Q.The equilibrium constant for a reaction at room temperature is $K_1$ and that at 700 K is $K_2$. If $K_1 > K_2$, then

Pressure only affects an equilibrium that has gaseous components, and only when the total moles of gaseous reactants differ from the total moles of gaseous products ($\Delta n_g\neq0$). Increasing the pressure (decreasing the volume) shifts the equilibrium toward the side with fewer moles of gas; decreasing the pressure (increasing the volume) shifts it toward the side with more moles of gas. For ammonia synthesis, $N_2(g)+3H_2(g)\rightleftharpoons2NH_3(g)$ (4 mol gas $\rightarrow$ 2 mol gas), compressing the system favours $NH_3$ formation; expanding it favours decomposition back to $N_2$ and $H_2$. When $\Delta n_g=0$, as in $H_2(g)+I_2(g)\rightleftharpoons2HI(g)$, pressure changes have no effect at all on the equilibrium position.

Temperature shifts equilibrium in the direction that absorbs the applied change -- raising the temperature always favours whichever direction is endothermic, and lowering the temperature always favours whichever direction is exothermic. For ammonia synthesis, the forward (synthesis) direction is exothermic, so raising the temperature decomposes some $NH_3$ back to $N_2$ and $H_2$ (the endothermic, heat-absorbing direction), while lowering the temperature favours more $NH_3$ formation. This same logic reads a colour change directly as a sign of $\Delta H$: for $[Co(H_2O)_6]^{2+}(\text{pink})+4Cl^-\rightleftharpoons[CoCl_4]^{2-}(\text{blue})+6H_2O$, if the mixture is blue (forward/product-favoured) at room temperature but turns pink (reverse/reactant-favoured) on cooling, then cooling is favouring the reverse reaction -- so heating must favour the forward reaction, meaning the forward reaction is endothermic, $\Delta H>0$. Comparing a reaction's equilibrium constant at two temperatures gives the same conclusion quantitatively: if $K_1$ (room temperature) is greater than $K_2$ (at 700 K), then K is falling as temperature rises -- exactly the signature of an exothermic forward reaction (a lower K at higher T means the endothermic, i.e. reverse, direction is being favoured as T rises, so the forward direction must be exothermic).

Two important distinctions separate temperature's effect from pressure's and concentration's. First, a pressure or concentration change shifts the equilibrium composition while leaving the numerical value of the equilibrium constant itself unchanged -- but a temperature change genuinely changes the value of K (made quantitative by the Van't Hoff equation). Second, dissolving more gas in cold water is itself a pressure effect in disguise -- Henry's law equilibrium $CO_2(g)\rightleftharpoons CO_2(\text{solution})$ shifts toward more dissolved gas when the partial pressure of $CO_2$ above the liquid is increased (this is why carbonated drinks are bottled under pressure and fizz out once opened, rather than a temperature effect).

Queries like "effect of pressure and temperature on equilibrium Le Chatelier's principle" and "ammonia synthesis equilibrium shift class 11" recur frequently around the Equilibrium chapter of the NCERT/CBSE Class 11 Chemistry curriculum, since this is one of the most heavily tested application areas in JEE Main and NEET chemistry. The colour-change reasoning shown here for reading off the sign of $\Delta H$ is a recurring assertion-reasoning question format in competitive exams.