THERMODYNAMICS-QnA

Q1: What is thermodynamics?

Thermodynamics is the branch of physics that deals with the relationship between heat, work, and internal energy of a system without considering its microscopic details.

Q2: Define thermodynamic system.

A thermodynamic system is a specified quantity of matter or a region of space chosen for study, separated from surroundings by a real or imaginary boundary.

Q3: What is a closed system?

A closed system can exchange energy with surroundings but does not allow transfer of mass across its boundary.

Q4: Define internal energy.

Internal energy is the total microscopic energy of a system due to random motion and interaction of its particles.

Q5: What is a thermodynamic process?

A thermodynamic process is the transformation of a system from one equilibrium state to another.

Q6: State the SI unit of heat.

The SI unit of heat is joule (J).

Q7: What is an isothermal process?

An isothermal process is one in which the temperature of the system remains constant throughout the process.

Q8: Define adiabatic process.

An adiabatic process is one in which no heat is exchanged between the system and surroundings.

Q9: What is thermodynamic equilibrium?

A system is in thermodynamic equilibrium when it is simultaneously in thermal, mechanical, and chemical equilibrium.

Q10: What is state variable?

A state variable is a physical quantity whose value depends only on the state of the system, not on the path followed.

Q11: Name a path variable.

Heat and work are examples of path variables.

Q12: What does PV diagram represent?

A PV diagram represents the variation of pressure with volume during a thermodynamic process.

Q13: Define cyclic process.

A cyclic process is one in which the system returns to its initial state after passing through a series of changes.

Q14: What is absolute zero?

Absolute zero is the lowest possible temperature at which molecular motion is minimum.

Q15: Write the mathematical form of first law of thermodynamics.

?Q = ?U + ?W

Q1: Explain the significance of internal energy.

Internal energy accounts for all microscopic forms of energy within a system and determines how heat and work affect its temperature and state.

Q2: Distinguish between isothermal and adiabatic processes.

In isothermal processes temperature remains constant with heat exchange, while in adiabatic processes temperature changes with no heat exchange.

Q3: Why is work a path function?

Work depends on the specific manner in which the process occurs, not just on initial and final states.

Q4: What happens to internal energy in an isothermal process of ideal gas?

The internal energy remains constant because it depends only on temperature, which does not change.

Q5: Explain quasi-static process.

A quasi-static process proceeds infinitely slowly so that the system remains nearly in equilibrium at every stage.

Q6: Why is specific heat a state function?

Specific heat depends only on the state of the system and the nature of substance, not on the path followed.

Q7: Define heat capacity.

Heat capacity is the amount of heat required to raise the temperature of a body by one degree.

Q8: What is work done in free expansion?

In free expansion, work done is zero because there is no opposing external pressure.

Q9: What is a reversible process?

A reversible process is one that can be reversed exactly by infinitesimal changes without leaving any effect on surroundings.

Q10: State one limitation of first law.

The first law does not indicate the direction in which a thermodynamic process can occur.

Q1: Explain the first law of thermodynamics.

The first law states that heat supplied to a system is partly used to increase internal energy and partly to perform external work, expressing energy conservation.

Q2: Derive work done in an isothermal process.

For an ideal gas at constant temperature, work done equals nRT ln (V2/V1), obtained by integrating PdV using ideal gas equation.

Q3: Explain adiabatic process with equation.

In an adiabatic process, no heat is exchanged and the relation \(PV^{\gamma}\) = constant holds, where \(\gamma\) is ratio of specific heats.

Q4: Explain state and path variables with examples.

State variables depend only on the state of system like pressure and temperature, while path variables depend on the process such as heat and work.

Q5: Explain cyclic process using PV diagram.

In a cyclic process, the PV diagram forms a closed loop, and net work equals the area enclosed by the loop.

Q1: Describe the first law of thermodynamics in detail.

The first law of thermodynamics is a fundamental statement of the principle of conservation of energy applied to thermodynamic systems. It states that energy can neither be created nor destroyed; it can only be transferred from one form to another. When a system undergoes a thermodynamic process, the energy supplied to it in the form of heat does not vanish but appears partly as work done by the system and partly as a change in its internal energy.

Mathematically, the first law of thermodynamics is expressed as:

\[ \begin{aligned} \Delta U &= Q - W \end{aligned} \]

Here, \(\Delta U\) represents the change in internal energy of the system, \(Q\) denotes the heat supplied to the system, and \(W\) is the work done by the system on its surroundings. If heat is supplied to the system, \(Q\) is taken as positive, and if heat is rejected, \(Q\) is negative. Similarly, work done by the system is considered positive, while work done on the system is negative.

Internal energy is the total microscopic energy possessed by a system due to the random motion and interactions of its molecules. It depends only on the state of the system and not on the path by which the system reaches that state. Therefore, a change in internal energy depends solely on the initial and final states of the system.

The physical interpretation of the first law is that whenever heat is supplied to a system, it is used in two possible ways. One part of the heat increases the internal energy of the system, which may result in a rise in temperature or a change of state, while the remaining part is used to perform external work. If no heat is supplied and the system still does work, the work is performed at the expense of its internal energy.

Thus, the first law of thermodynamics provides a quantitative relationship between heat, work, and internal energy, and it establishes that all thermodynamic processes strictly obey the law of conservation of energy.

Q2: Explain internal energy and factors affecting it.

Internal energy is the total energy possessed by a thermodynamic system due to the microscopic motion and interactions of its constituent particles. It arises from the kinetic energy associated with the random translational, rotational, and vibrational motion of molecules, as well as the potential energy resulting from intermolecular forces. Internal energy does not include the macroscopic kinetic energy of the system as a whole or the potential energy due to its position in an external field.

For a given system, internal energy is a state function, which means its value depends only on the present state of the system and not on the path by which that state is achieved. As a result, a change in internal energy between two states is independent of the process involved and is determined solely by the initial and final conditions of the system.

In the case of an ideal gas, internal energy depends only on temperature. When the temperature of an ideal gas increases, the average kinetic energy of its molecules increases, leading to an increase in internal energy. This relationship may be expressed mathematically as:

\[ \begin{aligned} \Delta U &= n C_V \Delta T \end{aligned} \]

Here, \(n\) is the number of moles of the gas, \(C_V\) is the molar heat capacity at constant volume, and \(\Delta T\) is the change in temperature. For real gases and other systems, internal energy may also depend on volume and pressure because intermolecular forces contribute significantly to the potential energy of the system.

The factors affecting internal energy therefore include the temperature of the system, which directly influences molecular motion, the nature of the substance, which determines the strength of intermolecular interactions, and the physical state of the system, such as solid, liquid, or gas. In addition, the amount of substance present also affects internal energy, since a larger number of particles implies a greater total microscopic energy.

Hence, internal energy represents the microscopic energy content of a system and is governed primarily by temperature, the nature of the material, and the quantity and state of the substance.

Q3: Discuss isothermal and adiabatic processes comparatively.

An isothermal process is a thermodynamic process in which the temperature of the system remains constant throughout the transformation. This is possible only when the system is allowed to exchange heat freely with its surroundings. During such a process, any heat absorbed by the system is exactly used to perform external work, so there is no change in internal energy. For an ideal gas, this condition is expressed mathematically by the relation:

\[ \begin{aligned} \Delta U &= 0 \\ Q &= W \end{aligned} \]

As the temperature remains constant, the pressure and volume of the gas change in accordance with Boyle’s law, and the pressure–volume curve for an isothermal process is a smooth hyperbola. The system must be transformed slowly so that thermal equilibrium with the surroundings is maintained at every stage.

An adiabatic process, on the other hand, is a process in which no heat is exchanged between the system and its surroundings. This may occur when the system is perfectly insulated or when the process takes place very rapidly, leaving no time for heat transfer. In an adiabatic process, the work done by the system is entirely at the expense of its internal energy, leading to a change in temperature. This behavior is described by:

\[ \begin{aligned} Q &= 0 \\ \Delta U &= -W \end{aligned} \]

For an ideal gas undergoing an adiabatic change, pressure, volume, and temperature are related by a different equation, and the pressure–volume curve is steeper than that of an isothermal process. As a result, for the same change in volume, the pressure falls more rapidly in an adiabatic expansion than in an isothermal one.

Comparatively, the essential difference between the two processes lies in heat exchange and temperature variation. In an isothermal process, temperature remains constant due to continuous heat exchange, whereas in an adiabatic process, the absence of heat transfer causes the temperature to change. Both processes play an important role in understanding the working of heat engines and natural thermodynamic phenomena.

Q4: Explain work done in expansion and compression of gas.

In thermodynamics, work is said to be done by a gas when it expands against an external pressure, and work is done on the gas when it is compressed. Consider a gas enclosed in a cylinder fitted with a frictionless movable piston. When the gas expands, the piston moves outward due to the force exerted by the gas pressure, resulting in displacement. This displacement under the action of force constitutes mechanical work.

If the pressure exerted by the gas at any instant is \(P\) and the corresponding small change in volume is \(dV\), the infinitesimal work done by the gas during expansion is given by:

\[ \begin{aligned} dW &= P\,dV \end{aligned} \]

To find the total work done when the volume of the gas changes from an initial value \(V_1\) to a final value \(V_2\), the above expression is integrated over the given limits. Thus, the work done in expansion is expressed as:

\[ \begin{aligned} W &= \int_{V_1}^{V_2} P\,dV \end{aligned} \]

In an expansion process, the volume increases, so \(dV\) is positive, and hence the work done by the gas is positive. Physically, this means that the gas transfers energy to the surroundings by pushing the piston outward.

In the case of compression, the external pressure is greater than the pressure of the gas, and the piston is pushed inward, reducing the volume of the gas. Here, the small change in volume \(dV\) is negative. As a result, the work done by the gas becomes negative, indicating that work is done on the gas by the surroundings. The same mathematical expression applies, but the sign of work changes due to the decrease in volume.

Thus, the work done in expansion and compression of a gas depends on the pressure-volume behavior of the system and the nature of the thermodynamic process. The concept is fundamental to understanding energy transfer in engines, compressors, and other thermodynamic devices.

Q5: Explain limitations of first law of thermodynamics.

The first law of thermodynamics establishes a fundamental energy balance by relating heat, work, and internal energy, yet it has important conceptual and practical limitations. While the law states that energy is conserved during any thermodynamic process, it does not provide information about the direction in which a process will occur. For example, it cannot explain why heat naturally flows from a hotter body to a colder one and never spontaneously in the reverse direction, even though both directions would satisfy energy conservation.

Mathematically, the first law is written as:

\[ \begin{aligned} \Delta U &= Q - W \end{aligned} \]

This equation accounts only for the quantity of energy involved and not for the feasibility or spontaneity of a process. As a result, the first law alone cannot distinguish between reversible and irreversible processes, nor can it predict whether a given process can occur under real physical conditions.

Another major limitation is that the first law places no restriction on the conversion of heat into work. According to it, it appears possible to convert the entire amount of heat absorbed by a system into work, provided the energy balance is maintained. However, practical experience shows that no heat engine can achieve one hundred percent efficiency. The first law fails to explain this inherent limitation on energy conversion.

Furthermore, the first law does not address the quality or usefulness of energy. It treats all forms of energy as equivalent, even though some forms, such as mechanical work, are more useful than others, such as low-temperature heat. It also cannot explain why certain processes are irreversible in nature, despite satisfying energy conservation.

Thus, although the first law of thermodynamics is essential for energy accounting, it is incomplete in explaining natural processes. These shortcomings necessitate the second law of thermodynamics, which introduces the concepts of direction, irreversibility, and efficiency in thermodynamic phenomena.