Explain collision theory, transition state, activation energy, and factors affecting reaction rates.

You're on a roll — here are the explanations for your questions: Explain collision theory Collision theory states that for a chemical reaction to occur, reactant particles must collide with each other. These collisions must be effective*, meaning they must have sufficient energy (equal to or greater than the activation energy) and the correct orientation for bonds to break and form. Explain transition state The transition state* (or activated complex) is a high-energy, unstable intermediate structure that exists momentarily during a chemical reaction. In this state, old bonds are partially breaking, and new bonds are partially forming. It represents the peak of the energy barrier between reactants and products. Explain activation energy Activation energy* (E_a) is the minimum amount of energy that reactant molecules must possess to undergo a chemical reaction. It is the energy barrier that must be overcome for reactants to transform into products, allowing the formation of the transition state. List 5 factors that affect chemical reaction and explain 2 Five factors: 1. Concentration of reactants 2. Temperature 3. Surface area of reactants 4. Presence of a catalyst 5. Nature of reactants Explanation of two factors: Concentration of reactants: Increasing the concentration of reactants leads to a higher number of reactant particles per unit volume. This increases the frequency of collisions between reacting particles, thereby increasing the rate of reaction. Temperature: Raising the temperature increases the average kinetic energy of reactant particles. This results in more frequent collisions and, more importantly, a higher proportion of collisions having energy equal to or greater than the activation energy, leading to a faster reaction rate. Drop the next question 📸