Coaching Centre
Reversible reactions and dynamic equilibrium
Reversible and irreversible reactions, open, closed and isolated systems, static versus dynamic equilibrium, how concentrations and rates change as equilibrium is reached, modelling equilibrium, and non-equilibrium systems such as combustion and photosynthesis.
An unopened bottle of soft drink can sit on a shelf for a year and still fizz when you open it. Inside, carbon dioxide is constantly escaping from the liquid into the gas above it and dissolving back again, at exactly the same rate. Take the cap off and that balance is gone: the gas leaves, nothing comes back, and the drink goes flat. The difference between those two bottles is the whole idea of this module.
Reactions that don't finish
Some reactions go essentially to completion. Burn magnesium ribbon and you end up with white magnesium oxide; no amount of waiting will turn it back into magnesium and oxygen. Many others are reversible: the products can react with each other to re-form the reactants. We show this with a double arrow.
$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \qquad \Delta H = -92 \text{ kJ mol}^{-1}$
The reaction from left to right is called the forward reaction and the one from right to left the reverse reaction. Their enthalpy changes are equal in size and opposite in sign: if the forward reaction releases 92 kJ per mole of equation, the reverse absorbs 92 kJ.
Whether a reaction is "reversible" in practice depends on the conditions. In a sealed container, almost any reaction has some reverse reaction. It becomes effectively irreversible when the reverse reaction is far too slow to notice or when a product escapes.
Open, closed and isolated systems
- An open system exchanges both matter and energy with its surroundings: an open beaker, a Bunsen flame, a living cell.
- A closed system exchanges energy but not matter: a stoppered flask, a sealed syringe, an unopened bottle.
- An isolated system exchanges neither. A good thermos flask is close; truly isolated systems don't exist in the lab.
Chemical equilibrium can only be established in a closed system. In an open system a gaseous product escapes, so the reverse reaction can never keep pace.
Static and dynamic equilibrium
"Equilibrium" means a state where nothing appears to change. There are two very different ways that can happen.
| Static equilibrium | Dynamic equilibrium | |
|---|---|---|
| Rates | Forward and reverse rates are both zero | Forward rate = reverse rate, and neither is zero |
| Particles | Nothing is reacting | Particles keep reacting in both directions |
| Observable (macroscopic) properties | Constant | Constant: colour, pressure, concentration, temperature |
| Chemistry examples | A reaction that has gone to completion (burnt magnesium), or graphite at room temperature, which could in principle become diamond but doesn't at any measurable rate | N₂O₄ ⇌ 2NO₂ in a sealed tube; a saturated salt solution; CO₂ in a sealed bottle |
The key point for exams: at dynamic equilibrium the macroscopic properties are constant while the microscopic processes continue. Concentrations of reactants and products stay constant, but they are generally not equal to each other.
Reaching equilibrium
Mix the reactants in a sealed flask. At the start there are no products, so the reverse rate is zero and the forward rate is at its highest. As reactants are used up the forward rate falls; as products build up the reverse rate rises. When the two rates become equal, concentrations stop changing: equilibrium has been reached. From then on the concentration–time lines are flat and the rate–time lines meet and stay together.
The same equilibrium can be reached from either direction. Start with pure NO₂ in a sealed tube, or pure N₂O₄ containing the same number of nitrogen atoms, keep the temperature the same, and you end with the same mixture.
- Leave the cap on and watch the graph. When do the numbers of dissolved and gaseous CO₂ molecules stop changing? Compare the two rate bars at that point.
- Press Start with all CO₂ as gas. Does the bottle end up with roughly the same split as before? What does that show about equilibrium?
- Warm the bottle to 35 °C and wait. Where is more of the CO₂ now? Explain why warm drinks go flat faster.
- Take the cap off. Why does the system never reach equilibrium now?
Investigating reversibility
Modelling dynamic equilibrium
A classic model uses two 50 mL measuring cylinders. Fill one ("reactants") with 50 mL of water and leave the other ("products") empty. One student uses a 10 mL pipette to move water from the reactant cylinder to the product cylinder; at the same moment a second student uses a 5 mL pipette to move water back. Each pipette is dipped to the bottom, so the volume transferred depends on how much water is in that cylinder, just as the rate of a reaction depends on concentration.
After a few transfers the volumes stop changing, even though water keeps moving both ways. The final volumes are not equal: the faster transfer (the bigger pipette) leaves more water in the products cylinder. That is exactly how a dynamic equilibrium behaves. Like every model it has limits: real particles move continuously and randomly, not in turns, and the transfers aren't caused by collisions.
Non-equilibrium systems
Some important processes never reach equilibrium because they happen in open systems and are strongly driven in one direction.
- Combustion, such as $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$, is highly exothermic and the hot gaseous products spread out into the surroundings. The products have far less chemical energy than the reactants and are dispersed through the air, so both enthalpy and entropy (5B) favour the forward direction overwhelmingly, and in an open flame the products are gone anyway.
- Photosynthesis, $6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$, is endothermic and produces fewer gas molecules than it uses, so neither enthalpy nor entropy favours it. It only proceeds because the plant continuously absorbs light energy and constantly exchanges gases and sugars with its surroundings. A living cell is an open system kept far from equilibrium; when the inputs stop, the cell dies.
Test yourself
Which statement is true for a system at dynamic equilibrium?
Equal, non-zero rates keep concentrations constant.
Calcium carbonate is heated strongly in an open crucible. CaCO₃(s) → CaO(s) + CO₂(g). Why does this reaction not reach equilibrium?
It is an open system. In a sealed container the same reaction does reach equilibrium.
A sealed flask of NO₂ and N₂O₄ has a constant brown colour. Which observation would show that the equilibrium is dynamic rather than static?
Constant properties fit both kinds of equilibrium. Tracers moving between species prove reactions are still happening.
Which is the best description of photosynthesis in a leaf?
It is endothermic and only continues while light is supplied and matter is exchanged.
Label the concentration and rate graphs.
Classify each system.
A reaction that can proceed in both directions is . Equilibrium can only be established in a system. At equilibrium the forward and reverse rates are equal and non-zero, so properties such as colour stay constant. Combustion is an example of a process.
A student seals some colourless N₂O₄ gas in a flask. Over a few minutes the gas turns pale brown, and then the colour stays constant. N₂O₄(g) ⇌ 2NO₂(g). Explain these observations in terms of the rates of the forward and reverse reactions, and sketch the shape of the rate–time graph you would expect.
At the start the flask contains only N₂O₄, so the forward reaction occurs at its maximum rate and the reverse rate is zero. Brown NO₂ forms, so the gas becomes coloured. As N₂O₄ is used up its concentration falls, so the forward rate decreases; as NO₂ builds up, the reverse reaction (2NO₂ → N₂O₄) speeds up. Eventually the two rates become equal. From then on NO₂ is formed as fast as it is used, so its concentration, and the colour, stay constant even though both reactions continue. This is dynamic equilibrium, possible because the flask is a closed system. On a rate–time graph the forward rate starts high and curves down, the reverse rate starts at zero and curves up, and the two meet at a non-zero rate and then continue as one horizontal line.