Thermodynamics (SAQs)

Chemistry-1 | 6. Thermodynamics – SAQs:
Welcome to SAQs in Chapter 6: Thermodynamics. This page features the key FAQs for Short Answer Questions. Answers are provided in simple English, with a Telugu explanation, and formatted according to the exam style. This will aid in understanding the material and achieving top marks in your final exams.

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SAQ-1 : State and explain “Hess law of constant heat summation” with example.

Introduction

Hess’s Law of Constant Heat Summation states that the total enthalpy change for the reaction is the same, regardless of the number of steps the reaction is carried out in. This law is based on the principle that enthalpy is a state function.

Explanation of Hess’s Law

The law implies that the change in enthalpy for a chemical process is independent of the pathway taken from the initial to the final state. In other words, whether a reaction takes place in one step or multiple steps, the total enthalpy change remains constant. This principle is useful in calculating the enthalpy changes of reactions where direct measurement is difficult.

Application and Example of Hess’s Law

To apply Hess’s Law, one can break down a complex reaction into a series of simpler steps whose enthalpy changes are known. By summing the enthalpy changes of these steps, the overall enthalpy change of the reaction can be determined.

Example: Consider the formation of carbon dioxide (CO₂). It can occur in a single step:

$$C(s) + O₂(g) \rightarrow CO₂(g)$$

Or in two steps:

1. $$C(s) + \frac{1}{2}O₂(g) \rightarrow CO(g)$$
2. $$CO(g) + \frac{1}{2}O₂(g) \rightarrow CO₂(g)$$

By Hess’s Law, the total enthalpy change of these two steps will equal the enthalpy change of the direct formation of CO₂ from carbon and oxygen.

Summary

Hess’s Law is a fundamental principle in thermochemistry. It enables chemists to calculate enthalpy changes for reactions that are difficult to study directly, by using known enthalpy changes of other related reactions.

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SAQ-2 : Define heat capacity? What are C_p and C_v? Show that C_p – C_v = R. 

Introduction

In the realm of thermodynamics, understanding how substances absorb and store heat is crucial. This brings us to the concept of heat capacity and the specific heat capacities at constant pressure (C_p​) and constant volume (C_v​). Here’s a simplified exploration of these ideas and their relationships.

Heat Capacity Explained

Heat Capacity (C): This is a measure of how much heat a substance can absorb for its temperature to rise by 1°C. Mathematically, it’s represented as $$C = \frac{q}{\Delta T}$$ where:

• ‘q’ is the heat absorbed.
• ‘ΔT’ is the temperature increase.

Specific Heat Capacities: C_p​ and C_v​

1. At Constant Volume (C_v​): Imagine you’re heating a substance, but its volume remains unchanged. The heat you add in this scenario, for every 1°C rise, is the heat capacity at constant volume. Its mathematical expression is: $$C_v = \left(\frac{dU}{dT}\right)_v​$$
2. At Constant Pressure (C_p​): Now, think of a situation where you’re heating the substance, but the pressure stays constant. The heat added for every degree rise under this condition is the heat capacity at constant pressure. It’s mathematically defined as: $$C_p = \left(\frac{dH}{dT}\right)_p$$​

The Connection Between C_p​ and C_v​

For a mole of an ideal gas, the enthalpy (H) can be expressed as the sum of its internal energy (U) and the product of pressure (P) and volume (V). Using the ideal gas law (PV = RT), we can rewrite H as: H = U + RT
Differentiating with respect to temperature, we find: $$\frac{dH}{dT} = \frac{dU}{dT} + R$$
Using our earlier definitions of C_p​ and C_v​, we can equate the above to: C_p​ = C_v​ + R This leads to the pivotal relationship: C_p​ − C_v​ = R

Summary

The concept of heat capacity provides insight into how substances store heat, with C_p​ and C_v​ offering more specific scenarios at constant pressure and volume, respectively. The relationship between these specific heat capacities, represented by the equation C_p​ − C_v​ = R, is fundamental in thermodynamics and underscores the consistency in how gases behave under different conditions.

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SAQ-3 : What are intensive and extensive properties?

Introduction

In the study of matter and its properties, it’s important to distinguish between the characteristics that depend on the amount of substance and those that don’t. This distinction leads us to two classifications: intensive and extensive properties.

Intensive Properties Explained

1. Definition: Intensive properties are characteristics of a system that remain unchanged regardless of the amount of substance present.
2. Examples:
   • Density: How compact a substance is.
   • Viscosity: How thick or thin a liquid is.
   • Specific Heat: Amount of heat a substance can absorb without changing its temperature significantly.
   • Temperature: Average kinetic energy of particles in a substance.
   • Pressure: Force exerted by a substance per unit area.
   • Vapor Pressure: Pressure exerted by a vapor in equilibrium with its condensed phase.

Extensive Properties Unveiled

1. Definition: Extensive properties are those that change based on how much of the substance you have.
2. Examples:
   • Mass: Amount of matter in a substance.
   • Volume: Space that a substance occupies.
   • Internal Energy: Total energy (kinetic + potential) of a system.
   • Enthalpy: Heat content of a system.
   • Entropy: Measure of disorder or randomness in a system.
   • Heat Capacity: Heat required to raise the temperature of the entire substance by a certain amount.

Summary

When analyzing the properties of matter, it’s crucial to understand if a particular characteristic is dependent on the quantity (extensive) or is inherent to the substance irrespective of its amount (intensive). Recognizing this difference can significantly impact how we approach and study various physical and chemical processes.

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SAQ-4 : What are open, closed and isolated systems? Give one example for each?

Introduction

In the realm of thermodynamics, systems are often classified based on their ability to exchange matter and energy with their surroundings. This classification leads to the identification of open, closed, and isolated systems.

Open Systems

1. Definition: Open systems can freely exchange both matter and energy with the outside environment.
2. Example:
   • A cup of tea in a saucer: In this setup, not only can heat be transferred between the tea and the environment, but there’s also potential for matter transfer, like tea spilling out or evaporating.

Closed Systems

1. Definition: Closed systems can only exchange energy with the surroundings, but not matter.
2. Example:
   • A chilled sealed drink bottle: While the drink inside can gain or lose heat from/to the surroundings, the matter (liquid inside) remains confined within the bottle, preventing any exchange of liquid or gas with the external environment.

Isolated Systems

1. Definition: Isolated systems are completely cut off from their surroundings. They neither exchange energy nor matter with the outside environment.
2. Example:
   • Hot coffee in a thermos flask: Here, the thermos prevents heat exchange between the coffee and the outside environment. At the same time, because it’s sealed, there’s no way for the coffee to spill or for anything from the environment to get in.

Summary

Understanding the distinctions between open, closed, and isolated systems provides a foundation for many concepts in thermodynamics. Recognizing the characteristics of each type of system and their real-world examples helps in comprehending various physical and chemical transformations in diverse scenarios.

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SAQ-5 : State the third law of thermodynamics. What do you understand by it?

Introduction

In the study of thermodynamics, laws are established that govern the behavior of energy and matter. One such law, known as the third law of thermodynamics, has significant implications for the behavior of substances at extremely low temperatures.

Third Law of Thermodynamics

1. Definition: The energy of a pure and perfectly crystalline substance approaches zero as the temperature gets closer to absolute zero.
2. Implications:
   • Zero Entropy: This essentially means there’s a perfect order or the least possible disorder in the system.
   • Entropy Limitation: The law sets a lower bound on the value of entropy. As a substance heats up, its entropy increases. Conversely, as it cools, its entropy decreases.
   • Calculating Entropy: This law offers a way to determine the entropy (S) of a substance at any temperature. If the temperature dependence of heat capacity (C_p​) is known, it becomes possible to ascertain the absolute value of entropy.
   • Mathematical Representation: The entropy at any given temperature can be represented as: $$S_T = \int_0^T \frac{C_p}{T} dT$$ Here, the equation helps to integrate the temperature-dependent heat capacity over the temperature range to derive the entropy.

Summary

The third law of thermodynamics provides insights into the behavior of substances at very low temperatures, highlighting the significance of order and entropy. It lays the foundation for understanding the absolute values of entropy and offers mathematical tools to determine these values for substances across different temperatures.