Definition
A reversible reaction can proceed in both directions () — products can reform the reactants under the same conditions.
Dynamic equilibrium: forward and reverse reaction rates become equal, so concentrations of reactants and products stop changing — but both reactions keep happening (unlike a static, unchanging state).
Remember
Combustion reactions are generally irreversible because their products (usually gases that escape, or very stable low-energy compounds) don't readily reform the original reactants under normal conditions.
Rule
Increasing a reactant's concentration shifts equilibrium toward the products (forward); removing a product also pulls the reaction forward.
Increasing pressure shifts equilibrium toward the side with fewer gas molecules; increasing temperature shifts it toward the endothermic direction.
A catalyst speeds up both forward and reverse reactions equally — it helps equilibrium get established faster, but does not change the equilibrium position or the yield.
Example 1 — a fizzy drink can
Explain why gas bubbles out when you open a can of fizzy drink.
- Dissolved is in equilibrium with gas under the can's pressurized seal.
- Opening the can drops the pressure suddenly, shifting the equilibrium toward the gas side (fewer gas molecules more gas molecules isn't favoured by pressure, so lowering pressure favours the side with more gas molecules).
Answer: Releasing pressure shifts equilibrium toward releasing gas
Quick Recap
- ●Reversible reaction: can go both ways; dynamic equilibrium: forward/reverse rates equal, concentrations stop changing
- ●More reactant or removing product: shifts forward. More pressure: favours fewer gas molecules. More heat: favours the endothermic direction
- ●A catalyst speeds up reaching equilibrium but doesn't change its position or yield
- ●Combustion is typically irreversible — its products don't reform the reactants easily
