Q.In the presence of a catalyst, the heat evolved or absorbed during the reaction ___________.

The key idea is that a catalyst lowers the activation energy but does not affect the thermodynamic quantities of the reaction.

1. The heat evolved or absorbed during a reaction is the enthalpy change ($\Delta H$), which depends only on the initial and final states of the reactants and products.
2. A catalyst provides an alternative reaction pathway with a lower activation energy, but it does not change the initial or final states.
3. Therefore, $\Delta H$ remains the same whether a catalyst is present or not.

The key idea here is simple but often misunderstood: a catalyst affects the path of a reaction, not its destination. Let’s unpack why.

The Concept: What a Catalyst Actually Does

A catalyst provides an alternative reaction pathway with a lower activation energy. This means more reactant molecules have enough energy to cross the energy barrier per unit time, so the reaction rate increases.

But here’s the crucial point: the initial and final states of the reaction — the reactants and products — are exactly the same with or without the catalyst. The catalyst is not consumed and does not appear in the overall balanced equation.

The heat evolved or absorbed in a reaction is the enthalpy change, $\Delta H = H_{\text{products}} - H_{\text{reactants}}$. Since the reactants and products are identical in both the catalysed and uncatalysed reactions, $\Delta H$ must be the same.

Step-by-Step Reasoning

1. Identify what the question asks

   The “heat evolved or absorbed” is the enthalpy change of the reaction, $\Delta H$. This is a thermodynamic property, not a kinetic one.

2. Recall the role of a catalyst

   A catalyst speeds up the reaction by lowering the activation energy ($E_a$) for both the forward and reverse reactions equally. It does not alter the energies of the reactants or products themselves.

3. Visualise the energy profile

   Draw an energy diagram: the reactants are at some energy level, the products at another. The catalyst lowers the peak (the transition state) but leaves the two flat ends exactly where they were. The vertical drop (or rise) from reactants to products — that’s $\Delta H$ — stays the same.

4. Apply the principle

   Since $\Delta H$ depends only on the initial and final states, and those are unchanged, the heat evolved or absorbed remains unchanged.

Concept: Effect of Catalyst on Reaction Enthalpy

A catalyst provides an alternative reaction pathway with a lower activation energy, but it does not change the initial and final energy states of the reactants and products.

Method: First Law of Thermodynamics / Hess’s Law Approach

Step 1: Recall the definition of enthalpy change ($\Delta H$)

$\Delta H = H_{\text{products}} - H_{\text{reactants}}$

It depends only on the initial and final states, not on the path taken.

Step 2: Recognize the role of a catalyst

A catalyst speeds up the reaction by lowering the activation energy barrier — it participates in the reaction mechanism but is regenerated unchanged at the end.

Step 3: Apply the principle

Since the catalyst does not alter the identity or energy of reactants or products, the difference $H_{\text{products}} - H_{\text{reactants}}$ remains the same.

Step 4: Conclude

Therefore, the heat evolved or absorbed ($\Delta H$) remains unchanged.

Final Answer:

(iii) remains unchanged.

Here is the breakdown of the common mistakes students make on this concept, along with how to avoid them.

The Core Concept

A catalyst provides an alternative pathway (mechanism) for the reaction. This new pathway has a lower activation energy ($E_a$), which is why the reaction speeds up.

Crucially, a catalyst does not change the initial and final states of the reactants and products. Since the enthalpy change ($\Delta H$) depends only on the difference in energy between these initial and final states (Hess's Law), the catalyst cannot change $\Delta H$.

Therefore, the correct answer is (iii) remains unchanged.

Common Mistake #1: Confusing Rate with Energy Change

• The Mistake: Students think that because a catalyst makes a reaction happen faster, it must also change how much heat is released or absorbed. They assume "faster" means "more" or "less" heat.
• Why it happens: The word "catalyst" is often associated with "speed" in everyday language. Students fail to separate the kinetics (how fast) from the thermodynamics (how much energy).
• How to Avoid:
  • Draw the Energy Profile Diagram: Always sketch the reaction coordinate diagram. Draw the curve for the uncatalyzed reaction (high peak) and the catalyzed reaction (lower peak). Notice that the starting point (reactants) and ending point (products) are at the exact same energy levels on both curves.
  • Remember the Definition: $\Delta H = H_{products} - H_{reactants}$. The catalyst does not change $H_{products}$ or $H_{reactants}$.
  • Key Phrase: "A catalyst affects the path, not the destination."

Common Mistake #2: Thinking a Catalyst Absorbs or Releases Heat

• The Mistake: Students believe the catalyst itself participates in the reaction by absorbing heat (making it "less exothermic") or releasing heat (making it "more exothermic").
• Why it happens: Some students know that catalysts can be involved in the reaction mechanism (e.g., forming an intermediate) and assume this involvement changes the overall energy balance.
• How to Avoid:
  • The "Regeneration" Rule: A catalyst is chemically unchanged at the end of the reaction. If it were to absorb or release a net amount of heat, its own chemical structure or energy state would have to change permanently. Since it is regenerated, its net energy contribution to the system is zero.
  • Think of a "Middleman": A catalyst is like a middleman who helps two people trade goods. The middleman facilitates the trade but doesn't keep any of the goods or money. The net value of the trade is the same whether the middleman is there or not.

Common Mistake #3: Confusing Catalyst with an "Initiator" or "Igniter"

• The Mistake: Students think a catalyst "starts" a reaction that otherwise wouldn't happen, and therefore must supply the initial energy (heat) to get it going.
• Why it happens: This is a confusion between a catalyst and an initiator (like a spark plug or a match). An initiator provides the initial activation energy and is consumed in the process.
• How to Avoid:
  • Compare and Contrast:
    • Initiator: Provides energy, gets consumed, changes $\Delta H$ of the overall process (e.g., burning a match to start a fire adds the match's energy to the system).
    • Catalyst: Lowers the energy barrier, is not consumed, does not change $\Delta H$.
  • The "Free Pass" Analogy: A catalyst is like a "free pass" that lowers the entrance fee to a concert. It doesn't change the price of the ticket (the $\Delta H$), it just makes it easier to get in (lowers $E_a$).