Explain the importance of collisions between reacting particles as a criterion for determining reaction rates.
In order for a reaction to occur, the reactants must first collide with each other. This is necessary in order to induce bond breakage (between the reactants) and bond formation (to form the products).
Note that not all collision would result to the formation of the products. There are certain factors: for instance, energy of collision, or the proper orientation of the molecules of the reactant. An increase in the number of successful collisions lead to an increase in the reaction rate.This is so because if the probability of having sucessful collision (collisions of the proper orientation with sufficient energy) is higher, then the reaction would proceed faster since there's a higher rate of 'converting' reactant species to products.
Going back to the factors...
High energy is necessary because this would initiate the reaction by breaking the bonds of the reactants (which is endothermic) prior to forming the bonds that constitute the products. This is called the activation energy. If collisions have low energy, then it's possible that they won't overcome the activation energy, hence a collision will not result to product formation. To overcome this, one could increase the temperature of the system, hence increase average energy of each molecules (and hence, collisions) and a higher proportion of molecules/collisions will produce products. Hence, higher temperature increases the rate of the reaction. Adding catalysts lowers the activation energy, such that two colliding particles will need lower energy to form products. Both of these result to higher reaction rate.
The orientation of molecules is important as the atoms to be 'connected' must be in close proximity with each other. This is random, really. Increase in the number of total collisions (by increasing temperature) also increases the probability of having successful collisions, and hence higher reaction rates. The same goes for addition of catalysts.