Alright class, let's embark on an exciting journey into the world of Chemical Kinetics, specifically focusing today on a very special type of substance that can magically speed up reactions: the catalyst! Think of them as the unsung heroes of many chemical processes, from the ones happening in your body to industrial production plants.

### What is a Catalyst? Your Reaction's Best Friend!

Have you ever been stuck in a traffic jam, wishing there was a secret shortcut to get to your destination faster? Or perhaps you've seen a difficult task become incredibly easy with the right tool or guide? That's precisely what a catalyst does for a chemical reaction!

In simple terms, a catalyst is a substance that increases the rate of a chemical reaction without itself being consumed in the reaction. It's like that helpful friend who shows you a quicker, easier path to climb a mountain. You still climb the mountain, but it takes less effort and less time because of the path they showed you.

Let's break down this definition with some key points:

1. Speeds up reaction: This is their primary job. They make reactions go faster, sometimes thousands or even millions of times faster!
2. Not consumed: This is super important! The catalyst participates in the reaction, helps it along, but at the end, it comes out exactly as it was at the beginning, ready to help another set of reactants. It's like the guide who takes you up the mountain; they are still there, perfectly fine, ready to guide the next group.
3. Required in small amounts: Because they aren't consumed, a tiny amount of catalyst can process a huge amount of reactants over time.

### How Does a Catalyst Work? The "Shortcut" Mechanism!

Now, the big question: how do catalysts perform this magic? To understand this, we need to recall our discussion on Activation Energy (Ea). Remember, activation energy is the minimum amount of energy that reacting particles must possess for a reaction to occur. It's like the "energy barrier" or the "hump" reactants need to get over to transform into products.

Imagine you want to push a heavy box over a hill. The height of the hill is your activation energy. You need to put in a lot of effort (energy) to get it to the top, and once it's over, it rolls down to the other side (products).

What a catalyst does is provide an alternative reaction pathway or a "new mechanism" for the reaction to proceed. This new pathway has a lower activation energy than the original, uncatalyzed pathway.

Let's go back to our hill analogy:
* Uncatalyzed reaction: You push the box over the original, tall hill. Lots of effort needed.
* Catalyzed reaction: The catalyst *doesn't make the box lighter* or *give you more strength*. Instead, it magically creates a tunnel through the hill or finds a much lower, easier pass. Now, you need much less effort to get the box from one side to the other.

So, in chemical terms:
1. Reactants interact with the catalyst, forming temporary intermediate complexes.
2. These intermediates then break down to form the products and regenerate the catalyst.
3. Crucially, all the steps in this new pathway, involving the catalyst, require less activation energy than the single, high-energy step of the uncatalyzed reaction.

Because more reactant molecules will have sufficient energy to overcome this lower energy barrier at any given temperature, the reaction proceeds much faster! It's as simple as that – lower barrier, faster reaction.

Key takeaway for JEE: A catalyst *does not change the energy of the reactants or products*. It only changes the *path* between them, specifically lowering the activation energy of the forward and reverse reactions by the same amount.

### Important Characteristics of Catalysts

Beyond just lowering activation energy and not being consumed, catalysts have several other fascinating properties:

1. Specificity: Many catalysts are highly specific. This means a particular catalyst might only work for one specific reaction or a specific type of reaction, and not for others. Think of a key fitting only one lock. Enzymes (biological catalysts) are prime examples of this!
* Example: The enzyme *urease* specifically breaks down urea into ammonia and carbon dioxide, but it won't catalyze the breakdown of other compounds like glucose.
2. Does not initiate reaction: Generally, a catalyst cannot start a reaction that is thermodynamically impossible (i.e., a reaction that wouldn't happen on its own even if given infinite time). It can only speed up a reaction that *is already going to happen*, just very slowly.
3. Does not change Equilibrium Position: This is a crucial point for JEE! A catalyst speeds up both the forward and the reverse reactions equally. This means it helps the system reach equilibrium faster, but it does not change the final equilibrium concentrations of reactants and products. The position of equilibrium remains unaffected.
* Think of it this way: if you have two teams pulling a rope (forward and reverse reactions), and a catalyst makes both teams pull faster, they will still end up in the same balanced position, just quicker.
4. Promoters and Poisons:
* Promoters are substances that enhance the activity of a catalyst. They are not catalysts themselves but help the catalyst do its job better.
* Example: In the Haber process for ammonia synthesis (N₂ + 3H₂ → 2NH₃), iron is the catalyst, and molybdenum (Mo) acts as a promoter, enhancing the activity of iron.
* Poisons are substances that decrease or destroy the activity of a catalyst. They can block active sites on the catalyst or alter its surface.
* Example: Carbon monoxide (CO) is a notorious poison for many metal catalysts, including those used in catalytic converters in cars.

### Types of Catalysis (A Quick Glimpse for Fundamentals)

We broadly classify catalysis based on the physical state (phase) of the catalyst compared to the reactants:

1. Homogeneous Catalysis: This occurs when the catalyst and the reactants are in the same phase (e.g., all are gases, or all are liquids).
* Example: The decomposition of hydrogen peroxide (H₂O₂) can be catalyzed by iodide ions (I⁻) in an aqueous solution. Here, both H₂O₂, I⁻, and the resulting H₂O and O₂ are in the aqueous phase.
* 2 H₂O₂(aq) + I⁻(aq) → 2 H₂O(l) + O₂(g) + I⁻(aq) (I⁻ is regenerated)
2. Heterogeneous Catalysis: This occurs when the catalyst is in a different phase from the reactants. Usually, the catalyst is a solid, and the reactants are gases or liquids.
* Example: The synthesis of ammonia in the Haber process uses gaseous nitrogen and hydrogen reacting on a solid iron catalyst. The gases adsorb onto the surface of the solid iron, react, and then desorb as ammonia.
* N₂(g) + 3H₂(g) --(Fe catalyst)--> 2NH₃(g)

### Real-World Examples to Cement Your Understanding

Catalysts are everywhere!

* Your Car's Catalytic Converter: This is a fantastic example of heterogeneous catalysis. It uses precious metals like platinum, palladium, and rhodium (solid catalysts) to convert harmful gases (like carbon monoxide and nitrogen oxides from car exhaust) into less harmful ones (carbon dioxide, nitrogen, and water vapor).
* Enzymes in Your Body: These are biological catalysts, mostly proteins, that facilitate nearly all biochemical reactions essential for life. For example, the enzyme *amylase* in your saliva helps break down complex carbohydrates (starch) into simpler sugars, even at body temperature. Without it, your digestion would be incredibly slow!
* Industrial Production: Many large-scale industrial processes rely heavily on catalysts to make production efficient and cost-effective. We already discussed Haber's process for ammonia and the Contact Process for sulfuric acid (using V₂O₅ as catalyst).

### Connecting to JEE: Why is This Important?

Understanding catalysts is absolutely crucial for JEE, not just conceptually but also for solving numerical problems:

* You'll encounter questions asking about the effect of catalysts on activation energy, reaction rate, equilibrium constant, and enthalpy changes. Always remember: Catalyst lowers Ea, increases rate, *no change* to Keq or ΔH.
* The concept of providing an alternative mechanism is fundamental to understanding how catalysts function at a molecular level.
* Knowledge of different types of catalysis and specific examples will help you with multiple-choice questions.

So, in essence, catalysts are like the master chefs of the chemical world. They don't add ingredients or change the final dish, but they ensure the cooking happens much faster and more efficiently by providing a better recipe (alternative pathway) and the right tools (lowering activation energy)!

Keep these fundamentals in mind, and you'll have a strong base for our deeper dive into catalyst mechanisms!

Welcome, future engineers and scientists! Today, we're diving deep into one of the most fascinating and industrially crucial concepts in Chemical Kinetics: the Concept of Catalysis and its Basic Mechanism. This isn't just a theoretical idea; it's the backbone of countless industrial processes, from producing fertilizers to refining petroleum, and even fundamental to life itself through enzymes.

### 1. What Exactly is a Catalyst? Unraveling the Mystery

Let's start from the very beginning. Imagine you have a chemical reaction that is thermodynamically favorable – meaning it *wants* to happen – but it's excruciatingly slow. It might take hours, days, or even centuries to produce a significant amount of product. Why? Because there's often an energy barrier that needs to be overcome.

This is where a catalyst comes in. A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed in the overall reaction. It participates in the reaction, changes its pathway, but is regenerated at the end, ready to facilitate more reactions.

Think of it like this: You want to travel from one side of a mountain to the other. You *could* climb over the mountain (the uncatalyzed path), which takes a lot of energy and time. Or, you could find (or build) a tunnel through the mountain. The tunnel (the catalyzed path) allows you to get to the other side much faster and with less effort. The tunnel builder (the catalyst) isn't "used up" in the process; it remains, allowing countless others to use the path.

Key Definition:
A catalyst is a substance that alters the rate of a chemical reaction without undergoing any permanent chemical change itself. It typically speeds up reactions by providing an alternative reaction pathway with a lower activation energy.

### 2. The Energy Story: How Catalysts Affect Activation Energy (Ea)

To understand *how* a catalyst works, we need to revisit our understanding of activation energy (Ea) and reaction coordinate diagrams.

Recall that for a reaction to occur, reactant molecules must collide with sufficient energy (greater than or equal to Ea) and in the correct orientation. The activation energy is the minimum energy required for reactants to be converted into products via a transition state.

Consider an uncatalyzed reaction:
$ ext{A} + ext{B}
ightarrow ext{P}$

This reaction proceeds through a high-energy transition state, requiring a significant Ea.

Now, with a catalyst (Cat) involved, the reaction proceeds through a different, multi-step mechanism:
1. $ ext{A} + ext{Cat}
ightarrow ext{A-Cat}$ (Intermediate formation)
2. $ ext{A-Cat} + ext{B}
ightarrow ext{P} + ext{Cat}$ (Product formation and catalyst regeneration)

Each of these new steps has its own activation energy. The crucial point is that the highest activation energy among these new steps is lower than the activation energy of the uncatalyzed reaction.

Feature	Uncatalyzed Reaction	Catalyzed Reaction
Reaction Pathway	Direct, single step (often conceptual)	Alternative, multi-step pathway
Activation Energy (Ea)	High	Lower (for the rate-determining step)
Reaction Rate	Slow	Faster
ΔH (Enthalpy Change)	Unaffected	Unaffected
ΔG (Gibbs Free Energy Change)	Unaffected	Unaffected
Equilibrium Position	Unaffected	Unaffected (speeds up both forward and reverse equally)

Visualizing with a Reaction Coordinate Diagram:
Imagine a hill. Without a catalyst, you have to climb over a very tall hill to get to the product side. With a catalyst, you might have to climb several smaller hills, but the highest point you reach on this new path is significantly lower than the peak of the original tall hill.

This lowering of activation energy is the fundamental reason why catalysts speed up reactions. More molecules at a given temperature will possess the necessary activation energy to react, leading to a higher frequency of successful collisions and thus a faster reaction rate.

JEE Focus: Understanding that a catalyst *does not* change $Delta H$ or $Delta G$ is absolutely crucial for JEE. It only affects the *kinetics* (rate) by altering the *pathway* and thus the *activation energy*. It speeds up the attainment of equilibrium but does not shift the equilibrium position itself. Why? Because if it lowered the activation energy for the forward reaction, it must also lower it by the *same amount* for the reverse reaction.

### 3. Key Characteristics of Catalysts

Beyond just lowering Ea, catalysts possess several important features:

1. Small Amount Required: A very small amount of catalyst can catalyze a large amount of reactants. Since it's regenerated, it acts as a continuous cycle.
2. Specificity: Many catalysts are highly specific in their action. For example, an enzyme that catalyzes the hydrolysis of an ester will not catalyze the synthesis of a protein. This is often referred to as "lock and key" specificity, particularly for biological enzymes.
3. No Initiation: A catalyst does not initiate a reaction that is thermodynamically impossible (i.e., if $Delta G$ is positive, it won't magically make it negative). It only speeds up reactions that are already thermodynamically feasible.
4. No Change in Equilibrium: As discussed, a catalyst speeds up both the forward and reverse reactions equally. Therefore, it helps the system reach equilibrium faster but does not change the position or composition of the equilibrium mixture.
5. Regenerated at the End: The catalyst remains chemically unchanged in quantity and composition at the end of the reaction. While it participates in intermediate steps, it is always reformed.
6. Can be Promoted or Poisoned:
* Promoters: Substances that enhance the activity of a catalyst. For example, molybdenum (Mo) acts as a promoter in the Haber process for ammonia synthesis, enhancing the activity of the iron catalyst.
* Poisons: Substances that decrease or destroy the activity of a catalyst. For instance, carbon monoxide (CO) is a common poison for many metal catalysts, as it strongly adsorbs to the active sites, preventing reactants from binding.

### 4. Basic Mechanism Idea: The Intermediate Complex Theory

The most common basic mechanism to explain catalysis is the Intermediate Complex Theory.

Consider a reaction: $ ext{A} + ext{B}
ightarrow ext{Products}$

Uncatalyzed reaction:
$ ext{A} + ext{B} xrightarrow{k} ext{Transition State}
ightarrow ext{Products}$ (High $ ext{E}_{ ext{a}}$)

Catalyzed reaction (two steps):
1. Formation of an Intermediate Complex: The catalyst (Cat) first reacts with one of the reactants (say, A) to form an unstable intermediate complex (A-Cat).
$ ext{A} + ext{Cat}
ightarrow ext{A-Cat}$ ($ ext{E}_{ ext{a1}}$)
2. Reaction of Intermediate and Regeneration of Catalyst: This intermediate then reacts with the other reactant (B) to form the products, and simultaneously regenerates the catalyst.
$ ext{A-Cat} + ext{B}
ightarrow ext{Products} + ext{Cat}$ ($ ext{E}_{ ext{a2}}$)

The key here is that both $ ext{E}_{ ext{a1}}$ and $ ext{E}_{ ext{a2}}$ are significantly lower than the activation energy of the uncatalyzed reaction.

Example: Decomposition of Hydrogen Peroxide ($ ext{H}_2 ext{O}_2$) catalyzed by Iodide Ions ($ ext{I}^-$)

The uncatalyzed decomposition of hydrogen peroxide is slow:
$2 ext{H}_2 ext{O}_2(aq)
ightarrow 2 ext{H}_2 ext{O}(l) + ext{O}_2(g)$ (Slow, High $ ext{E}_{ ext{a}}$)

In the presence of iodide ions, the reaction speeds up significantly, and the $ ext{I}^-$ ions are regenerated:
Step 1: $ ext{H}_2 ext{O}_2(aq) + ext{I}^-(aq)
ightarrow ext{H}_2 ext{O}(l) + ext{IO}^-(aq)$ (Faster, lower $ ext{E}_{ ext{a1}}$)
*(Here, $ ext{IO}^-$ is the intermediate complex)*

Step 2: $ ext{H}_2 ext{O}_2(aq) + ext{IO}^-(aq)
ightarrow ext{H}_2 ext{O}(l) + ext{O}_2(g) + ext{I}^-(aq)$ (Faster, lower $ ext{E}_{ ext{a2}}$, $ ext{I}^-$ regenerated)
*(The $ ext{I}^-$ catalyst is regenerated)*

Overall reaction: $2 ext{H}_2 ext{O}_2(aq) xrightarrow{ ext{I}^-(aq)} 2 ext{H}_2 ext{O}(l) + ext{O}_2(g)$

Notice how the catalyst, $ ext{I}^-$, is consumed in the first step and then regenerated in the second step, thus not being consumed in the overall reaction. The new pathway involving the $ ext{IO}^-$ intermediate has a much lower activation energy barrier.

### 5. Types of Catalysis and Their Mechanisms

Catalysis is broadly classified based on the physical state (phase) of the reactants and the catalyst:

#### a) Homogeneous Catalysis

In homogeneous catalysis, the catalyst and the reactants are in the same physical phase (e.g., all are gases, or all are liquids/solutions).

Mechanism Idea: The intermediate complex theory, as discussed above, is a perfect fit for homogeneous catalysis. The catalyst molecule directly interacts with reactant molecules in the same phase to form an intermediate, which then breaks down to give products and regenerate the catalyst.

Examples:
1. Acid-base catalysis (e.g., Hydrolysis of Esters):
$ ext{CH}_3 ext{COOC}_2 ext{H}_5(aq) + ext{H}_2 ext{O}(l) xrightarrow{ ext{H}^+(aq)} ext{CH}_3 ext{COOH}(aq) + ext{C}_2 ext{H}_5 ext{OH}(aq)$
Here, $ ext{H}^+$ ions (from an acid like $ ext{HCl}$ or $ ext{H}_2 ext{SO}_4$) are in the same liquid phase as the ester and water. The $ ext{H}^+$ protonates the carbonyl oxygen of the ester, making the carbon more electrophilic and susceptible to nucleophilic attack by water, lowering the Ea.

2. Oxidation of Carbon Monoxide by Nitric Oxide:
$2 ext{CO}(g) + ext{O}_2(g) xrightarrow{ ext{NO}(g)} 2 ext{CO}_2(g)$ (in the stratosphere, where $ ext{NO}$ acts as a catalyst)
Mechanism:
* $2 ext{NO}(g) + ext{O}_2(g)
ightarrow 2 ext{NO}_2(g)$ (NO consumed)
* $2 ext{NO}_2(g) + 2 ext{CO}(g)
ightarrow 2 ext{CO}_2(g) + 2 ext{NO}(g)$ ($ ext{NO}$ regenerated)
Here, $ ext{NO}_2$ is the intermediate, and $ ext{NO}$ is the regenerated catalyst.

#### b) Heterogeneous Catalysis

In heterogeneous catalysis, the catalyst is in a different physical phase from the reactants. Typically, the catalyst is a solid, and the reactants are gases or liquids. This type of catalysis is extremely important in industrial processes.

Mechanism Idea: Adsorption Theory of Catalysis
This theory explains how solid catalysts function. It involves a series of steps on the catalyst surface:

1. Adsorption of Reactant Molecules: Reactant molecules from the gas or liquid phase diffuse to the surface of the solid catalyst and get weakly adsorbed onto active sites. Active sites are specific locations on the catalyst surface with unsatisfied valencies, typically corners, edges, or structural defects.
2. Activation of Adsorbed Molecules: The adsorption process weakens the bonds within the reactant molecules, making them more reactive. This effectively lowers the activation energy for their conversion into products.
3. Reaction on the Catalyst Surface: The adsorbed reactant molecules interact with each other on the catalyst surface, forming new bonds and eventually products.
4. Desorption of Product Molecules: Once formed, the product molecules desorb (detach) from the catalyst surface and diffuse away, leaving the active sites free for new reactant molecules.

Example: Haber-Bosch Process for Ammonia Synthesis
$ ext{N}_2(g) + 3 ext{H}_2(g) xrightarrow{ ext{Fe}(s)} 2 ext{NH}_3(g)$

Here, solid iron acts as the catalyst, and nitrogen and hydrogen are gases.
* $ ext{N}_2$ and $ ext{H}_2$ molecules adsorb onto the surface of the iron catalyst.
* The strong $ ext{N} equiv ext{N}$ and $ ext{H}- ext{H}$ bonds are weakened (or even broken) upon adsorption.
* Activated nitrogen and hydrogen atoms/fragments react on the surface.
* $ ext{NH}_3$ molecules form and then desorb from the surface.

Other Important Examples for JEE:
* Hydrogenation of Vegetable Oils: $ ext{Vegetable Oil}(l) + ext{H}_2(g) xrightarrow{ ext{Ni}(s), ext{Pt}(s) ext{ or } ext{Pd}(s)} ext{Vegetable Ghee}(s)$
* Ostwald Process (Nitric Acid production): $ ext{NH}_3(g) + ext{O}_2(g) xrightarrow{ ext{Pt/Rh gauze}(s)} ext{NO}(g)$
* Contact Process (Sulfuric Acid production): $2 ext{SO}_2(g) + ext{O}_2(g) xrightarrow{ ext{V}_2 ext{O}_5(s)} 2 ext{SO}_3(g)$

JEE Advanced Focus: For heterogeneous catalysis, understanding the concept of active sites is critical. The efficiency of a heterogeneous catalyst often depends on the number and nature of these sites. Surface area of the catalyst is also a major factor – a larger surface area generally provides more active sites. The interaction strength (adsorption enthalpy) of reactants with the catalyst surface is a delicate balance; it shouldn't be too weak (no adsorption) or too strong (products can't desorb).

### 6. Autocatalysis

An interesting special case is autocatalysis, where one of the products of the reaction itself acts as a catalyst for the same reaction. This leads to a unique reaction profile where the rate of reaction initially increases as more product (and thus more catalyst) is formed, then eventually slows down as reactants are consumed.

Example:
The oxidation of oxalic acid by acidified potassium permanganate:
$2 ext{MnO}_4^-(aq) + 5 ext{C}_2 ext{O}_4^{2-}(aq) + 16 ext{H}^+(aq)
ightarrow 2 ext{Mn}^{2+}(aq) + 10 ext{CO}_2(g) + 8 ext{H}_2 ext{O}(l)$

Initially, the reaction is slow. However, as $ ext{Mn}^{2+}$ ions are formed, they act as an autocatalyst, significantly speeding up the reaction. The purple color of $ ext{MnO}_4^-$ disappears much faster once some $ ext{Mn}^{2+}$ has accumulated.

### 7. Enzymes: Nature's Catalysts

Finally, a brief mention of enzymes. These are biological catalysts, almost exclusively proteins, that catalyze biochemical reactions in living organisms. They are highly efficient and incredibly specific, often catalyzing only one specific reaction or a very small group of reactions. Their mechanism generally follows the intermediate complex theory, where the reactant (substrate) binds to the enzyme's active site to form an enzyme-substrate complex, which then transforms into product(s) and releases the enzyme.

CBSE vs. JEE Focus:
* CBSE: Focuses on the definition, characteristics, types (homogeneous/heterogeneous) with basic examples like Haber, Contact processes, and a general understanding of how catalysts lower Ea.
* JEE: Requires a much deeper conceptual understanding. This includes detailed mechanisms (intermediate formation, adsorption theory, multi-step pathways), the precise effect on reaction coordinate diagrams (multiple transition states, relative Ea values), the non-effect on thermodynamics ($Delta H$, $Delta G$, equilibrium constant K), the role of promoters and poisons, and sometimes specific industrial process details and their underlying catalytic principles. Be prepared for questions that involve drawing reaction coordinate diagrams for catalyzed vs. uncatalyzed reactions, or analyzing rate laws in the presence of a catalyst.

Catalysis is a vast and fascinating field, demonstrating how a small intervention can drastically alter the course of chemical events. Mastering this concept is key to understanding chemical kinetics at an advanced level and appreciating its profound impact on both nature and technology.

Here are some effective mnemonics and short-cuts to help you quickly recall the concept of catalysts and their basic mechanism for your JEE Main and board exams.

Understanding catalysts is crucial as they play a significant role in industrial processes and biological systems. These mnemonics will help you remember their key properties and how they function.

1. Mnemonic for Core Effects and Properties: The "CAT-ALYST" Mantra

This mnemonic covers the most important aspects of a catalyst's function and general properties.

• Cuts Activation energy (Ea).


• Trims reaction Time (speeds up the reaction, helping it reach equilibrium faster).


• Alternate pathway provided.


• Leaves $Delta G$ (Gibbs Free Energy) and $K_{eq}$ (Equilibrium Constant) Lone (unchanged).


• You (the catalyst) remain chemically Yourself at the end (chemically unchanged).


• Small Supply is sufficient (a small amount can catalyze a large reaction).


• Temperature and pressure dependent (catalytic activity is often optimal within specific ranges).

JEE Tip: Remember that a catalyst does NOT initiate a reaction, nor does it affect the thermodynamics (like $Delta G$) or the position of equilibrium. It only affects the kinetics (rate) of the reaction.

2. Mnemonic for Mechanism of Catalyst Action: The "I-P.A.L." Approach

This short mnemonic helps you recall the fundamental way a catalyst works at a molecular level.

• Intermediate formation: The catalyst typically forms a temporary intermediate complex with one or more reactants.


• Provides Alternate Low-energy pathway: This intermediate then reacts further to form products, regenerating the catalyst in the process. This new pathway has a lower activation energy barrier compared to the uncatalyzed reaction.

Think of a catalyst as your "PAL" (friend) who helps you find an "I-nnovative Pathway At Lower" energy.

3. Mnemonic for Key Characteristics of a Catalyst: "S.C.A.L.E."

This covers additional important attributes of catalysts.

• Specificity: Catalysts are often highly specific in their action, meaning a particular catalyst works for a specific reaction or a specific type of reaction. (e.g., enzymes are highly specific biological catalysts).


• Chemically unchanged: At the end of the reaction, the catalyst is recovered in its original chemical form and amount.


• Amount: Only a small quantity of catalyst is typically required to significantly alter the reaction rate.


• Lowers activation energy: This is the primary reason for the increased reaction rate.


• Equilibrium: A catalyst does not shift the position of equilibrium; it only helps the system attain equilibrium faster by speeding up both the forward and reverse reactions equally.

Short-cut/Quick Tip: The "Energy Hill" Analogy

Visualize the activation energy as an "energy hill" that reactants must climb to become products. A catalyst doesn't change the starting point (reactants' energy) or the ending point (products' energy), nor does it change the overall height difference between them ($Delta H$ or $Delta G$). Instead, it simply digs a tunnel or builds a ramp through or around the hill, making the climb easier and faster. The hill itself (the overall reaction) remains in the same location, but the path to get over it is now much quicker.

By using these mnemonics and the analogy, you can quickly recall the essential concepts of catalysts, which will be beneficial for both objective and subjective questions.

Here are some quick tips to master the concept of catalysts and their basic mechanism, crucial for both JEE Main and board exams.

Quick Tips: Catalyst Concept & Mechanism

• Catalyst Definition: A substance that increases the rate of a chemical reaction without being consumed in the overall process. It participates in the reaction but is regenerated in its original form.


• Mechanism - Lowering Activation Energy (Ea):
  
  
  
  • The primary function of a catalyst is to provide an alternative reaction pathway with a lower activation energy (Ea).
  
  
  • A lower Ea means a larger fraction of reactant molecules possess the necessary activation energy, leading to more effective collisions and a faster reaction rate.
  
  
  • JEE Tip: Always visualize the energy profile diagram. The 'hump' (Ea) is lowered, but the initial and final energy levels (reactants and products) remain unchanged.
  
  
  
  


• What a Catalyst DOES NOT Change:
  
  
  
  • Enthalpy Change (ΔH): The overall energy difference between reactants and products.
  
  
  • Gibbs Free Energy Change (ΔG): The spontaneity of the reaction.
  
  
  • Equilibrium Constant (K_eq): It speeds up both forward and reverse reactions proportionally, so equilibrium is reached faster, but its position is unaltered.
  
  
  • Nature of Reactants/Products: It does not change what the reactants are or what products are formed.
  
  
  • Initiation: A catalyst cannot initiate a non-spontaneous reaction.
  
  
  
  


• What a Catalyst DOES Change:
  
  
  
  • Rate of Reaction: Significantly increases.
  
  
  • Rate Constant (k): A catalyst increases the rate constant 'k' as per the Arrhenius equation (k = Ae^-Ea/RT) due to the reduced Ea.
  
  
  • Time to Reach Equilibrium: Reduces the time required to achieve equilibrium.
  
  
  • Reaction Mechanism: Provides a new, alternative reaction mechanism, often involving the formation of an intermediate.
  
  
  
  


• Basic Mechanism Idea (Intermediate Formation):
  
  
  
  • In many catalyzed reactions, the catalyst reacts with one or more reactants to form an unstable intermediate.
  
  
  • This intermediate then decomposes to form products and regenerates the catalyst.
  
  
  • Example:
    
    A + C → AC (intermediate, lower Ea)
    
    AC + B → AB + C (catalyst regenerated)
    
  
  
  • Boards Focus: Understand this intermediate formation idea.
  
  
  
  


• Characteristics of Catalysts:
  
  
  
  • Specificity: Many catalysts are highly specific, catalyzing only certain reactions (especially enzymes).
  
  
  • Small Amount: Only a small amount of catalyst is usually sufficient to catalyze a large quantity of reactants.
  
  
  • Promoters: Substances that enhance the activity of a catalyst.
  
  
  • Poisons: Substances that decrease or inhibit the activity of a catalyst.
  
  
  
  

Remember: For JEE, focus on the quantitative impact (on Ea, k, reaction rate) and the non-impact on thermodynamic quantities. For boards, a clear understanding of the definition, properties, and the basic mechanism is key.

Welcome to the 'Intuitive Understanding' section for Catalysts!

In chemical kinetics, understanding the role of a catalyst is crucial for both JEE Main and board exams. It often appears in questions related to reaction mechanisms and energy profiles. Let's grasp its core concept intuitively.

What is a Catalyst? The "Helper" Molecule

Imagine you have a task (a chemical reaction) that's difficult or slow to complete. A catalyst is like a helpful tool or a guide that makes the task much easier and faster, without actually being used up in the process. It doesn't become part of the final product, and you can theoretically reuse it for many more tasks.

• A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed in the overall reaction.


• It can be recovered chemically unchanged at the end of the reaction.

How Does It Work? Lowering the Energy Barrier

Think of a chemical reaction as climbing over a mountain to get from one valley (reactants) to another (products). The peak of the mountain represents the activation energy (E_a) – the minimum energy required for reactants to transform into products.

Intuitive Analogy: The Mountain Pass

Suppose you want to travel from city A to city B. There's a tall mountain range between them (this is your activation energy). You could climb over the mountain, which is slow and requires a lot of effort (high E_a). A catalyst, however, is like discovering a tunnel or a lower, easier mountain pass through the range. You still go from city A to city B, but the journey is much faster and requires less effort (lower E_a). The tunnel itself isn't consumed; you can use it repeatedly.

• Catalysts provide an alternative reaction pathway that has a lower activation energy than the uncatalyzed reaction.


• By lowering E_a, a greater fraction of reactant molecules possess the necessary activation energy at a given temperature, leading to more effective collisions and thus a faster reaction rate.

Basic Mechanism Idea: Providing a New Route

The "alternative pathway" means that the catalyst interacts with the reactants in a specific way, forming an intermediate (often called the catalyst-reactant complex). This intermediate then breaks down to form the products and regenerates the original catalyst.

Consider a simple reaction: R → P

Without catalyst: R → [Activated Complex] → P (High E_a)

With catalyst (C):

1. R + C → RC (Intermediate complex formation, lower E_a1)


2. RC → P + C (Product formation and catalyst regeneration, lower E_a2)

The overall energy barrier for the two steps (E_a1 + E_a2) is significantly lower than the single step in the uncatalyzed reaction. The catalyst (C) is present at the start and regenerated at the end, unchanged.

Key Characteristics and Exam Pointers:

• Unchanged in Mass and Chemical Composition: The catalyst's identity remains the same at the end of the reaction.


• Small Amount Required: Even a tiny quantity of catalyst can affect the rate of a large amount of reactants.


• Specificity: Catalysts are often highly specific; one catalyst might accelerate one reaction but not another.


• No Change in Equilibrium: Catalysts speed up both forward and reverse reactions equally. Therefore, they do not alter the position of equilibrium in a reversible reaction; they only help the system reach equilibrium faster. This is a crucial concept for JEE and CBSE.


• Initiation: Catalysts do not initiate a reaction that is thermodynamically not feasible; they only speed up existing reactions.

Understanding catalysts intuitively, especially their role in lowering activation energy and providing alternative pathways, is fundamental. It will help you tackle questions involving reaction energy diagrams and predict reaction outcomes more effectively in your exams.

Prerequisites for Understanding Catalyst Concept and Reaction Mechanisms

To effectively grasp the concept of catalysts and the intricacies of reaction mechanisms, a solid foundation in the fundamental principles of Chemical Kinetics and related topics is essential. Students should be comfortable with the following concepts:

• 
  Basic Chemical Kinetics:
  
  
  
  • 
    Reaction Rate: Understanding what reaction rate means (change in concentration of reactants or products per unit time) and its units.
    
  
  
  • 
    Factors Affecting Reaction Rate: Awareness of how temperature, concentration (or pressure for gases), surface area, and the nature of reactants generally influence reaction rates.
    
  
  
  • 
    Rate Law and Order of Reaction: The ability to write a rate law from experimental data or a given mechanism, determine the order of reaction with respect to individual reactants, and the overall order. Knowledge of zero, first, and second-order reactions.
    
  
  
  
  


• 
  Activation Energy (E_a):
  
  
  
  • A clear understanding of activation energy as the minimum energy required for reactant molecules to transform into products.
  
  
  • The concept of an activated complex or transition state.
  
  
  • How a reaction profile (energy vs. reaction coordinate diagram) is constructed, showing reactants, products, and the transition state. This is crucial as catalysts modify this profile.
  
  
  
  


• 
  Collision Theory:
  
  
  
  • The basic postulates of collision theory: molecules must collide, collisions must have sufficient energy (E_a), and collisions must have proper orientation.
  
  
  • Understanding the role of the frequency factor (A) and exponential factor (e^-Ea/RT) in the Arrhenius equation.
  
  
  
  


• 
  Elementary Reactions and Molecularity:
  
  
  
  • Distinction between an elementary reaction (occurring in a single step) and a complex reaction (occurring in multiple steps).
  
  
  • Understanding molecularity (number of reacting species in an elementary step) and how it differs from the order of reaction (which is experimentally determined for an overall reaction).
  
  
  
  


• 
  Rate Determining Step (RDS):
  
  
  
  • The concept that in a multi-step reaction mechanism, the slowest elementary step dictates the overall rate of the reaction. This step is known as the rate-determining step or rate-limiting step.
  
  
  • Ability to identify the RDS in a proposed mechanism.
  
  
  
  


• 
  Basic Thermodynamics (JEE Focus):
  
  
  
  • A brief idea of enthalpy change (ΔH) for a reaction. It's important to know that catalysts affect the rate, not the thermodynamics (ΔH, ΔG) or equilibrium constant (K_eq) of a reaction.
  
  
  
  

JEE Tip: A strong grasp of these prerequisites, especially Activation Energy and the Rate Determining Step, will significantly simplify your understanding of how catalysts function and how reaction mechanisms are deduced and validated.

Navigating the concepts of catalysts and reaction mechanisms can be tricky, and competitive exams often design questions to exploit common misconceptions. Be vigilant against these traps to secure your marks.

• 
  Misconception: Catalysts shift equilibrium or change K_eq.
  
  

  Exam Trap: Questions might ask about the effect of a catalyst on the equilibrium constant (K_eq) or the position of equilibrium. Students often mistakenly believe catalysts favor product formation, hence shifting equilibrium.

  
  

  Correction: Catalysts do not alter the equilibrium constant (K_eq) or the position of equilibrium. They increase the rate of both the forward and reverse reactions equally, allowing the system to reach equilibrium faster. The final equilibrium composition remains unchanged. This is a crucial concept for both JEE and CBSE.

  
  


• 
  Misconception: Catalysts are consumed or permanently changed.
  
  

  Exam Trap: Problems might present a reaction mechanism where a species acts as a catalyst but appears to be consumed, leading students to mark it as a reactant or product.

  
  

  Correction: A catalyst participates in the reaction mechanism but is regenerated in a subsequent step. Its concentration at the end of the reaction is the same as at the beginning. It lowers the activation energy (E_a) by providing an alternative reaction pathway, but it is not consumed in the overall process.

  
  


• 
  Incorrectly determining the Rate Law from overall stoichiometry.
  
  

  Exam Trap: For multi-step reactions, students often use the stoichiometric coefficients of the overall balanced equation to write the rate law, which is incorrect unless the reaction is elementary.

  
  

  Correction: The rate law for a multi-step reaction is determined by the slowest step, known as the Rate-Determining Step (RDS). Only the molecularity of the elementary steps can be directly used to write their rate laws. If a catalyst is involved in the RDS, its concentration will appear in the rate law.

  
  


• 
  Confusing Reaction Intermediates with Transition States.
  
  

  Exam Trap: Questions often probe the difference between intermediates and transition states, which are both fleeting species in a reaction mechanism.

  
  

  Correction:
  

  
  
  • Reaction Intermediates: These are species formed in one elementary step and consumed in a subsequent step. They have a finite lifetime and can, in principle, be isolated (though often difficult). They correspond to energy minima between two transition states on a reaction profile diagram.
  
  
  • Transition States (Activated Complexes): These are unstable, high-energy arrangements of atoms at the peak of an energy barrier. They have an extremely short lifetime and cannot be isolated. They correspond to energy maxima on a reaction profile diagram. Catalysts lower the energy of the transition state.
  
  
  
  
  
  


• 
  Catalysts and Thermodynamics: No change in ΔG, ΔH, ΔS.
  
  

  Exam Trap: Many questions try to mislead students into believing catalysts affect thermodynamic parameters like Gibbs free energy change (ΔG), enthalpy change (ΔH), or entropy change (ΔS).

  
  

  Correction: Catalysts only affect the kinetics (rate) of a reaction, not its thermodynamics. The initial and final states of the reactants and products remain the same, regardless of the pathway. Therefore, all state functions like ΔG, ΔH, and ΔS for the overall reaction remain unchanged in the presence of a catalyst. This is a frequent trap in both objective and subjective questions.

  
  

Understanding these common pitfalls will help you approach catalyst and mechanism problems with greater precision and avoid losing marks unnecessarily. Always think critically about what a catalyst *does* and *does not* affect.

Understanding the concept of a catalyst and its basic mechanism is fundamental to Chemical Kinetics, frequently tested in both CBSE board exams and JEE Main.

Key Takeaways: Catalyst Concept & Mechanism

• Definition: A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed in the overall reaction. It can be recovered chemically unchanged at the end of the reaction.


• Mechanism of Action: Lowering Activation Energy (E_a)
  
  
  
  • Catalysts provide an alternative reaction pathway or mechanism.
  
  
  • This alternative pathway has a lower activation energy (E_a) compared to the uncatalyzed reaction.
  
  
  • By lowering E_a, a greater fraction of reactant molecules possess the minimum required energy for reaction at a given temperature, leading to a significantly faster reaction rate.
  
  
  • This effect can be visualized on a potential energy diagram, where the catalyst effectively "creates a new, lower mountain pass" for the reaction to proceed.
  
  
  
  


• Effect on Equilibrium:
  
  
  
  • A catalyst does NOT change the position of equilibrium (K_eq). It increases the rate of both the forward and reverse reactions equally.
  
  
  • Its only effect is to help the system reach equilibrium faster.
  
  
  
  


• No Change in Thermodynamics:
  
  
  
  • A catalyst does NOT alter the overall enthalpy change (ΔH) or Gibbs free energy change (ΔG) of the reaction. These are state functions dependent only on the initial and final states, not the pathway.
  
  
  • It only affects the kinetics, not the thermodynamics, of the reaction.
  
  
  
  


• Key Characteristics of Catalysts:
  
  
  
  • Specificity: Many catalysts exhibit specificity, meaning they are effective for only a particular reaction or class of reactions (e.g., enzymes).
  
  
  • Small Amount: Only a small amount of catalyst is generally required to significantly affect the reaction rate.
  
  
  • Intermediate Formation: Catalysts often work by forming unstable intermediate compounds with reactants, which then decompose to form products and regenerate the catalyst.
  
  
  • Catalytic Promoters: Substances that enhance the activity of a catalyst. For example, Mo in the Haber process.
  
  
  • Catalytic Poisons (Inhibitors): Substances that decrease or destroy the activity of a catalyst. For example, CO poisoning of transition metal catalysts.
  
  
  
  


• Types of Catalysis (Basic Idea):
  
  
  
  • Homogeneous Catalysis: Catalyst and reactants are in the same phase (e.g., liquid-phase reactions with an acid catalyst).
  
  
  • Heterogeneous Catalysis: Catalyst and reactants are in different phases, typically a solid catalyst and gaseous or liquid reactants (e.g., hydrogenation of ethene using Ni, Pt, or Pd).
  
  
  
  

JEE & CBSE Focus: Expect questions on the effect of a catalyst on E_a, reaction rate, K_eq, ΔH, and ΔG. Understanding intermediate formation in a catalytic cycle is also important for advanced problems. Distinguishing between promoters and poisons is a common factual question.

Keep these core ideas clear to master this concept for your exams!

Welcome to the "Problem Solving Approach" for catalysts and basic reaction mechanisms! This section will guide you on how to tackle common problems encountered in JEE and CBSE exams.

1. Approach to Catalyst-Related Problems

Catalysts are crucial in chemical kinetics. Understanding their fundamental action helps solve a variety of problems.

• Identify the Catalyst's Role: A catalyst lowers the activation energy (E_a) of a reaction by providing an alternative reaction pathway. It does NOT change the overall enthalpy change ($Delta H$) or the equilibrium constant (K_eq).


• Effect on Reaction Rate:
  
  
  
  • Since E_a is lowered, a greater fraction of molecules possess the necessary energy for reaction, leading to a faster reaction rate.
  
  
  • In rate law expressions, the catalyst generally does not appear in the final, overall rate law if it's regenerated and doesn't affect the stoichiometry, but it can be involved in elementary steps.
  
  
  
  


• Equilibrium Considerations: A catalyst increases the rates of both forward and reverse reactions equally. Therefore, it does not shift the equilibrium position or change the final yield; it only helps equilibrium to be attained faster.


• Types of Catalysis:
  
  
  
  • Homogeneous: Reactants and catalyst are in the same phase.
  
  
  • Heterogeneous: Reactants and catalyst are in different phases (e.g., solid catalyst, gaseous reactants). Adsorption is often a key step.
  
  
  
  


• JEE Specific: Problems might involve comparing reaction rates with and without a catalyst using the Arrhenius equation (k = A e^-Ea/RT). If E_a' is the activation energy with catalyst, then k' = A e^-Ea'/RT.

2. Approach to Basic Reaction Mechanism Problems

A reaction mechanism describes the sequence of elementary steps that constitute the overall reaction. The goal is often to derive the rate law from a proposed mechanism.

• Identify Elementary Steps: Each step in a mechanism is an elementary reaction, meaning its rate law can be directly written from its stoichiometry.


• Locate the Rate-Determining Step (RDS): The slowest step in a reaction mechanism determines the overall rate of the reaction. This is the most crucial step for deriving the rate law.
  
  
  
  • Write the rate law for the RDS based on its molecularity (e.g., if A + B $ o$ Products is RDS, Rate = k[A][B]).
  
  
  
  


• Handle Intermediates: Intermediates are species produced in one elementary step and consumed in a subsequent step. They do not appear in the overall balanced equation or the final rate law.
  
  
  
  • If the RDS involves an intermediate, you must express its concentration in terms of reactants or products.
    
  • Method 1: Pre-equilibrium Approximation (Fast Equilibrium): If an elementary step preceding the RDS is a fast equilibrium, express the intermediate concentration using the equilibrium constant of that step.
    
  • Method 2: Steady-State Approximation (SSA): Assume the concentration of the intermediate remains constant over most of the reaction (rate of formation = rate of consumption). This is more complex and usually involves solving for the intermediate's concentration. (More common in JEE Advanced, less in Main/CBSE).
  
  
  
  


• Verify the Derived Rate Law: Compare your derived rate law with experimental data or options provided. Ensure it only contains species that are reactants, products, or catalysts (if involved in the RDS).


• Distinguish Intermediates vs. Catalysts:
  
  
  
  • Intermediate: Formed and then consumed.
  
  
  • Catalyst: Consumed and then regenerated in a subsequent step, appearing as a reactant in one step and a product in another, maintaining its overall concentration.
  
  
  
  

Mastering these approaches will enable you to confidently tackle problems on catalysts and reaction mechanisms. Always look for the RDS and properly deal with intermediates!

CBSE Focus Areas: Concept of Catalyst and Basic Mechanism Idea

The concept of catalysts is a vital part of Chemical Kinetics for CBSE Board Exams. Students are expected to understand not just what a catalyst is, but also its key characteristics and how it influences reaction rates at a fundamental level, particularly concerning activation energy.

1. Definition of a Catalyst

A catalyst is a substance that alters the rate of a chemical reaction without itself being consumed in the overall reaction. Most commonly, catalysts *increase* the reaction rate (positive catalysts), but some can *decrease* it (negative catalysts or inhibitors).

2. Key Characteristics of a Catalyst (CBSE Important)

For CBSE, understanding these points is crucial:

• 
  Unchanged Chemically and Quantitatively: A catalyst participates in the reaction but is regenerated at the end, remaining unchanged in its chemical composition and mass.
  


• 
  Small Quantity Required: Even a tiny amount of catalyst can catalyze a large amount of reactants.
  


• 
  Specific in Action: A catalyst is usually specific for a particular reaction. For example, MnO₂ catalyzes the decomposition of KClO₃ but not that of H₂O₂.
  


• 
  Does Not Initiate Reaction: A catalyst cannot start a reaction that is not thermodynamically feasible (i.e., spontaneous). It only changes the rate of an existing reaction.
  


• 
  Does Not Change Equilibrium: A catalyst accelerates both forward and backward reactions equally. Thus, it helps in achieving equilibrium faster but does not alter the position of equilibrium or the value of the equilibrium constant.
  


• 
  Optimum Temperature/pH: For enzyme catalysts, there's an optimum temperature and pH at which their activity is maximum. Deviations can lead to denaturation.
  

3. Basic Mechanism Idea: Lowering Activation Energy

The most important concept for CBSE regarding catalyst mechanism is its effect on activation energy (E_a).

• 
  A catalyst provides an alternate reaction pathway or mechanism that has a lower activation energy than the uncatalyzed reaction.
  


• 
  By lowering E_a, a greater fraction of reactant molecules possess the minimum energy required to react.
  


• 
  This leads to an increase in the rate of effective collisions, thereby speeding up the reaction.
  


• 
  Energy Profile Diagrams: Be prepared to interpret and draw energy profile diagrams showing the difference between catalyzed and uncatalyzed reactions. The diagram will show the reactants and products at the same energy levels, but the peak (transition state) will be lower for the catalyzed reaction.
  
  
  CBSE Tip: Always show that the ΔH (enthalpy change) of the reaction remains the same for both catalyzed and uncatalyzed pathways.
  

4. Promoters and Poisons

These terms are also relevant for CBSE:

• 
  Promoters: Substances that enhance the activity of a catalyst. E.g., Mo as a promoter for Fe catalyst in Haber's process for ammonia synthesis.
  


• 
  Poisons: Substances that decrease or destroy the activity of a catalyst. E.g., H₂S poisoning the iron catalyst in Haber's process.
  

CBSE Exam Success Mantra:

Master the definitions, characteristics, and the energy profile diagram explaining the role of activation energy. Focus on conceptual understanding rather than complex mechanistic derivations for board exams. Good luck!

JEE Focus Areas: Concept of Catalyst and Basic Mechanism Idea

Understanding catalysts and their mechanism is crucial for JEE Main, as it frequently involves conceptual questions, energy profile diagrams, and their impact on reaction kinetics.

1. Definition and Role of a Catalyst

• A catalyst is a substance that alters (usually increases) the rate of a chemical reaction without itself being consumed in the overall reaction.


• It participates in the reaction but is regenerated at the end, appearing both as a reactant and product in the mechanistic steps.


• Key Point: A small amount of catalyst can catalyze a large amount of reactants.

2. Basic Mechanism Idea: Lowering Activation Energy (E_a)

The fundamental mechanism by which a catalyst works is by providing an alternative reaction pathway with a lower activation energy (E_a).

• The catalyst forms an intermediate compound with one or more reactants, which then decomposes to yield the products and regenerate the catalyst.


• The energy required to reach the transition state for this alternative pathway is significantly less than that of the uncatalyzed reaction.


• This lower E_a leads to a higher fraction of molecules possessing sufficient energy to react, thus increasing the reaction rate.

Example: Consider an uncatalyzed reaction A + B → AB. A catalyst (C) might facilitate the reaction through a two-step mechanism:

1. A + C → AC (Intermediate, lower E_a1)


2. AC + B → AB + C (Products formed, catalyst regenerated, lower E_a2)

The overall reaction remains A + B → AB, but the path is altered.

3. Characteristics of Catalysts (JEE Relevance)

• Specificity: Catalysts are often highly specific in their action, catalyzing only particular reactions or classes of reactions.


• Small Amount: Only a minute quantity of catalyst is generally required.


• No Change in Equilibrium: A catalyst affects the rate of both forward and reverse reactions equally, hence it does not change the position of equilibrium or the equilibrium constant (K_eq). It only helps in achieving equilibrium faster.


• No Change in Thermodynamics: Catalysts do not alter the change in Gibbs free energy (ΔG), enthalpy (ΔH), or entropy (ΔS) of the reaction. These are state functions determined by the initial and final states, not the path.


• Initiation: A catalyst cannot initiate a reaction that is thermodynamically non-spontaneous (ΔG > 0). It can only speed up an already spontaneous reaction.


• Optimum Conditions: Catalytic activity is often maximal under specific conditions of temperature, pH (for enzymes), and pressure.

4. Effect on Energy Profile Diagrams

Parameter	Effect of Catalyst	JEE Significance
Activation Energy (Ea)	Decreases	This is the primary function of a positive catalyst. Expect questions involving comparing energy diagrams for catalyzed vs. uncatalyzed reactions.
Reaction Rate	Increases	Direct consequence of lower Ea.
Equilibrium Constant (Keq)	No Change	Common JEE Misconception: Catalysts do NOT shift equilibrium.
ΔH (Enthalpy Change)	No Change	Thermodynamic property, independent of reaction pathway.
ΔG (Gibbs Free Energy)	No Change	Thermodynamic property, independent of reaction pathway.

JEE Tip: Always remember that catalysts affect the 'kinetics' (how fast) but not the 'thermodynamics' (how far/spontaneous) of a reaction. Pay close attention to energy profile diagrams in questions, identifying the activation energy barrier for both catalyzed and uncatalyzed pathways.