Effective collisions lead to chemical reactions, while ineffective collisions do not.

Effective collisions occur when reactant molecules collide with sufficient energy and the correct orientation, allowing them to break existing bonds and form new ones. In essence, these successful collisions result in a chemical reaction. The minimum energy required for a reaction to take place is known as the activation energy, denoted as $E_a$. A higher activation energy corresponds to a lower probability of achieving an effective collision. Effective collisions are crucial for chemical reactions and are influenced by several factors, including temperature, concentration, and surface area.

On the other hand, ineffective collisions take place when reactant molecules collide without enough energy or at an incorrect orientation, resulting in no reaction. Such collisions do not contribute to the overall reaction rate and are therefore considered a waste of energy. Ineffective collisions can arise from conditions such as low temperature, low concentration, or insufficient surface area. By increasing the temperature, concentration, or surface area, we can enhance the frequency of effective collisions, thereby increasing the reaction rate.

In summary, effective collisions occur when reactants collide with adequate energy and the appropriate angle to form new bonds, which is contingent upon overcoming the activation energy barrier. Factors like temperature, concentration, and surface area can significantly increase the likelihood of these effective interactions. Conversely, ineffective collisions lack the necessary energy or orientation, resulting in no reaction. By boosting temperature, concentration, or surface area, we can convert more collisions into effective ones, thus accelerating the reaction rate.