First off, heating an exothermic reaction causes an abundance of products (heat is a product). To reestablish equilibrium, heat must be consumed and the reaction must shift to the left. Unfortuantely though, this could get a bit confusing because the shift to the left may or may not result in an increase in pressure. If the pressure increases, there may also be a shift in the forward direction to reduce the moles of gas. If the container were not fixed (meaning the volume could expand), the pressure would not change and there would be no shift in the forward direction. I'm assuming for a locked container, an increase in temperature dominates and that the tendency of the reaction to shift to the right (due to increased pressure via temp increase) is negligible.
Choice D would be true if we're considering how the reaction would respond to the increase in temperature only. The question though is asking how we could increase the percent yield of ammonia NOT how to achieve equilibrium. Raising the temperature (adding more heat) forces the reaction to shift left to use up some of that heat, so more reactants are produced. Even if you somehow heated the reaction and then placed the container in a cool external environment to remove excess heat (via conduction or something), this does not increase the yield of ammonia. Instead heat would be lost faster and equilibrium would be reached sooner.
Choice C is wrong because increasing the volume will decrease the pressure. To reestablish equilibrium, the reaction will shift to the side with more moles of gas. For this question, that would be a shift to the left. This does not increase the yield of ammonia.
Choice B is essentially the issue I brought up earlier regarding the type of container. If the volume of the container could expand, there would be no change in pressure because the kinetic energy supplied by the increase in temperature is being used up to expand the volume of the container and thus cool it down until equilibrium is once again reached. In other words, this change won't shift the reaction in the forward direction (no increase of ammonia).
Choice A is correct. The only way we could produce ammonia by heating up this reaction is if we applied a significant amount of pressure as well. By increasing the overall pressure, there would be a shift in the forward direction to relieve the pressure increase by reducing the moles of gas. Choice D, C, and B result in no net increase in ammonia. By process of elimination, choice A is right.
I would imagine though that the yield still wouldn't be as high since there's both a forward and reverse reaction to consider. However, like Graffiti pointed out, it's probably due to the kinetics of the reaction. Some companies increase production by producing a low yield of ammonia at much quicker rate than a high yield of ammonia that would takes ages.
