A catalyst speeds up a reaction by providing an alternative reaction pathway that has a lower activation energy than the uncatalysed route. Because a greater fraction of molecules possess energy exceeding this reduced barrier at a given temperature, both the forward and reverse rates increase, and the reaction reaches equilibrium faster. The catalyst participates in intermediate steps but is regenerated, so it is not consumed overall. The option of increasing activation energy describes a negative effect and would slow the reaction. Raising temperature is a separate factor and is not how a catalyst works at constant temperature. Shifting the equilibrium is incorrect because a catalyst speeds both directions equally and leaves the equilibrium position and the value of K unchanged. This mechanism, illustrated by energy-profile diagrams, is a central NCERT concept. Plausibility check: a lower barrier necessarily increases the rate constant through the Arrhenius factor while leaving the thermodynamic equilibrium untouched, consistent with the chosen answer.