Equilibrium
Chemical equilibrium is the dynamic balance point where reactions continue in both directions at equal rates, producing no net change in reactant and…
Start with the simplest version: this lesson is about Equilibrium. If you can explain the core idea to a friend using everyday language, examples, and one clear reason why it matters, you have moved from memorising to understanding.
Chemical equilibrium is the dynamic balance point where reactions continue in both directions at equal rates, producing no net change in reactant and product amounts. Understanding equilibrium explains why reactions don't always go to completion, how shifting conditions affects outcomes, and why your blood maintains constant pH despite continuous chemical activity.
What is Equilibrium? Dynamic vs. Static
When a reaction reaches equilibrium, it appears static—amounts of reactants and products stop changing. But on a molecular level, the reaction continues vigorously in both directions. Molecules constantly form and break apart; we simply can't detect net change because forward and reverse rates balance.
Consider the reaction: N₂O₄(g) ⇌ 2NO₂(g). Brown nitrogen dioxide forms from colorless dinitrogen tetroxide. Heat the system: the brown color intensifies (equilibrium shifts right, producing more NO₂). The reaction hasn't "stopped"—it's shifted to a new equilibrium point where both reactions occur at the same rate.
This is dynamic equilibrium: molecules react continuously, but macroscopic properties remain constant.
Physical Equilibria: Phase Changes
Physical equilibria involve no chemical bonds breaking:
When liquid water evaporates in a closed container, molecules escape the liquid surface. Simultaneously, gas molecules strike the surface and condense back to liquid. Eventually, evaporation and condensation rates balance. The liquid level stays constant while molecules continuously exchange between phases.
This explains why sealed containers have constant vapor pressure at constant temperature: equilibrium is reached between evaporation and condensation. This principle is why perfume in an open bottle eventually evaporates completely (no equilibrium established—gases escape to atmosphere) but perfume in a sealed bottle reaches equilibrium (not all evaporates).
Chemical Equilibria: Reaction Balance
Chemical equilibria occur when reactant and product amounts stabilize. Consider:
2NO₂(g) ⇌ N₂O₄(g)
If we start with pure NO₂, some converts to N₂O₄. As N₂O₄ forms, it begins decomposing back to NO₂. Eventually, rates balance: NO₂ forms from N₂O₄ as fast as NO₂ combines into N₂O₄. Concentrations stabilize.
Crucially, equilibrium is reached regardless of starting point. Start with pure N₂O₄, pure NO₂, or a mixture—the system always reaches the same equilibrium point (at constant temperature and pressure). This predictability is because equilibrium depends only on the system's conditions, not its history.
The Equilibrium Constant: Quantifying Balance
The equilibrium constant (K) quantifies how far a reaction proceeds before reaching equilibrium. It's the ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients.
For: aA + bB ⇌ cC + dD
K = [C]^c[D]^d / [A]^a[B]^b
[X] denotes concentration in moles per liter at equilibrium.
For the NO₂/N₂O₄ reaction: K = [N₂O₄]/[NO₂]²
Large K (say, K = 1000) means products dominate at equilibrium—the reaction proceeds nearly to completion.
Small K (say, K = 0.001) means reactants dominate—the reaction barely proceeds.
K = 1 means roughly equal concentrations of reactants and products.
K varies with temperature but not with pressure or initial concentrations. This immutability is powerful: measure K once at a given temperature, use it to predict equilibrium compositions under any starting conditions.
Le Chatelier's Principle: Responding to Stress
When a system at equilibrium is "stressed" (conditions changed), it shifts to counteract that stress. This is Le Chatelier's Principle.
Temperature change: If an exothermic reaction is heated, equilibrium shifts left (reverse direction produces heat, cooling the system). If an endothermic reaction is heated, equilibrium shifts right (forward reaction absorbs heat).
Pressure change: If a reaction has more gas molecules as products than reactants (like N₂ + 3H₂ ⇌ 2NH₃, left side has 4 molecules, right has 2), increasing pressure shifts equilibrium right (fewer molecules mean lower pressure—the system "tries" to reduce pressure).
Concentration change: If you add reactants, equilibrium shifts right to consume the added material. If you remove products, equilibrium also shifts right to replace them. The system opposes the change.
Catalyst effect: Catalysts don't shift equilibrium—they speed both forward and reverse reactions equally. Equilibrium position remains unchanged; it's reached faster.
These shifts aren't mysterious. They're consequences of how K works: if you disturb a system, reactions shift until K is re-established at the new conditions.
Applications of Equilibrium
Gas Phase Reactions: The Haber process (N₂ + 3H₂ ⇌ 2NH₃) produces ammonia for fertilizers. High pressure and low temperature favor products, but low temperature slows reaction. Industry uses moderate conditions, catalysts, and continuous removal of ammonia to shift equilibrium right.
Solubility Equilibria: When salt dissolves, it reaches equilibrium: NaCl(s) ⇌ Na⁺(aq) + Cl⁻(aq). The equilibrium constant is called the solubility product (Ksp). Knowing Ksp, you can predict whether a precipitate forms when solutions mix.
Acid-Base Equilibria: Water ionizes slightly: H₂O ⇌ H⁺ + OH⁻. This equilibrium determines pH. Buffers work by maintaining pH through equilibrium between weak acids and their conjugate bases.
Biological Systems: Hemoglobin binds oxygen reversibly: Hb + O₂ ⇌ HbO₂. In lungs (high oxygen pressure), the reaction shifts right (oxygen binds). In tissues (low oxygen pressure), it shifts left (oxygen releases). Your body's oxygen transport is fundamentally an equilibrium process.
Quantitative Calculations: Finding Equilibrium Compositions
Given K and initial concentrations, you can calculate equilibrium concentrations using an ICE table (Initial, Change, Equilibrium):
For: A ⇌ B, starting with 1.0 M of A and K = 0.4
| State | A | B |
|---|---|---|
| Initial | 1.0 | 0 |
| Change | -x | +x |
| Equilibrium | 1.0-x | x |
At equilibrium: K = [B]/[A] = x/(1.0-x) = 0.4
Solving: x = 0.4(1.0-x) → x = 0.4 - 0.4x → 1.4x = 0.4 → x ≈ 0.29
So [B] ≈ 0.29 M and [A] ≈ 0.71 M at equilibrium.
This method extends to complex reactions, allowing prediction of exact equilibrium compositions.
Socratic Questions
- Why must equilibrium be "dynamic" rather than static? What would it mean physically if the forward and reverse reactions actually stopped occurring?
- If you reach the same equilibrium position regardless of starting point, what determines where that position is? What property of the system creates this invariance?
- Why does increasing pressure shift equilibrium in a reaction with different numbers of gas molecules on each side, but a catalyst doesn't shift equilibrium at all?
- In hemoglobin binding oxygen, the equilibrium shifts right in lungs and left in tissues. Why does the oxygen "know" which direction to shift based on its concentration?
- If a reaction has K = 0.01 (strongly favoring reactants), why would a chemist still be interested in it? When might such unfavorable reactions be useful?