Q.Predict, using Le Chatelier's principle, the effect of increasing the total pressure (by decreasing the volume) on the position of the equilibrium $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$.

Both perturbations — adding a reactant and removing a product — push the equilibrium $2H_2(g) + CO(g) \rightleftharpoons CH_3OH(g)$ towards the product side, per Le Chatelier's principle.

The equilibrium is:

$2H_2(g) + CO(g) \rightleftharpoons CH_3OH(g)$

Le Chatelier's principle: if a system at equilibrium is subjected to a change (in concentration, pressure, temperature, etc.), the equilibrium shifts in the direction that tends to counteract (undo) that change.

(i) Addition of H$_2$: Increasing the concentration of a reactant (H$_2$) disturbs the equilibrium. To counteract (reduce) this added H$_2$, the equilibrium shifts in the forward direction (towards the product, CH$_3$OH), consuming the extra H$_2$ (along with CO) and producing more methanol. Net effect: more CH$_3$OH is formed.

(ii) Removal of CH$_3$OH: Removing some of the product (CH$_3$OH) decreases its concentration below the equilibrium value. To counteract this loss and restore some of the removed product, the equilibrium again shifts in the forward direction, converting more $H_2$ and $CO$ into $CH_3OH$. Net effect: the reaction proceeds further forward, and more CH$_3$OH is continuously formed (this is, in fact, how continuously removing product can be used to drive a reaction to near completion).

Le Chatelier's principle governs how a system at equilibrium responds to a disturbance: it shifts to partially counteract the change.

Le Chatelier's principle states that if a system at equilibrium is subjected to a change in concentration, pressure, volume, or temperature, the equilibrium shifts in the direction that tends to counteract (partly undo) that change. The main factors are:

1. Concentration: Adding more of a reactant shifts equilibrium towards the products (to consume the extra reactant); removing a product (or adding more product) similarly shifts the equilibrium to compensate.

2. Pressure/volume (relevant for gas-phase equilibria involving unequal moles of gas on the two sides): Increasing pressure (decreasing volume) shifts equilibrium towards the side with fewer moles of gas, since that side occupies less volume; decreasing pressure shifts it towards the side with more moles of gas. If the moles of gas are equal on both sides, pressure/volume changes have no effect on the position of equilibrium.

3. Temperature: Increasing temperature shifts equilibrium in the endothermic direction (the direction that absorbs the added heat); decreasing temperature shifts it in the exothermic direction.

4. Catalyst: A catalyst speeds up both the forward and reverse reactions equally, so it helps the system reach equilibrium faster, but it does not change the position of equilibrium (the equilibrium constant Kc/Kp is unaffected).

5. Addition of an inert gas: at constant volume, adding an inert gas does not change the partial pressures of the reacting species, so equilibrium is unaffected; at constant pressure, adding inert gas increases the total volume, which is equivalent to a pressure decrease and can shift equilibrium towards the side with more moles of gas.

Le Chatelier's principle: the system counteracts any imposed change; both actions here push the equilibrium forward (toward more CH3OH).

For the equilibrium $2H_2(g) + CO(g) \rightleftharpoons CH_3OH(g)$:

1. Addition of H2: Increasing the concentration of a reactant (H2) disturbs the equilibrium. By Le Chatelier's principle, the system shifts in the direction that consumes the added reactant — i.e., the forward direction — to partially offset the increase. This produces more CH3OH and consumes more CO.
2. Removal of CH3OH: Removing a product decreases its concentration below the equilibrium value. The system responds by shifting in the direction that replenishes the removed product — again the forward direction — so more H2 and CO react to form more CH3OH. In both cases, the equilibrium shifts to the right (forward direction), increasing the amount of CH3OH formed.

By Le Chatelier's principle, higher temperature reduces NH3 yield (since the forward reaction is exothermic) while higher pressure increases NH3 yield (since the forward reaction decreases the total moles of gas).

The reaction N2(g) + 3H2(g) ⇌ 2NH3(g) (the Haber process) is exothermic in the forward direction (ΔH < 0), and the forward reaction reduces the total number of gas moles from 4 (1 + 3) to 2.

Effect of temperature: Since the forward reaction releases heat, by Le Chatelier's principle, raising the temperature shifts the equilibrium in the direction that absorbs the added heat — i.e., backward, toward N2 and H2 — decreasing the yield of NH3. Lowering the temperature favours the forward (exothermic) reaction and increases NH3 yield, although at very low temperature the rate of reaction becomes too slow, so industrially a moderate temperature (around 700 K) is used as a compromise.

Effect of pressure: Since the forward reaction decreases the number of gas moles (4 mol reactants → 2 mol product), increasing the total pressure shifts the equilibrium in the direction that reduces the number of gas moles — i.e., forward, toward NH3 — increasing its yield. Decreasing pressure favours the reverse reaction and lowers NH3 yield. Industrially, high pressure (around 200 atm) is used to maximise NH3 yield.