A catalyst is part of the reaction, otherwise it wouldn't have any effect. The thing that makes it a catalyst and not an additional reaction is that it is not used up.
A reaction works through a certain mechanism. This mechanism forms an intermediate. This intermediate is higher in energy. So a reaction does need activation energy, but it doesn't come from the catalyst. It comes from the temperature. When the temperature gets higher, the odds for a single atom with more than average kinetic energy become better. These are the molecules that react.
But if the activation energy is too high, if the intermediate is too unstable and too high in energy, the reaction will not happen, even if it is exothermic/gives off heat.
Catalysts lower activation energy by stabilizing the intermediate step, making the intermediate lower in energy. The activation energy is the initial rise in energy needed to activate the reaction. The molecule needs to go up the hill in energy before it can go downhill on the other side.
Often a catalyst forms a complex with the reactant. Reactant is now more likely to become the product. When that happens, the catalyst is released unchanged.
So reactant A reacts with catalyst C; A + C -> AC
AC is the complex of reactant and catalyst. This is the intermediate, which is high in energy and unstable.
AC - > B + C
Or it will just go back down the same slope of the hill where it just went up: AC -> A +C
In either case, activation energy is released and the catalyst neither adds or releases energy.
If a possible 'catalyst' were to add energy, that energy would be used up and it wouldn't be a catalyst.