H2 gas is adsorbed on the metal surface like tungsten. This follows ..... order reaction.
(a) third

Q: Why doesn't the rate increase if we add more hydrogen?

Because the reaction happens on the surface, not in the air. Once all the 'seats' (active sites) on the surface are taken, extra hydrogen molecules just bounce off and have to wait their turn. The rate is bottlenecked by the available surface area.

Q: Is this always a zero-order reaction?

No! It is only zero-order at high pressures when the surface is completely saturated. At very low pressures, there are plenty of empty seats, so adding more gas directly increases the rate, making it a first-order reaction.

Q: What exactly are 'active sites'?

Active sites are specific microscopic locations on the catalyst's surface where the reactant molecules can chemically bind. You can think of them as specialized docking stations that facilitate the reaction.

Q: Does the type of metal matter in this process?

Absolutely. Different metals have different binding energies. Tungsten binds hydrogen very strongly, which makes it relatively easy to saturate the surface, quickly leading to zero-order kinetics.

Q: What is the actual role of the tungsten catalyst here?

The tungsten surface provides an alternative chemical pathway with a much lower activation energy for the reaction (such as the dissociation of $H_2$ into $H$ atoms), thereby speeding up the overall process significantly.

Question

$H_2$ gas is adsorbed on the metal surface like tungsten. This follows ..... order reaction.

Options:

(A) third

(B) second

(C) zero

(D) first

Sigma Percentile · Accepted Answer

### The Magic of Heterogeneous Catalysis

Imagine you are observing a solid piece of tungsten metal placed inside a sealed container filled with hydrogen gas. The hydrogen molecules are in constant, chaotic motion, zipping around and frequently colliding with the walls of the container and the metal itself. When these gas molecules strike the tungsten surface, something fascinating happens: they don't just bounce off; they stick. This phenomenon of gas molecules adhering to a solid surface is known as **adsorption**.

In the realm of heterogeneous catalysis, the actual chemical reaction—whether it is the decomposition of the gas or its reaction with another species—does not occur in the empty space of the bulk gas phase. Instead, it takes place exclusively on the surface of the solid catalyst. Therefore, to understand the kinetics of this reaction, our entire focus must shift from the gas floating around to the microscopic events happening right on the tungsten surface.

### The Movie Theater Analogy

Can these hydrogen molecules stick just anywhere on the metal? Not exactly. The surface of a catalyst is not a uniform, sticky floor. It possesses specific, highly reactive locations known as **active sites**.

You can visualize these active sites as a strictly limited number of seats in a movie theater. A chemical reaction can only proceed if a reactant molecule successfully finds an empty seat and sits down. Consequently, the overall rate of the reaction is directly proportional to the fraction of the surface that is currently occupied by the reactant molecules. In chemical kinetics, we denote this fractional surface coverage by the Greek letter $	heta$ (theta).

### The Saturation Point

Now, let us conduct a thought experiment. What would happen if we drastically increase the pressure of the hydrogen gas inside the container?

A massive crowd of hydrogen molecules will rush towards the tungsten surface. Because the pressure is so high, every single active site on the surface will be rapidly occupied. The theater is completely sold out! At this stage, the fractional coverage $	heta$ reaches its maximum possible value, which is approximately $1$.

If you were to pump even more hydrogen gas into the container at this point, would it speed up the reaction? Absolutely not. There are simply no empty seats left for the new molecules to occupy. The extra gas molecules will just bounce off the saturated surface and return to the gas phase. The surface has reached its absolute **saturation point**.

### The Mathematical Translation

How does this physical reality translate into the mathematics of chemical kinetics? We established earlier that the rate of the reaction depends on the fraction of the surface that is covered:

$$	ext{Rate} \propto 	heta$$

$$	ext{Rate} = k \cdot 	heta$$

However, at high pressures, the entire surface is covered, meaning $	heta$ becomes a fixed constant ($	heta \approx 1$). Substituting this into our rate equation gives:

$$	ext{Rate} = k \cdot (1) = k$$

This mathematical result is profound. It tells us that the rate of the reaction no longer depends on the concentration or the pressure of the hydrogen gas outside. Whether you double the pressure or triple it, the rate remains absolutely constant. It has reached its maximum possible speed.

When the rate of a reaction is completely independent of the reactant's concentration, we express it as:

$$	ext{Rate} = k [H_2]^0$$

Anything raised to the power of zero is one. This is the very definition of a **zero-order reaction**. Therefore, the adsorption and subsequent reaction of a gas on a metal surface at high pressure strictly follows zero-order kinetics.

### The Low-Pressure Twist

Before we conclude, there is a critical nuance you must understand—a favorite trap in competitive exams like JEE. Will this reaction *always* remain zero-order?

The answer is no. What if we drop the pressure of the hydrogen gas to an extremely low value? Under low-pressure conditions, there are very few gas molecules, which means there are plenty of empty seats (active sites) available on the tungsten surface.

In this scenario, the surface is far from being saturated. If you add a little more gas, those new molecules will easily find empty seats, increasing the covered area $	heta$. Mathematically, at low pressures, the coverage $	heta$ is directly proportional to the pressure of the gas ($P_{H_2}$):

$$	ext{Rate} = k \cdot 	heta \propto k \cdot P_{H_2}$$

$$	ext{Rate} = k' [H_2]^1$$

Suddenly, the rate is directly dependent on the concentration of the gas! The reaction has transitioned and now behaves exactly like a **first-order reaction**.

To master this concept, remember this golden rule: **Reactions on catalyst surfaces are first-order at low pressures and zero-order at high pressures.**

Animated Solution for Chemistry - Chemical Kinetics: $H_{2}$ gas is adsorbed on the metal surface like tungsten. This follows ..... order reaction.

Select Answer:

Visualized Solution

$H_{2} Adsorption$

$Active Sites$

$Saturation$

$Rate Equation$

$Zero Order$

$Low Pressure Case$

The Sigma Insight: Order and Molecularity

The Magic of Heterogeneous Catalysis

The Movie Theater Analogy

The Saturation Point

The Mathematical Translation

The Low-Pressure Twist

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Animated Solution for Chemistry - Chemical Kinetics: .math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } H2​ gas is adsorbed on the metal surface like tungsten. This follows ..... order reaction.