Explain why, even though a collision may have energy in excess of the activation energy, reaction may not occur.

Collision theory is at the core of why chemical reactions occur. It explains the dynamics behind reactions and provides a scientific basis for understanding reaction rates. According to this theory, particles such as atoms, ions, or molecules must collide in order to react. But it's not just any collision that will do; these particles must collide with sufficient energy, known as the activation energy, which is necessary for breaking the original bonds so new ones can form.

The theory also emphasizes that only a percentage of collisions result in a reaction. This is because only those collisions with energy greater than or equal to the activation energy, as well as proper orientation, can lead to successful reactions. It's akin to the idea that only a key with the right shape and force can turn a lock. If either element is off, the lock won't turn, just as a chemical reaction won't proceed.

Rate of Reactions

Chemical kinetics is about understanding the speeds (or rates) at which reactions occur and the factors that affect these rates. Activation energy is one crucial factor. Think of it like a hill that reactants must climb over to transform into products. The height of this hill dictates how much energy is needed for a reaction to proceed.

Imagine pushing a ball up a hill—the higher the hill, the more effort you need. Similarly, in chemical kinetics, the higher the activation energy, the slower the reaction rate, as fewer particles will have enough energy to reach the top and 'roll down' to products. The relationship between the rate of reaction and activation energy is fundamental in predicting how fast a reaction will occur under different conditions.

The concept of reaction orientation adds another layer to our understanding of why reactions occur or fail to occur. As previously mentioned, even if particles collide with energy above the threshold of activation energy, they must be oriented in a particular way for a successful reaction to take place. It's a bit like building with blocks—if the blocks don't fit together in the correct way, you cannot build your structure.

The specific atomic or molecular arrangement required for a reaction to proceed is crucial, as it ensures that the right atoms bond together. Location, position, and timing all play a part—imagine a puzzle where each piece must align perfectly for completion. This is a vivid analogy for reaction orientation because, without correct alignment, no product is formed, even with enough energy.