This lesson explores why some chemical reactions happen instantly (like an explosion) while others take years (like rusting). Understanding the rate of reaction helps us control chemical processes in industry and everyday life.

1. Short Notes: Core Concepts

A. What is the Rate of a Reaction?

• The rate of a reaction is a measure of how quickly reactants are used up or how quickly products are formed.

• It's essentially the speed of the reaction.

B. Collision Theory

• For a reaction to occur, reactant particles (atoms, ions, or molecules) must collide with each other.

• The collisions must have enough energy to break old bonds and form new ones.

• Anything that increases the frequency or energy of these collisions will increase the reaction rate.

C. Factors Affecting the Rate of a Reaction

1. Surface Area of Reactants:

   • Rule: Increasing the surface area of a solid reactant increases the reaction rate.

   • Why? A powder has a much larger total surface area than a single lump of the same mass. This exposes more particles to the other reactant, leading to more frequent collisions.

   • Example: Powdered chalk reacts much faster with acid than a solid piece of chalk.

2. Concentration of Reactants:

   • Rule: Increasing the concentration of a reactant in a solution increases the reaction rate.

   • Why? Higher concentration means there are more reactant particles packed into the same volume. This increases the chances of particles colliding, so collisions become more frequent.

   • Example: A magnesium ribbon fizzes faster in concentrated HCl than in dilute HCl.

3. Temperature:

   • Rule: Increasing the temperature increases the reaction rate.

   • Why? (Two reasons!)

     1. Particles gain more kinetic energy and move faster, leading to more frequent collisions.

     2. More importantly, the collisions are more energetic, so a larger proportion of them have enough energy to result in a reaction.

   • Example: Food spoils faster at room temperature than in a refrigerator because the decomposition reactions are slower at lower temperatures.

4. Catalyst:

   • Definition: A substance that increases the rate of a reaction without being chemically changed or used up in the process.

   • How? A catalyst provides an alternative pathway for the reaction that requires less energy.

   • Example: Manganese dioxide (MnO₂) is a catalyst for the decomposition of hydrogen peroxide (H₂O₂). Enzymes in our body are biological catalysts.

2. Tips & Tricks for the Exam

• The Magic Words: When explaining why a rate changes, always use the word "collisions". The correct answer will almost always involve "more frequent collisions" or "more energetic collisions".

• Concentration vs. Pressure: For gases, increasing the pressure is the same as increasing the concentration. It forces the gas particles closer together, leading to more frequent collisions.

• Catalysts are Specific: A catalyst that works for one reaction might not work for another. It is not a reactant or a product, so it's often written above the arrow in a chemical equation.

3. Important Points & Common Exam Questions

• Comparing Rates: "An experiment is done by adding a lump of calcium carbonate to dilute acid. Suggest two ways to increase the rate of production of carbon dioxide gas."

  • Answer: 1. Use powdered calcium carbonate instead of a lump (increases surface area). 2. Gently heat the acid (increases temperature). / Use more concentrated acid.

• Explaining Observations: "Why does food last longer when kept in a refrigerator?"

  • Answer: The low temperature slows down the rate of chemical reactions (like decomposition by bacteria) that cause food to spoil.

• Role of a Catalyst: "In the decomposition of hydrogen peroxide (2H₂O₂ → 2H₂O + O₂), manganese dioxide is added. What is the role of manganese dioxide?"

  • Answer: It acts as a catalyst to speed up the rate of the reaction.