Identify the correct statement regarding a spontaneous process.
(a) For a spontaneous process in an

Q: Why can't an isolated system exchange energy?

By definition, an isolated system has perfectly insulated and rigid walls, preventing any transfer of heat or work with the surroundings.

Q: Can an endothermic process really be spontaneous?

Yes! Melting of ice at room temperature is endothermic (absorbs heat) but spontaneous because the increase in entropy (solid to liquid) outweighs the energy cost.

Q: Why is option (d) wrong?

Lowering energy (negative $\Delta H$) favors spontaneity, but it's not the *only* criterion. A process can be spontaneous even if energy increases, provided the entropy increases significantly.

Q: What happens to the surroundings in an isolated system?

Nothing! Since the system is isolated, it doesn't interact with the surroundings. Therefore, $\Delta S_{\text{surroundings}} = 0$.

Q: Is $\Delta S_{\text{system}}$ always positive for spontaneous processes?

No, only in an *isolated* system. In an open or closed system, $\Delta S_{\text{system}}$ can be negative as long as $\Delta S_{\text{surroundings}}$ is positive enough to make $\Delta S_{\text{universe}} > 0$.

Question

Identify the correct statement regarding a spontaneous process.

Options:

(A) For a spontaneous process in an isolated system, the change in entropy is positive

(B) Endothermic processes are never spontaneous

(C) Exothermic processes are always spontaneous

(D) Lowering of energy in the reaction process is the only criteria for spontaneity

Sigma Percentile · Accepted Answer

### The Ultimate Decider of Spontaneity

When we talk about chemical thermodynamics, the word "spontaneous" doesn't mean a reaction happens instantly or explodes. It simply means the process has a natural tendency to occur without any continuous external input of energy. But what exactly drives this natural tendency? Let's dive into the fascinating world of isolated systems to find out.

### The Isolated Universe

Imagine a perfectly sealed, rigid, and insulated box. This is our **isolated system**. By definition, an isolated system cannot exchange mass or energy with its surroundings. This means the heat exchange $q = 0$, the work done $w = 0$, and consequently, the change in internal energy $\Delta U = 0$.

Now, if the energy is strictly locked and cannot decrease, what could possibly drive a spontaneous change inside this box? The answer lies in the Second Law of Thermodynamics.

### The Second Law in Action

The Second Law states that for any spontaneous process, the total entropy of the universe must increase:

$$\Delta S_{	ext{universe}} > 0$$

We know that the universe is made up of the system and its surroundings:

$$\Delta S_{	ext{universe}} = \Delta S_{	ext{system}} + \Delta S_{	ext{surroundings}}$$

Here is the beautiful catch: because our system is completely isolated, it has absolutely zero interaction with the surroundings. Therefore, the surroundings experience no change whatsoever, meaning $\Delta S_{	ext{surroundings}} = 0$.

Substituting this back into our equation, we get:

$$\Delta S_{	ext{universe}} = \Delta S_{	ext{system}}$$

Since $\Delta S_{	ext{universe}}$ must be positive for a spontaneous process, it mathematically forces the entropy of the system to be positive:

$$\Delta S_{	ext{system}} > 0$$

This proves that in an isolated system, the **only** driving force for a spontaneous process is the increase in entropy (randomness or disorder).

### Debunking the Energy Myth

Many students fall into the trap of thinking that systems always want to lower their energy (exothermic processes), and that's the only way a reaction can be spontaneous. This is a massive misconception!

The true master equation for spontaneity at constant temperature and pressure is the Gibbs Free Energy equation:

$$\Delta G = \Delta H - T\Delta S$$

For a process to be spontaneous, $\Delta G$ must be negative.

- **Endothermic processes** (where $\Delta H$ is positive) can absolutely be spontaneous if the temperature is high enough and the entropy increases significantly ($\Delta S > 0$). A classic example is the melting of ice at room temperature.
- **Exothermic processes** (where $\Delta H$ is negative) are not always spontaneous. If they lead to a massive decrease in entropy, they might not occur at all.

Therefore, lowering energy is not the only criterion for spontaneity. Entropy is an equally, if not more, powerful player in the grand cosmic dance of thermodynamics.

Animated Solution for Chemistry - Chemical Thermodynamics: Identify the correct statement regarding a spontaneous process.

Select Answer:

Visualized Solution

$The Isolated System$

$Driving Force for Spontaneity$

$System as the Universe$

$Conclusion for Isolated System$

$Evaluating Options$

The Sigma Insight: Entropy and Free Energy

The Ultimate Decider of Spontaneity

The Isolated Universe

The Second Law in Action

Debunking the Energy Myth

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