Explain why, unlike gas-phase reactions, a reaction in solution is hardly affected at all by pressure.

When we delve into the world of gas-phase reactions, we open up a fascinating segment of chemistry where reactions take place between gaseous reactants. In these kinds of reactions, pressure plays a pivotal role in determining the rate at which the reaction occurs. This ties back to the concept that gases are compressible: when you increase the pressure on a gas, you effectively crowd the molecules closer together. This increase in molecular density is bound to lead to more collisions between the gas molecules.

Most gas-phase reactions are highly dependent on these collisions for the reactants to transform into products. An example can be seen in the synthesis of ammonia through the Haber process, where nitrogen and hydrogen gases react under high pressure to increase the chance of collisions and hence the reaction rate. The step by step solution provided illustrates this concept well, focusing on the direct correlation between pressure and the frequency of molecular collisions.

The Collision Theory paints a vivid picture of how chemical reactions occur at a microscopic level. It suggests that for a reaction to take place, reactant particles must collide with each other with sufficient energy and an appropriate orientation. This is known as an effective collision. One of the key points of this theory is that increasing the concentration of reactants will naturally lead to a higher frequency of collisions, thus accelerating the reaction rate.

When pressure is increased in a system involving gases, it's akin to packing more particles into a fixed space, which subsequently enhances the opportunity for effective collisions. However, the theory also stipulates that not all collisions result in a chemical reaction. Particles must overcome the activation energy barrier during their collision, which is why increasing the temperature can also play a significant role in chemical reaction rates, a concept interlinked with pressure in gas-phase reactions.

Diving into the realm of solutions in chemistry, it's important to distinguish how reactions here differ from those in the gas phase when subjected to pressure changes. A solution is a homogeneous mixture composed of a solute dissolved in a solvent. In liquid solutions, the particles are in constant, close contact due to the intermolecular forces at play, and varying the pressure doesn’t significantly alter this closeness or quantum of interaction.

Therefore, while gases respond dramatically to changes in pressure due to their compressible nature, liquids and solutions, which are virtually incompressible, showcase minimal reactivity changes when subjected to the same pressure alterations. This fundamental difference provides insight into why a reaction in solution is hardly affected at all by pressure, as the step by step solution correctly implies. It's the incompressibility of liquids that shields the reactions in solution from the influence of pressure, a nuance critical to understanding chemical processes in different states of matter.