In chemistry, reactions can be:
Irreversible (one way), orReversible (two-way).An irreversible reaction is a reaction that effectively goes in one direction only :
Reactants → Products Once the reactants have formed products, they do not easily change back into the original substances. These reactions often release a lot of energy and produce very stable products . $$\mathrm{CH}_4(\mathrm{~g})+2 \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})+2 \mathrm{H}_2 \mathrm{O}$$
Methane burns in oxygen to form carbon dioxide and water . Under normal conditions, CO₂ and H₂O do not react together to reform methane and oxygen. The reaction is practically one-way → irreversible . $$2 \mathrm{Mg}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{MgO}(\mathrm{~s})$$
Magnesium burns in air with a bright white flame to form magnesium oxide . Magnesium oxide is very stable; it does not spontaneously decompose back into Mg and O₂. The reaction is effectively irreversible in normal conditions. A reversible reaction can go in both directions :
Reactants ⇌ Products The products can react again to re-form the original reactants . Both the forward and reverse reactions $$\mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{HI}(\mathrm{~g})$$
Forward reaction: H₂ and I₂ combine to form HI.Reverse reaction: HI decomposes back into H₂ and I₂.If this reaction is carried out in a closed container , both reactions occur, and over time the system reaches a state where:
H₂ and I₂ are still reacting to form HI. HI is still decomposing back into H₂ and I₂. But the concentrations of all three gases become constant . This state is called dynamic equilibrium .
Irreversible reactions Proceed in one direction only . Products are usually very stable and do not easily revert to reactants. Often release a large amount of energy and go to completion. Reversible reactions Can proceed in both forward and reverse directions. Products can react to reform reactants . Can reach a state of dynamic equilibrium in a closed system , wherethe rate of the forward reaction = the rate of the reverse reaction , and the concentrations of reactants and products remain constant (but not necessarily equal). Reversible reactions are written with a double arrow , e.g. ⇌ . Irreversible reactions are written with a single arrow , e.g. → . A general reversible reaction is written as: $$\mathrm{A}+\mathrm{B} \rightleftharpoons \mathrm{C}+\mathrm{D}$$
Forward reaction: $$\mathrm{A}+\mathrm{B} \to \mathrm{C}+\mathrm{D}$$Reverse reaction: $$\mathrm{C}+\mathrm{D} \to \mathrm{A}+\mathrm{DB}$$The double arrow (⇌) tells us the reaction can go both ways .
We can represent the dynamic nature of reversible reactions with arrows and rates .
At the start (only reactants present):Forward reaction rate is high (long arrow).Reverse reaction rate is low (short arrow or almost zero), because there are almost no products yet. As time passes:Reactants are used up → forward rate decreases . Products are formed → reverse rate increases . At dynamic equilibrium :The rates of the forward and reverse reactions become equal . Arrows can be drawn the same length to show this. The system looks unchanged at the macroscopic level, but particles are still reacting both ways. Key point: At equilibrium, rates are equal, but concentrations are constant (not necessarily equal).
Rates of forward and reverse reactions become equal at dynamic equilibrium. $$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$
Forward: Nitrogen + hydrogen → ammonia. Reverse: Ammonia → nitrogen + hydrogen. In a closed system , an equilibrium mixture is formed containing N₂, H₂, and NH₃ . At equilibrium:Ammonia is being formed and decomposed at the same rate . The mixture of gases remains constant in composition. $$\mathrm{N}_2 \mathrm{O}_4(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})$$
N₂O₄ is colourless . NO₂ is brown . At equilibrium in a closed tube:Both gases are present. The overall colour depends on the relative amounts of N₂O₄ and NO₂. Changing conditions (like temperature) can shift the balance and make the mixture more brown (more NO₂) or more colourless (more N₂O₄). This is a nice visual model for equilibrium shifts. Hydrated cobalt(II) chloride: CoCl₂·6H₂O (pink). Anhydrous cobalt(II) chloride: CoCl₂ (blue). Dehydration (heating): $$\mathrm{CoCl}_2 \cdot 6 \mathrm{H}_2 \mathrm{O}(\text { pink }) \rightleftharpoons \mathrm{CoCl}_2(\text { blue })+6 \mathrm{H}_2 \mathrm{O}(\mathrm{~g})$$ Heating drives the reaction to the right (removing water). Adding water drives the reaction back to the left (re-forming the hydrated salt). These reversible colour changes are used in humidity indicators and desiccant (drying) packs . For a reversible reaction in a closed system equilibrium mixture . The position of equilibrium tells us whether the mixture contains mostly reactants , mostly products , or significant amounts of both . Le Chatelier’s Principle states:
If a system at equilibrium is disturbed by a change in conditions, the system will adjust to oppose the change and establish a new equilibrium. We’ll focus on:
Concentration Temperature (and briefly) Pressure and catalysts For a general reaction:
$$\mathrm{A}+\mathrm{B} \rightleftharpoons \mathrm{C}+\mathrm{D}$$
Increase [A] or [B] (add reactant) The system responds by using up extra reactant . Equilibrium shifts to the right , making more products C and D.Decrease [C] or [D] (remove product) The system responds by making more product . Again, equilibrium shifts to the right . Increase [C] or [D] (add product) The system responds by using up extra product . Equilibrium shifts to the left , making more A and B.Decrease [A] or [B] (remove reactant) The system responds by reforming reactants . Equilibrium shifts to the left .Haber process
$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$
Increasing [N₂] or [H₂] → shifts equilibrium to the right , producing more NH₃. Removing NH₃ as it forms also shifts equilibrium to the right , favouring ammonia production. Temperature changes affect which direction is favoured , depending on whether the forward reaction is exothermic or endothermic Exothermic reaction (releases heat): $$\text { Reactants ⇌ Products + heat }$$Endothermic reaction (absorbs heat): $$\text { Reactants + heat ⇌ Products }$$ Add heat (increase temperature): The system will try to use up the extra heat . Equilibrium shifts in the endothermic direction . Remove heat (decrease temperature): The system will try to produce heat . Equilibrium shifts in the exothermic direction . Haber process (exothermic forward reaction)
$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})+\text { heat }$$
Increasing temperature → shifts equilibrium to the left , reducing ammonia yield. Decreasing temperature → shifts equilibrium to the right , increasing ammonia yield (but slows the rate). This is why industrial conditions are a compromise : moderately high temperature for a reasonable rate , but not so high that the yield is too low.
For reactions involving gases, changing the pressure (by changing volume) can affect the equilibrium position.
Le Chatelier’s Principle:
Increasing pressure → equilibrium shifts to the side with fewer gas molecules .Decreasing pressure → equilibrium shifts to the side with more gas molecules .Haber process
$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$
Left side: 1 + 3 = 4 moles of gas. Right side: 2 moles of gas. Thus:
Increasing pressure → shifts equilibrium to the right (towards fewer moles of gas) → more ammonia. Decreasing pressure → shifts equilibrium to the left . A catalyst :
Speeds up both the forward and reverse reactions equally. Helps the system reach equilibrium faster . Does not change the position of equilibrium or the final equilibrium composition. In your own words, what is the difference between a reversible and an irreversible reaction? Write the reversible reaction for the Haber process with the correct symbol. What does “dynamic equilibrium” mean in this context? For the reaction $$\mathrm{N}_2 \mathrm{O}_4(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})$$ predict what happens to the colour of the gas mixture if:more NO₂ is added, the mixture is cooled. In the Haber process, why are high pressure and moderate temperature used in industry?