If you increase the temperature, the endothermic reaction is favored. The Le Chatelier's principle states that the endothermic reaction is favored in order to minimize the effect of an increase in temperature. Endothermic reactions absorb heat so the temperature must, in a way, remain constant in that system. So that way, the equilibrium constant is never changed due to a change in temperature as the temperature doesn't change. I know I went wrong somewhere but can someone explain where. Please don't include equations like the ones with Gibbs free energy and Van't hoff equations and stuff as we didn't learn all of that in school yet. Please make the explanation as intuitive as possible.
The right way to do this is to understand the Van't Hoff equation.
Barring that, I can provide some intuition, but it'll have to be a bit hand wavy.
The way I think about this is that at higher temperatures, the intrinsic preference for reactants or products becomes less. That's because the thermal energy of the system is becoming more relevant than the energy difference between reactants and products.
For an exothermic reaction, when we increase the temperature, we'll reduce some of the preference for the product, thereby pushing the reaction backwards.
For an endothermic reaction, when we increase the temperature, we'll reduce some of the preference for reactant, thereby pushing the reaction forwards.
Of course, it's pretty hard to truly justify these kinds of quantitative relationships without the use of proper mathematical tools. In this case, the relationships are most clearly captured in the Van't Hoff equation.