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Chemistry 2e · Fundamental Equilibrium Concepts

Chemical Equilibria

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On this page 9 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Worked example
  6. Key takeaway
  7. Check yourself
  8. Study tools
  9. Sources & references

In 30 seconds

Many chemical reactions are reversible: the products can collide and react to re-form the reactants. When such a reaction runs in a — one that exchanges energy but not matter with its surroundings — the forward and reverse processes do not stop. Instead, they come to a balance. At , the forward and reverse reaction rates are equal, so the concentrations of all reactants and products stop changing even though molecules keep reacting in both directions.

The classic example is dinitrogen tetroxide, a colorless gas, which converts to brown nitrogen dioxide:

N2O4(g) ⇌ 2NO2(g)

Heat a sealed flask and the mixture darkens as more NO2 forms; cool it and the color fades as N2O4 returns. The double arrow ⇌ signals reversibility. This topic establishes what equilibrium is — and what it is not — before the chapter moves on to equilibrium constants, Le Châtelier's principle, and equilibrium calculations.

Why this matters

Equilibrium thinking explains systems that are everywhere: the oxygen–hemoglobin exchange in your blood, the carbon dioxide dissolved in a sealed soft drink, the ammonia produced industrially by the Haber process, and the carbonate chemistry of the oceans. In every case, amounts stop changing not because chemistry has stopped but because opposing processes have balanced.

For exams and later chapters, this topic prevents a costly misunderstanding: equilibrium does not mean "no reaction." Nearly every acid–base, solubility, and redox problem in later chapters is an equilibrium problem in disguise.

The college version

Core Concepts

Reversible reactions and the double arrow

Many reactions written with a single arrow are reversible: under the same conditions, products re-form reactants. Chemists signal this with a double arrow, ⇌ . Examples include the N2O4/NO2 interconversion, esterification and its hydrolysis, and oxygen binding and release by hemoglobin. A reaction that cannot run backward (most combustions) cannot reach equilibrium; it simply runs until a reactant runs out.

Dynamic equilibrium: equal rates, constant amounts

In a closed flask of N2O4 and NO2, both processes — N2O4 → 2NO2 and 2NO2 → N2O4 — occur continuously. Early on, one rate dominates; as concentrations change, the rates move toward each other. At equilibrium they are exactly equal, so for every N2O4 molecule that splits, a pair of NO2 molecules recombines. Net change is zero and amounts are constant. This is dynamic equilibrium: constant at the macroscopic level, busy at the molecular level.

Reaching equilibrium from either direction

The same is reached whether you start with pure N2O4, pure NO2, or any mixture — provided the temperature is the same and the container is closed. Starting with pure N2O4, the forward reaction dominates at first; starting with pure NO2, the reverse reaction dominates. Either way, the system settles at the same final concentrations. This behavior is strong evidence that equilibrium reflects a balance of rates, not the direction in which the reaction was started. The next topic quantifies that balance with the equilibrium constant.

Equilibrium requires a closed system

An open bottle of soda never reaches equilibrium: CO2 escapes, the reverse process (gas dissolving back in) cannot keep up, and the drink keeps losing fizz. Sealed, the same bottle settles into an equilibrium among dissolved CO2, carbonic acid, and headspace gas. The rule: equilibrium concentrations are only constant when matter is neither lost nor added. Energy, however, can cross the boundary — heating a sealed flask changes the equilibrium position, as the color change of N2O4/NO2 demonstrates.

Equilibrium is not limited to chemical reactions

Any process with a measurable reverse can reach equilibrium. A saturated sugar solution in contact with undissolved sugar is an equilibrium: sugar dissolves and crystallizes at equal rates. Liquid water and water vapor in a sealed container at constant temperature form another. Recognizing these as equilibria makes the same rate-balance reasoning useful far beyond reaction equations.

Common Confusions

Do Not ConfuseWithDifference
Equilibrium means the reaction has stoppedDynamic equilibriumForward and reverse reactions continue at equal rates; only the net change is zero.
At equilibrium, concentrations are equalAt equilibrium, concentrations are constantThe equilibrium constant fixes a ratio, not an equality; one side can dominate.
Adding a catalyst changes the equilibrium positionA catalyst changes how fast equilibrium is reachedCatalysts lower activation barriers in both directions; K and the equilibrium composition at fixed temperature are unchanged.
An open container reaches the same equilibriumEquilibrium requires a closed systemWhen products escape, the reverse reaction cannot balance them, so the system keeps shifting.
The double arrow means both directions run equally fastThe double arrow means the reaction is reversibleThe arrow only signals reversibility; the actual balance depends on rates and K.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Imagine a hallway with people walking both ways. At first more people walk one way, but after a while the same number walk each way every minute. Nothing changes, yet people keep moving. A reaction at equilibrium is like that: molecules keep reacting, but the amounts stay the same.

Worked example

Example 1: Reading an equilibrium composition

A sealed 1.00 L flask initially contains 0.100 mol N2O4 and no NO2. At equilibrium, 0.040 mol N2O4 remains. How much NO2 is present?

Write the stoichiometric relationship first: every mole of N2O4 that reacts produces two moles of NO2:

n(NO2) = 2 × (n0(N2O4) - neq(N2O4))

Substitute the given amounts:

n(NO2) = 2 × (0.100 mol - 0.040 mol) = 2 × 0.060 mol = 0.120 mol

The units check: mol  ×  (dimensionless ratio) = mol. Because NO2 is brown, the equilibrium mixture is brown, and its intensity reflects the NO2 concentration. (For practice: the equilibrium constant would be Kc = [NO2]2/[N2O4] = (0.120)2/0.040 = 0.36; the method comes in topic 4.)

Example 2: Forward rate equals reverse rate

For the same system, suppose the forward step is first order with kf = 4.0 × 10-4 s-1 and the reverse step is second order with kr = 2.5 × 10-2 M-1s-1. At equilibrium [N2O4] = 0.040 M. Find [NO2] at equilibrium.

Write the rate expressions first, then set them equal:

rateforward = kf[N2O4]   ratereverse = kr[NO2]2

At equilibrium rateforward = ratereverse:

kf[N2O4] = kr[NO2]2

Substitute the known values:

(4.0 × 10-4 s-1)(0.040 M) = (2.5 × 10-2 M-1s-1)[NO2]2

1.6 × 10-5 M s-1 = (2.5 × 10-2 M-1s-1)[NO2]2

[NO2]2 = 1.6 × 10-5 M s-12.5 × 10-2 M-1s-1 = 6.4 × 10-4 M2

[NO2] = 2.5 × 10-2 M

The unit algebra works out: (M s-1)/(M-1s-1) = M2, and the square root gives M. As a check, K = kf/kr = (4.0 × 10-4)/(2.5 × 10-2) = 0.016, matching [NO2]2/[N2O4] = (2.5 × 10-2)2/0.040 = 0.016. Equal rates really do produce a constant ratio of concentrations.

Key takeaways

  • Dynamic equilibrium means forward and reverse rates are equal; concentrations are constant — not zero, and not necessarily equal.
  • The same equilibrium position is reached from pure reactants, pure products, or any mixture at a given temperature.
  • Equilibrium requires a closed system; open systems that lose matter never settle.
  • A catalyst speeds the approach to equilibrium but does not change the equilibrium position or K at a fixed temperature.
  • For an elementary reversible pair, the equilibrium constant is the ratio of rate constants, K = kforward/kreverse — a preview of the next topic.
  • Macroscopic constancy hides microscopic activity: molecules keep reacting, with no net change.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

  1. What is the defining condition of dynamic equilibrium?

    Show answer

    Forward and reverse reaction rates are equal, so reactant and product concentrations remain constant over time.

  2. Why do concentrations stop changing at equilibrium if molecules are still reacting?

    Show answer

    Because each product molecule formed is matched, on average, by a product molecule reacting back; the net change is zero.

  3. Can equilibrium be reached starting with only products? Explain.

    Show answer

    Yes. The reverse reaction converts products into reactants until forward and reverse rates balance, settling at the same equilibrium position as any other starting mixture at that temperature.

  4. A sealed N2O4/NO2 flask at equilibrium is heated and the mixture darkens. Which species has increased?

    Show answer

    NO2 (brown) has increased; heating shifts this equilibrium toward products.

  5. Does adding a catalyst change the equilibrium concentrations? Why or why not?

    Show answer

    No. A catalyst lowers activation barriers in both directions equally, so it only changes how quickly equilibrium is reached — not the equilibrium concentrations at a given temperature.

  6. Why does an open bottle of soda keep losing CO2 instead of reaching equilibrium?

    Show answer

    CO2 escapes into the room, so the reverse process (gas dissolving back in) cannot balance it; the open system never reaches the closed-system equilibrium state.

Keep learning

Ready to build on this? Continue to the next lesson.

Study tools & related lessonsKey vocabulary · Related

Key vocabulary

reversible reaction
A reaction that can proceed in both forward and reverse directions under the same conditions.
closed system
A system that can exchange energy but not matter with its surroundings.
dynamic equilibrium
A state in which forward and reverse reaction rates are equal, so concentrations are constant over time.
forward reaction / reverse reaction
The reactant-to-product process and the product-to-reactant process.
equilibrium position
The set of concentrations a reaction settles at under given conditions.
equilibrium constant, K
The ratio of product to reactant concentrations at equilibrium (introduced next).
equilibrium constant (K)
A temperature-dependent ratio of product-to-reactant concentrations (or pressures) at equilibrium

Sources & references

  1. openstax.org — Chemistry 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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