Collision Theory, etc.

Kinetics- how factors affect chemical reactions. Lots of variations, lots of reasons.

For a reaction to happen, particles have to collide. This collision must have the energy to break the bonds, and must take place between the parts that are actually going to react. For many collisions, a high concentration of particles in a small space is ideal. They also need to move quick, to collide with enough force.

According to collision thoery;

Most collisions between molecules or other particles do not lead to reaction. They either do not have enough energy, or are in the wrong orientation.

This makes a fair amount of sense. If the particles did not have enough energy, the eA could not be reached and the reaction could not take place, as bonds could not be broken to form products. (eA being 'Activation Energy'). Also, orientation wise, as stated above, the collsion must take place between the parts of the molecule that are actually going to react, therefore, if must be orientated the right way, or, in english, pointing itself in the right direction. Kind of makes me think of car accidents. If the collision was represented as someone getting hit by a car, and the car hit them on the shoulder 'bond' this wouldn't be likely to break their leg... 'bond'.

Did that make any sense? xP

Many factors affect the rate of reaction.

• Temperature - it increases the kinetic energy of the particles. This increase the speed with which they move, which increases the amount of force (and in turn, energy) the collide with, and also the amount of collisons, as they are moving quicker.
• Increasing the concentration of a solution - it means there are more particles, and that means that collisions are more likely to happen. As a reaction goes on, the reactants are used up, which means (as the concentration drops) the rate slows down as the reaction happens.
  
• Increasing surface area of solid reactants - this give more of a surface of which to react, so the reaction will take place quicker. At first, I found this a tad confusing. But, thinking of it like baking cakes kinda helped. If you're baing cakes, you'll get more made in a bigger oven then a small one, right? So, if the cakes are reactants, and the oven is the solid reactant, if it's bigger, you can make more cakes. And everyone likes cakes. ^_____^
  

I mean products... xP

• Anyway, increasing the pressure - it forces more molecules into a give volume, so, like increasing the concentration, there is a natural increase in collsions, which means more chance of sucessful collsions, and more reactions taking place quicker.
  
• Catalysts - They lower the activation energy of the reaction by providing an alternate pathway. This means that, with a lower activation energy more particles with be able to colide with /that/ amount of energy to break the bonds, so the reaction will happen faster.
  

So anyway, the minimum energy required to kick off a reaction is a Activation Energy. In what I've jsut read up, there's another of those funny little comparison things that science textbooks seem to love (and me, but meh. :P) But yea, this ones pretty good. It goes on about ball rolling down a hill.
If there's a large activation energy, it takes a lot to get the ball to the top of the hill, to start rolling down, but if its a lower activation energy it takes less to get it to the top of the hill to get it rolling down. The rolling down part being the reaction.

In endothermic reactions, the energy of the products is higher then the energy of the reactants, as endothermic reactions take /in/ heat, and therefore energy, so there is going to be more.