (acids) or $K_b$ (bases). A larger $K_a$ means a stronger acid. For a conjugate pair, $K_a \times K_b = K_w$, linking the strength of an acid to that of its conjugate base.

Water itself ionises slightly: $\text{H}_2\text{O} \rightleftharpoons \text{H}^+ + \text{OH}^-$, with the ionic product of water $K_w = [\text{H}^+][\text{OH}^-] = 1.0\times10^{-14}$ at $25^{\circ}\text{C}$. In pure (neutral) water $[\text{H}^+] = [\text{OH}^-] = 10^{-7}\ \text{M}$.

Acidity is reported on the pH scale: $\text{pH} = -\log[\text{H}^+]$, with $\text{pOH} = -\log[\text{OH}^-]$ and $\text{pH} + \text{pOH} = 14$ at $25^{\circ}\text{C}$. So pH $< 7$ is acidic, $= 7$ neutral, and $> 7$ basic. A change of one pH unit means a tenfold change in $[\text{H}^+]$ — a point NEET likes to test numerically. Being fluent with $\text{pH} = -\log[\text{H}^+]$ and $K_w$ handles most acid-base questions.

A buffer solution resists changes in pH when small amounts of acid or base are added. It is made from a weak acid and its salt (acidic buffer, e.g. $\text{CH}_3\text{COOH} + \text{CH}_3\text{COONa}$) or a weak base and its salt (basic buffer). Buffers are vital biologically — blood is buffered near pH $7.4$, a fact NEET likes to reference.

The pH of an acidic buffer is given by the Henderson-Hasselbalch equation: $\text{pH} = \text{p}K_a + \log\dfrac{[\text{salt}]}{[\text{acid}]}$. When the salt and acid concentrations are equal, $\text{pH} = \text{p}K_a$ and the buffer has its maximum capacity. This equation is a guaranteed NEET calculation.

When a salt dissolves in water, its ions may react with water — salt hydrolysis — affecting the pH. A salt of a strong acid and strong base (e.g. NaCl) gives a neutral solution; a salt of a weak acid and strong base (e.g. $\text{CH}_3\text{COONa}$) is basic; and a salt of a strong acid and weak base (e.g. $\text{NH}_4\text{Cl}$) is acidic. Predicting the nature of a salt solution from its parent acid and base is a frequent NEET question.

For sparingly soluble salts, equilibrium between the solid and its dissolved ions is described by the solubility product $K_{sp}$. A precipitate forms when the ionic product exceeds $K_{sp}$; the solution is just saturated when they are equal. Adding a common ion shifts this equilibrium back, decreasing solubility — the common-ion effect, used to control precipitation and in qualitative salt analysis. Mastering the Henderson equation, salt-nature rules, and the $K_{sp}$/common-ion idea covers the bulk of ionic-equilibrium questions.