1. The number of collisions per unit volume of the reaction mixture is called: ✔ (c) Collision frequency
2. If the reactants possess energy higher than activation energy, then reaction will be: ✔ (b) Fast
3. Which explains increase in rate with catalyst? ✔ (b) Catalyst provides alternative pathway which lowers activation energy.
4. Which statement is correct? ✔ (a) Collisions with energy ≥ activation energy lead to reaction.
5. How concentrations change as reaction proceeds? ✔ (b) [Reactants] ↓ , [products] ↑
All multiple choices are solved. Correct options are ticked.
📌 Short Answer Questions (B)
17.1 What is a successful collision?
A collision that results in reaction; occurs when molecules collide with sufficient energy (≥ activation energy) and proper orientation.
17.2 How does increase in temperature increase the rate of a reaction?
Increase in temperature raises the average kinetic energy of molecules, so a larger fraction possesses energy ≥ activation energy, and collision frequency also increases slightly → rate increases.
17.3 Define activation energy.
Activation energy (Eₐ) is the minimum amount of energy required for reactant molecules to undergo a successful collision and form products.
17.4 Why does burning of sulphur proceed slower in air than in pure oxygen?
Air contains only about 21% oxygen (lower concentration), while pure oxygen is 100%. Lower concentration of O₂ in air reduces collision frequency between S and O₂, hence slower reaction.
17.5 Why is a catalyst used in a reaction preferably taken in powdered form?
Powdered form greatly increases the surface area of the catalyst, providing more active sites for reactant molecules to adsorb, which increases the rate of reaction.
17.6 Why is the rate of a reaction often very fast at the beginning?
At the start, the concentrations of reactants are maximum, so collision frequency is highest. As reaction proceeds, reactants are consumed and rate decreases.
17.7 Magnesium does not react with air at room temperature but reacts fast at high temperature giving intense white light. Explain.
At room temperature, Mg atoms do not have sufficient energy to overcome activation energy. At high temperature, they gain kinetic energy, many collisions have energy ≥ Eₐ, so reaction becomes fast.
17.8 What happens to the reactants after they climb the energy hill during a reaction?
After crossing the activation energy barrier (top of energy hill), reactants form an activated complex which then breaks down into products. So they convert to products.
✏️ Constructed Response Questions (C)
17.1 In what different ways can you increase successful collisions between particles?
(i) Increase concentration of reactants. (ii) Increase temperature. (iii) Increase surface area (by grinding solid). (iv) Use catalyst to lower activation energy. (v) Stir/mix the reactants.
17.2 Give an example of a reaction which proceeds with gain in mass.
Rusting of iron (or any corrosion): Iron combines with oxygen from air to form iron oxide, mass increases because oxygen is added.
17.3 From where do molecules get energy to attain higher energy state?
Molecules absorb energy from surroundings: from heat (thermal energy), light (photons), or electrical energy. In many reactions, it’s from collision with other fast-moving molecules (heat).
17.4 How does the presence of V₂O₅ catalyst lower the activation energy of 2SO₂ + O₂ → 2SO₃?
V₂O₅ provides an alternative surface mechanism: SO₂ and O₂ adsorb on V₂O₅ surface, bonds weaken, intermediate complexes form, and the reaction proceeds via lower energy pathway. V₂O₅ is not consumed.
17.6 Explain the catalytic action of an enzyme.
Enzymes are biological catalysts with an active site that fits specific substrate (lock-and-key). Substrate binds, forming enzyme-substrate complex, which strains bonds and lowers activation energy, then product released. Highly efficient and specific.
17.7 If you desire to stop a reaction going on at 60°C, what action will you take?
Lower the temperature (e.g., put the reaction vessel in ice bath). Decreasing temperature reduces kinetic energy, so fewer molecules have energy ≥ Eₐ, thus reaction slows and can effectively stop. Alternatively remove a reactant or add inhibitor.
📚 Descriptive Questions (D) – overview
17.1 Effect of surface area: Increasing surface area (by grinding solid into powder) exposes more particles to collisions, hence increases rate. Example: wood chips burn faster than log.
17.2 Main points of collision theory: (1) Molecules must collide. (2) They must have energy ≥ activation energy. (3) Proper orientation. (4) Rate depends on collision frequency and fraction of effective collisions.
17.3 Factors affecting rate: (a) Concentration: higher conc → more collisions → faster. (b) Temperature: higher T → more molecules above Eₐ → faster. (c) Surface area, catalyst, nature of reactants.
17.4 Role of chemical kinetics in food industry: Helps understand spoilage rates; used to set optimal storage temperatures, predict shelf life, design preservatives, and control food quality.
17.5 Role of enzymes in body: Enzymes like amylase, protease, lipase speed up digestion, metabolism, DNA replication; they lower activation energy at body temperature; highly specific.
17.9 How catalyst lowers activation energy: Provides alternative reaction pathway (mechanism) with lower activation energy. It may form intermediates, does not get consumed.
17.10 Two features of catalytic action of enzyme: (i) Highly specific (lock & key). (ii) Works under mild temperature and pH; active site binds substrate and weakens bonds.
🗓️ Lesson planner (reaction kinetics)
Collision theoryActivation energyEffect of concTemperatureCatalystSurface areaEnzymesReaction rateMCQ drill
📝 Student guidelines & memory tips
★ Successful collision = energy ≥ Eₐ + right aim.
★ Catalyst ≠ extra energy; it’s a shortcut (lower hill).
★ Powder increases surface = more collisions.
★ At start, rate is fastest because [reactants] max.
★ Burning S slower in air due to less O₂ (concentration).
★ After energy hill → activated complex → products.