Opposite effects of Le Chatelier's Principle How does Le Chatelier's Principle deal with a change in reaction conditions which can simultaneously effect the equilibrium in opposite ways.
For example, here's the equation for the conversion of $\ce{N2}$ and $\ce{H2}$ into $\ce{NH3}$. It's an exothermic reaction, so I've added heat as a product on the right side.

$$\ce{N2(g) + 3H2(g) &lt;=&gt; 2NH3(g) +\text{heat}}$$

If I increase the temperature (add heat) that would tend to push the reaction to the left because the reaction to the right is an exothermic reaction. But if I add heat (and the volume remains constant), then I'm also increasing the pressure, which should tend to push the reaction to the right, towards fewer moles of gaseous products.
Is there a simple explanation for how Le Chatelier's Principle works in such a situation?
It is important to know the externally caused change the system is dealing with.
If the primary change is increasing the pressure, then the equilibrium shifts to the lower volume, in case of gases the lower molar amount.
If the primary change is increasing the temperature, than the equilibrium shifts away from the direction that evolves thermal energy.
If external causes are combined, like pressurizing the system AND externally initiated heat transfer, the resulting direction of equilibrium shift would depend on what effect is stronger. As the Le Chatelier's rule is qualitative, it cannot provide answer here, but it must be quantitatively calculated via the equilibrium constant at particular temperature.