How a two-way reaction settles down: equilibrium hands-on

How to use this simulation

1. At the start there are only N₂ and H₂. In the upper plot the blue forward rate is at its highest and the red reverse rate is zero. Let it run: blue falls, red rises, and the two meet. That meeting point is equilibrium. In the lower plot the concentrations flatten out.
2. Once settled, compare “Q right now” with “Equilibrium constant Kc” in the readings: they are almost equal (1.00 and 1.00).
3. Now press “Add reactant (N₂)”. N₂ jumps by 1 mol/L, Q drops, and the message at the top says the reaction is shifting forward.
4. Then try heating, raising the pressure and removing product. Each time note which curve jumps and which way the mixture moves.
5. After any disturbance the concentrations take a while to level off. Try giving a second disturbance before the first has settled to see what happens.

Tap open, drain open, water level steady

Picture a bucket. A tap above pours in one litre a minute, and a hole at the bottom lets out exactly one litre a minute. Does the water level change? No. It stays put, yet the water is not standing still: it flows in and flows out.

Chemical equilibrium is exactly like that. The reaction never stops; it runs in both directions. Reactants break down into products and products break down into reactants. When the two rates match, nothing seems to change from outside: colour, pressure and concentrations all hold still. That is why it is called dynamic equilibrium.

One thing to get clear. At equilibrium the concentrations are constant, not equal. The tap and the hole move the same amount of water, but how full the bucket is is a different matter. In the simulation the two rates meet, yet the N₂, H₂ and NH₃ curves settle at different heights.

What equilibrium is, from the beginning

Some reactions are one-way: paper burns to ash and ash never turns back into paper. Many reactions are reversible, so reactants make products and products make reactants again. We show this with the ⇌ sign: N₂ + 3H₂ ⇌ 2NH₃.

Start such a reaction in a closed vessel and at first only the forward reaction runs, because there is no product yet. As product builds up, the reverse rate rises. As reactants get used up, the forward rate falls. Eventually the two rates are equal. That is equilibrium.

The simulation tells exactly this story. The starting mixture is N₂ 1 and H₂ 3 mol/L. At 500 K it settles with N₂ 0.46, H₂ 1.37 and NH₃ 1.09 mol/L, meaning about 54 percent of the starting nitrogen has become ammonia.

Key terms

You need to write these exactly in an exam.

Kc and Q: the arithmetic of equilibrium

Raise each concentration to the power of its coefficient and divide the products by the reactants to get Kc. At one temperature this value is constant whatever the concentrations are. Pure solids and pure liquids are left out.

Why this formula? The forward rate is kf[N₂][H₂]³ and the reverse rate is kr[NH₃]². At equilibrium they are equal, so kf ÷ kr = [NH₃]² ÷ ([N₂][H₂]³) = Kc. The simulation runs on exactly this: kr is set to kf ÷ Kc.

Q has the same formula as Kc, but with the concentrations right now. Comparing Q with Kc tells you which way the reaction will go.

Le Chatelier’s principle: what happens when you push it

Squeeze a spring and let go: it pushes back against you. An equilibrium behaves similarly. Disturb it and the system shifts so as to reduce the effect of the disturbance. The French chemist Henri Le Chatelier stated the principle.

The four disturbances in the simulation have four different effects. The table puts them side by side.

Temperature: the only thing that changes Kc

The Haber reaction has ΔH = −92 kJ, so it gives out heat. Raise the temperature and the system tries to absorb that heat by running backwards. The van ’t Hoff equation gives Kc at each temperature, and the simulation steps by 50 K.

Going from 500 K to 550 K, Kc drops from 1.00 to 0.134, which is 0.13 times its old value. The whole range:

Try it yourself in the simulation

Run these one at a time, writing a prediction before each.

• Experiment 1: do nothing and let it run. The rates meet and [NH₃] settles at 1.09 mol/L. In the readings Q and Kc are 1.00 and 1.00.
• Experiment 2: at equilibrium press “Add reactant”. Q becomes 0.314, below Kc, so the reaction goes forward. At the new equilibrium [NH₃] rises to 1.29.
• Experiment 3: press “Remove half the product”. NH₃ falls from 1.09 to 0.54 and Q is 0.250. The new equilibrium brings [NH₃] back to 0.72.
• Experiment 4: press “Raise the pressure”. Every concentration doubles at once and Q suddenly falls to 0.25 of its value, 0.250. The shift is forward because 4 moles of gas become 2.
• Experiment 5: press “Raise the temperature”. Kc falls, so Q > Kc and NH₃ starts to break down. Both rates jump at once, so equilibrium arrives faster than before.

Solved problems

Every figure here comes from the simulation’s own model, so you can check it against the screen.

Mistakes everybody makes

These cost the most marks in exams.

• Thinking the concentrations are equal at equilibrium: constant does not mean equal.
• Thinking the reaction has stopped: both rates are equal and neither is zero.
• Thinking pressure changes Kc: only temperature does.
• Thinking a catalyst raises the yield: it only gets you to equilibrium sooner.
• Forgetting to raise concentrations to the power of the coefficient: with 3H₂ you need [H₂]³.
• Putting a solid in the Kc expression: pure solids and liquids are left out.

Equilibrium in real life

The principle runs from factories to our own bodies.

• Fertiliser plants: the Haber process makes ammonia from the nitrogen and hydrogen in air, then urea. Plants use high pressure and a catalyst, and take the ammonia away as it forms so the equilibrium keeps moving forward.
• Fizzy drinks: while the bottle is sealed, CO₂ gas and dissolved CO₂ sit at equilibrium. Open the cap, the pressure drops, the equilibrium shifts and gas escapes as bubbles.
• Oxygen in blood: haemoglobin binding oxygen is reversible. In the lungs oxygen is plentiful so it binds; in tissues it is scarce so it is released.
• Sugar in hot water: dissolved and undissolved sugar are in equilibrium. Heating dissolves more, because dissolving takes in heat.
• Colour-change demos: some ion solutions change colour when warmed or cooled, a direct view of an equilibrium shifting.

Exam corner (CBSE, NCERT class 11 and others)

Try these on your own, then check the answers.

• Short answer: define dynamic equilibrium. State Le Chatelier’s principle. What does a catalyst do to equilibrium?
• Numerical: at equilibrium N₂ 0.5, H₂ 1.5 and NH₃ 0.4 mol/L. Find Kc. Answer: 0.095.
• For the reverse reaction with the same data, Kc = 1 ÷ Kc = 10.5.
• If the equation is halved to ½N₂ + 3⁄2H₂ ⇌ NH₃, the new Kc is the square root of Kc, 0.308.
• A mixture has N₂ = H₂ = NH₃ = 1 mol/L, so Q = 1.00. With Kc = 1.00, which way does it go? Q smaller than Kc means forward; larger means reverse.

Revision: at a glance

Read this list the night before the exam.

• At equilibrium the two rates are equal, but the reaction does not stop.
• Kc = product concentrations (to coefficient powers) ÷ reactant concentrations (to coefficient powers).
• Q < Kc: forward. Q > Kc: reverse. Q = Kc: equilibrium.
• Add reactant or remove product: forward. The opposite: reverse.
• More pressure shifts towards fewer moles of gas; more heat shifts the endothermic way.
• Only temperature changes Kc; a catalyst only makes equilibrium arrive sooner.

Frequently asked questions