What is Thermodynamic System and Surroundings, Thermodynamic Process, State of System, Point and Path of Function, Extensive and Intensive Properties, Latent, Specific and Sensible Heat

System and Surroundings:

In thermodynamics, the system and its surroundings are fundamental concepts used to analyze and describe the behavior of physical systems. Here's a brief explanation of these concepts: 

System: 

A system is the specific part of the universe that is of interest for study or analysis. It can be as simple as a gas contained within a cylinder or as complex as an entire chemical plant. 

Surroundings: 

The surroundings refer to everything outside the system that interacts with it. This includes anything that can exchange energy or matter with the system. The surroundings provide the external conditions that can influence the behavior of the system. For example, in the case of a gas in a cylinder, the surroundings would include the atmosphere outside the cylinder and the walls of the container. 

Imagine you're making a cup of coffee. The hot cup of coffee is system itself. This is the part of the universe you have to study. Everything else except that surrounds the cup is surrounding. That surrounding may be room or entire universe. Thermodynamics is simply the study of how energy i.e. heat and mass move between your System and its Surroundings. 

In thermodynamics, the relationship between the system and its surroundings is crucial for understanding how energy and matter flow in and out of the system, and how the system responds to changes in its environment. The study of these interactions forms the basis of thermodynamics and helps in predicting and analyzing the behavior of various physical systems.

What is boundary? 

A system is typically enclosed by a boundary, which can be real or imaginary, that separates it from its surroundings. The boundary can be fixed or movable, and it can allow for the transfer of energy or matter between the system and its surroundings. The boundary is the wall that you mentally draw to separate the System that you're studying from the Surroundings or everything else. It defines what can and cannot cross between the system and the surroundings. 

The physical nature of the boundary determines what kind of system you have. The crucial events happens at boundary like heat transfer occurs across the boundary. Work done involves the boundary moving. Mass flow happens through the boundary. 

• The boundary can be fixed or movable. 

• Fixed boundary is like a rigid steel box. The size and shape of the system cannot change. If you heat the gas inside, its pressure will increase, but the volume stays the same because the boundary is fixed. 

• While moveable boundary is like a piston in a cylinder or a balloon's skin. The system can change size. If you heat the gas inside a piston, the boundary i.e. piston head will move outward. The system has done work on the surroundings by pushing the piston. 

Boundary can be real or imaginary.

• Real boundary is like a thing you can touch. It's the actual container holding your system like skin of a balloon.

• Imaginary boundary is like you draw a line around a part of the universe to study it, even though there's no physical wall. It's a concept that helps you focus your analysis. It's like you imagine a box in the middle of the room. The air inside the box is your system; the air outside is the surroundings. 

While other boundary types depends on what it allows to pass.

• It is impermeable boundary if mass cannot cross this boundary. If mass cross boundary then it is known as permeable boundary.

• It is adiabatic boundary if boundary does not allow heat to pass through it. It is diathermal boundary if boundary allows heat to flow through it easily. 

Thermodynamic systems are classified as : 

Open System:

If the thermodynamic system has the capacity to exchange both matter and energy with its surroundings, it is said to be an open system.

e.g. A steam turbine 

Closed System:

A system which has the ability to exchange only energy with its surroundings and cannot exchange matter is known as a closed system.

e.g. A cylinder in which the valve is closed is an example of a closed system. When the cylinder is heated or cooled, it does not lose its mass. 

Isolated System:

A system which cannot exchange matter or energy with the surroundings is known as an isolated system. The zeroth law of thermodynamics states that thermodynamic processes do not affect the total energy of the system.

e.g. If the piston and cylinder arrangement in which the fluid like air or gas is being compressed or expanded is insulated, it becomes an isolated system.

Thermodynamic process:

A thermodynamic process is simply how a system changes from one state to another. It's the journey, not the destination.

Imagine you are driving from your home (State 1) to your friend's house (State 2). Your home and your friend's house are two different states, defined by their addresses like pressure and temperature define a state. The route you take and how you drive is the process. Did you take the highway or got stuck in traffic? Did you drive fast or slow? All of this is the process.

In thermodynamics, types of processes are defined by what we keep constant during the journey. It may be constant speed or constant altitude.

1. Isobaric Process (Constant Pressure):

An isobaric process is simply a change that happens under constant pressure. Iso means same and baric means pressure.

Imagine a cylinder with a gas inside, sealed by a piston that can move up and down without any friction. Now, place a constant weight on top of the piston. The weight provides a constant downward force. For the system to be in equilibrium, the pressure inside must always equal the weight pushing down, so it stays constant. The gas inside must push upward with an equal force. The piston's area are both constant, the pressure inside must also remain constant, no matter what else happens.

You start to heat the cylinder with a flame. When you add heat, the gas molecules gain energy and move faster. The upward force from the gas temporarily becomes greater than the downward force from the weight. This causes the piston to move upward. The gas expands. As the gas expands, the molecules have more space. Their collisions with the piston become less frequent. Once the upward force decreases until it once again exactly balances the downward force from the weight. As you heat the system, it expands or contracts, but the pushing force i.e. pressure remains the constant the entire time.

e.g. The pot is open to the atmosphere. The atmosphere is like a giant, constant weight pushing down on the water. As the water heats and turns to steam, it expands and escapes, but the entire process occurs at constant atmospheric pressure.

2. Isochoric Process (Constant Volume):

An isochoric process is simply a process that happens under constant volume. Iso means same and choric means volume. The volume cannot change.

Imagine a strong and closed metal tank filled with gas. The walls are so rigid and thick that they cannot bend, or expand. When you add heat to this tank by putting it on a flame. The gas molecules inside absorb the heat energy and start moving much, much faster. They are vibrating around with more kinetic energy. Normally, molecules would use this extra energy to push outward and expand by increasing volume. The walls of the tank won't bend. Since the molecules are moving faster but are trapped in the same small space, they start colliding with the walls far more frequently and violently. This causes increase in pressure. The only things that can change are Pressure (P) and Temperature (T). When you remove heat pressure and temperature decrease. This increase and decrease in value of pressure and temperature happens at constant volume.

e.g. Heating a sealed, rigid container of food in a microwave. The container's size doesn't change. The heat causes the pressure inside the can to rise extremely quickly. If it exceeds the can's strength, it will explode.

3. Isothermal Process (Constant Temperature):

An isothermal process is a change that happens at a constant temperature. Iso means same and thermal means heat. It's a transformation where the hotness or internal energy level of the system doesn't change.

Imagine a gas inside a cylinder with a piston. The entire cylinder is submerged in a giant, perfect thermal reservoir like a huge ocean of water at a fixed temperature i.e. 60°C. This reservoir acts as a perfect thermostat. When we slowly compress the piston.

As you compress the gas, you're doing work on it. This work adds energy to the gas molecules, making them move faster. If this were done in isolation, the temperature would rise. The moment the molecules start moving faster, their temperature tries to rise above the reservoir's temperature (50°C). Heat immediately flows out of the gas and into the surrounding reservoir to maintain equilibrium. This keeps the gas at exactly 50°C. The energy you put into the gas by compressing it i.e. work is perfectly balanced by the energy that flows out as heat. The net change in the gas's internal energy is zero. Its temperature stays constant.

In reverse if you let the gas expand very slowly. The gas does work on the piston, which requires energy. To get this energy, the gas's molecules would normally slow down, cooling the gas. But the thermostat prevents this. Heat flows from the reservoir into the gas to keep its temperature at a constant 50°C. The energy that flows in as heat is perfectly balanced by the energy that flows out as work.

The process happens so slowly that the system has time to adjust and stay at the same temperature throughout. To keep temperature constant during expansion, you must add heat. To keep it constant during compression, you must remove heat.

e.g. A slow, controlled phase change, like ice melting at 0°C or water boiling at 100°C. The temperature doesn't change until the entire process is complete. Imagine melting of ice in a glass of water at 0°C. As you add heat, the ice melts into water, but the temperature of the ice water mixture remains stubbornly at 0°C until all the ice is gone. The added heat energy is used to break the ice's bonds (doing work) rather than raising its temperature.

4. Adiabatic Process (No Heat Transfer):

An adiabatic process is a change that happens so fast and in such good insulation that no heat has time to enter or leave the system. It's a transformation where the system is completely self-reliant; it can only use its own internal energy.

Imagine our classic cylinder with a piston, but this time it's perfectly sealed and wrapped in the insulation. The system is perfectly insulated. No heat can enter or escape.

When you very rapidly compress the piston. You are doing work on the gas, adding energy to it. The process happens so fast that the added energy has no time to escape as heat through the insulation. The energy is trapped inside. The only place for this extra energy to go is into the gas molecules themselves. They start moving much faster. The gas gets hotter. Its temperature rises significantly, purely because of the work you did on it. The pressure also skyrockets because the molecules are both hotter and in a smaller space.

In reverse if you let the gas expand very rapidly. The expanding gas uses its own internal energy to push the piston outward and do work on the surroundings. The expansion happens so fast that no heat can flow in from the outside to replenish this lost energy. The only source of energy is the gas's own internal energy. The molecules have to slow down to provide this energy. The gas gets colder. Its temperature drops significantly.

The entire process, compression or expansion happened with zero heat transfer (Q = 0). All change in the system's energy comes only from work being done on it or by it. If workdone is quickly compression, the work done on it increases its energy, so its temperature rises dramatically like a bike pump getting hot. If workdone is quickly expansion, It does work on the surroundings, uses its own internal energy, and its temperature drops dramatically like spray from a deodorant can feeling cold.

e.g. When you pump up a tire quickly, the bottom of the pump gets very hot. You are doing work to compress the air rapidly. The process is fast enough that the heat doesn't have time to dissipate into the pump's metal walls, so the temperature of the air rises adiabatically.
When you spray from a deodorant can, the gas inside expands rapidly as it leaves the nozzle. This expansion does work on the surrounding air, using the gas's internal energy. This causes the remaining gas in the can and the nozzle itself to get very cold. You can often feel the chill.

State, Point and Path Function:

State of system:

The state of a system is like its complete identity card at a specific moment in time. It's the list of all its properties like temperature, pressure, volume, etc. that together uniquely describe what that system is like, at that time, right now. If you know the state, you know everything you can know about the system without having to know its history.

Imagine you have system like a gas inside a piston. The specific details visible in that some specific time are the state variables i.e. the properties that define the state. How much piston is expanded i.e. volume, how much hard is the gas pushing on the walls i.e. pressure, how fast are the molecules moving i.e. temperature, how much stuff is in there i.e. mass. The combination of these specific values like V=2 liter, P=200 kPa, T=300 K, n=0.04 moles, are the state of the gas in that instant.

The laws of thermodynamics are written in terms of changes in state functions (ΔU, ΔS). We can calculate the energy required to go from State 1 (cold water) to State 2 (steam) without caring about how we heated it. It may be on a stove or with a laser. P-V Diagrams are maps of possible states. A line on this diagram doesn't represent a path, but a series of equilibrium states the system passes through.

You don't need to list every single property to define a state. For a simple system like a gas, specifying just two independent intensive properties like Temperature and Pressure are enough to fix all other intensive properties like density, refractive index. T and P alone are sufficient only for intensive properties, not extensive ones. Temperature and pressure define the intensive state i.e. what kind of but we still need one extensive property like mass or volume to know the full extensive state i.e. how much. Once you set T and P, the intensive state is defined. If you also know one extensive property like mass or volume, you know everything about the system.

Thermodynamics simplifies reality by reducing the number of variables needed, thanks to the state postulate. This is why engineers and scientists can work with just a few measured properties. In Practice, an engineer designing a power plant doesn't need to individually measure the density,internal energy, enthalpy, and entropy of the steam at every point in the system. They simply measure Pressure and Temperature at various points. Using property tables like steam tables or equations of state, they can simply look up all the other intensive properties they need for their calculations like density, enthalpy, entropy for that specific P and T.

Point Function:

Point function is a property that only depends on your current state or position, not on how you got there. In Thermodynamics, properties like Pressure (P), Volume (V), Temperature (T), and Internal Energy (U) are point functions.

Imagine you are hiking up a mountain to a summit cabin. Your altitude is a point function. When you finally reach the top of mountain, your altitude is 3000 meters. No matter how you went there. You take short, steep trail, you are still at 3000 m. You take the long, scenic loop, you are still at 3000 m. The altitude only cares where you are, not your history.

The internal energy of a gas in a cylinder depends only on its current state like its current P, V, and T. It doesn't matter if you heated it quickly with a blowtorch or slowly over a candle to get it to that state. The energy contained within it is the same. Because the value only depends on the start and end points, the change in a point function is what matters.

The change in altitude from the base at 500m to the summit at 3000m is always 2500 m, no matter what path you take. Similarly, the change in internal energy ΔU of a system between two states is always the same, regardless of how the change happened.

Path Function:

Path function is a property that depends entirely on the path you took to get from the start to the end. In Thermodynamics, properties like Heat (Q) and Work (W) are path functions. You cannot say a system has a certain amount of heat or work inside it. Heat and work are modes of energy in transit. They describe the journey, not the destination.

In the hike, the distance you hiked is a path function. The distance from the base to the summit cabin is not a fixed number. You take short, steep trail, you might hike 5 km. You take the long, scenic loop, you might hike 12 km. The distance hiked is completely determined by your path.

You can get a gas to a certain temperature i.e. a state by adding a little heat and doing a lot of compression work on it, OR by adding a lot of heat and letting it expand and do work. The final state is the same, but the amounts of Heat (Q) and Work (W) that flowed during the process were completely different for each path.

Look at The First Law of Thermodynamics, ΔU = Q - W
• ΔU (Change in Internal Energy) is a Point Function. Its value is fixed for a given change in state.
• Q (Heat added) and W (Work done) are Path Functions. Their values can change dramatically.

This means there are infinitely many ways to achieve the same change in internal energy (ΔU). You can add a lot of heat (Q) and let the system do a lot of work (W). Or you can add no heat (Q=0) and only do work on it, meaning (-W). The difference (Q - W) will always be the same for a given ΔU, but the individual values of Q and W depend on the path.

We can analyze complex processes by simply comparing the start and end points, without worrying about the messy details in between.

Difference between state and point function:

State is the overall description of the system, while point function is a type of
property that defines that state. State is like the complete GPS coordinates including altitude, while point functions are the individual coordinates themselves like latitude, longitude, altitude. All point functions are state functions, and they work together to fully define a state. You use point functions like P, V, T to describe a state.

It's like person's complete profile vs their specific attributes like height, weight. A State is the complete description of what the system is like at a moment in time. It's the whole person. A Point Function or State Function is a property that helps to define that state. It's a specific attribute of that person, like their height or weight.

You can measure any point function and it will give you a single value that describes an aspect of the current state. If you know enough point functions (e.g., height, weight, and temperature), you have a very good description of the person's overall state.

Imagine gas in a piston. To describe it's state, we list its Point Functions like
Pressure = 200 kPa
Temperature = 300 K
Volume = 1 m³
Internal Energy = 10,000 J
The combination of [P=200 kPa, T=300 K, V=1 m³, U=10,000 J] is the state.

Extensive and intensive properties:

Extensive Properties:

Extensive properties depend on the amount of matter. They are additive. If you combine two identical systems, all extensive properties double. Extensive properties tell you the scale of a process. How big does your chemical reactor need to be to produce 1000 kg of product per day?

e.g. Mass, Volume, Number of Moles, Internal Energy, Enthalpy, Entropy, Gibbs Free Energy.

Intensive Properties:

Intensive properties do not depend on the amount of matter. They describe the intrinsic quality of the material. They are not additive. If you cut a system in half, all intensive properties remain the same. Intensive properties tell you the conditions needed to run the process. At what temperature and pressure does the reaction work best? You can design a small lab experiment using the right intensive properties, and then scale it up to a full-sized factory by focusing on the extensive properties.

e.g. Temperature, Pressure, Density, Specific Properties like Specific Heat Capacity (The property per unit mass. e.g., Heat needed per gram to raise the temperature by 1°C. The specific heat of water is ~4.18 J/g°C, whether for a drop or an ocean), Molar Properties like Molar Volume (The property per mole. e.g. Volume occupied by one mole of substance. One mole of an ideal gas always occupies 22.4 L at standard conditions, regardless of the gas type)


1. Imagine you have a cup of hot water. Extensive properties depends on how much water you have. The total volume of water If you pour two cups together, you have twice the volume.
While Intensive properties depend on what kind of water it is. If you pour two cups of the same hot cup of water together, the temperature doesn't double; it stays the same.

2. If you cut the cake in half or combine two cakes, these values change. The weight, mass and volume of the whole cake will change. Imagine cake is hot. If you cut it in half each half is still just as hot. Temperature will not change even you combine two cakes. Intensive properties tell you what state the material is in. e.g we define the boiling point of water (an intensive property, 100°C) regardless of whether it's in a teaspoon or a swimming pool.

Latent, Specific and Sensible Heat:

Latent Heat:

Latent heat is the energy needed to change a substance's state like from solid to liquid without changing its temperature. It's the energy used for rearranging molecules, not for making them move faster.

• When water on your skin evaporates, it pulls the latent heat of vaporization from your skin to power its change into a gas. This cools you down.
• Hurricanes get their immense energy from the latent heat released when water vapor in the atmosphere condenses into rain.
• A steamer is very efficient because the steam hitting the food releases a huge amount of latent heat as it condenses.

At the atomic level, the total energy you add to a system is distributed in two ways.
• Energy that makes molecules move or vibrate faster is kinetic energy. This is what we measure as temperature.
• Energy used to overcome the attractive forces (bonds) between molecules is potential energy. This energy is stored in the new configuration of the molecules and does not increase their speed. This is latent heat.

The Two Main Types of Latent Heat:

1. Latent Heat of Fusion (Solid → Liquid):

In a solid, molecules are locked in a rigid, crystalline structure by powerful intermolecular bonds. They vibrate around fixed positions but cannot move freely. When we add energy, initially it increases the vibrational kinetic energy of the molecules, raising the temperature. At the melting point, the molecules have enough kinetic energy to violently vibrate. Now, any additional energy we add is not used to vibrate faster. Instead, it is used to break the rigid bonds of the crystal lattice and to push neighboring molecules apart, creating space for molecules to slide past one another. This input energy increases the potential energy of the system. The molecules are now in a state where they are still close and attracted to each other (a liquid), but they have more freedom of movement. Their average kinetic energy (temperature) remains constant throughout this entire process. The energy spent on breaking bonds and creating space is the latent heat of fusion.

In reverse process i.e. freezing, when a liquid freezes, the molecules settle into a stable, ordered lattice. The potential energy they gained during melting the latent heat—is released back into the surroundings as thermal energy.

e.g. An ice cube at 0°C sits in a drink. The drink is at 0°C. When you add heat energy, the ice cube does not get warmer. Instead, it slowly melts into 0°C water. All the energy you're adding is going into breaking the ice's crystal lattice.

2. Latent Heat of Vaporization (Liquid → Gas):

In a liquid, molecules are close together and held by strong attractive forces (cohesion). They can slide past each other but are constantly colliding and interacting. When we add energy, initially it increases the translational and rotational kinetic energy of the molecules, raising the temperature. At the boiling point, the molecules are moving very quickly. Any additional energy is now used to overcome the much stronger cohesive forces that still bind the molecules in the liquid phase. This requires significant energy. It is also used to perform work against the external pressure to push the surrounding atmosphere aside and allow the molecule to escape into the vast space of the vapor phase. This is a massive increase in volume. This input energy drastically increases the potential energy of the system. A molecule in the gas phase is far from its neighbors and has very little potential energy from intermolecular forces. The energy required to achieve this separation and overcome external pressure is the latent heat of vaporization. Again, the average kinetic energy of molecules remaining in the liquid phase does not increase during this change.

In reverse process i.e. Condensation, when a gas condenses, a molecule moving towards the liquid surface is captured by the attractive forces of the liquid. It loses a large amount of potential energy, which is released as thermal energy. This is why steam burns are so severe, it's not just the heat, but the massive release of latent heat upon condensation.

e.g. A pot of water at 100°C is boiling. When you add more heat from the stove. The water does not get hotter than 100°C. Instead, it turns into steam. water vapor at 100°C. All the energy you're adding is going into giving the molecules enough energy to escape into the air. It takes much more energy than melting.

3. Latent Heat of Sublimation:

This is the energy needed to go directly from a solid to a gas, skipping the liquid phase entirely.
e.g. dry ice.

Specific Heat:

When you add a quantity of heat energy (Q) to a substance, that energy is distributed among the atoms and molecules. Specific heat is the measure of how much energy is required to raise the temperature of a unit mass of a substance by one degree. Specific heat tells you how much energy a substance needs to pack to change its temperature. The reason different substances require different amounts of energy for the same temperature increase lies in how the added energy is partitioned among the available degrees of freedom.

At the atomic level, the added thermal energy stored as

• Translational Kinetic Energy is energy that is associated with the molecule's velocity as it moves through space.
• Rotational Kinetic Energy is energy that is associated with the molecule's rotation around its center of mass.
• Vibrational Potential and Kinetic Energy is energy that is used to stretch and bend the bonds between atoms within a molecule.
• Intermolecular Potential Energy is energy that is used to overcome weak attractive forces between molecules, slightly increasing their average separation.

Types of Specific Heat:

• Specific Heat at Constant Volume (Cᵥ):

The substance is heated in a rigid, fixed volume container.

Because the volume is fixed, when you add heat energy (Q). The atoms/molecules cannot expand and do work on their surroundings. No energy is lost to pushing a piston or expanding boundaries. All of the added energy (Q) goes directly into increasing the internal energy of the system. This energy is distributed among all available molecular degrees of freedom i.e. translation, rotation, vibration. The result is a pure increase in molecular motion and interactions, leading to a rise in temperature.

In essence, Cᵥ measures the pure energy-storage capacity of the substance. It tells you how much energy is needed to raise the temperature when every single joule of input is used solely to energize the molecules.

The value of Cᵥ itself depends on the molecular structure:

• In Monatomic Gases like Argon, Helium, atoms are single points. They can only store energy via translation i.e. movement in x, y, z directions. They have very low specific heat because there are few degrees of freedom to feed.
• In Diatomic Gases like N₂, O₂, molecules are like dumbbells. They can store energy via translation and rotation means they can spin around two axes. They have a higher specific heat than monatomic gases.
• In Polyatomic Gases & Liquids/Solids like H₂O, CO₂, molecules have complex structures. They can store energy via translation, rotation, and vibration. This provides many more degree of freedom for energy, resulting in a high specific heat. Water has a very high specific heat because its molecular structure and hydrogen bonding allow for efficient energy storage in these vibrational modes. This why different substances have different specific heats.

• Specific Heat at Constant Pressure (Cₚ):

The substance is heated in a container with a movable piston, so its pressure remains constant.

When you add heat energy (Q). The energy does two things simultaneously. A portion of the energy goes into increasing the internal energy of the molecules i.e. increasing their kinetic and potential energy, just like in the Cᵥ case. This would tend to increase the pressure. Another potion of energy goes to keep the pressure constant, the system must expand. The energized molecules push the piston outward, increasing the volume. This requires energy. A significant portion of the added heat energy is used to perform this work of expansion. Therefore, for the same temperature increase, you must add more heat energy at constant pressure than at constant volume. The extra energy doesn't raise the temperature; it is spent on expanding the volume.

For Cv, all added energy goes into kinetic and potential energy of the molecules. For Cp, some energy is also used for expansion work against external pressure.

Imagine a sunny day at the beach. You have two different materials i.e. water in the ocean and sand on the shore. Both are exposed to the same amount of sunlight means the same heat input. While sand gets scorching hot very quickly. You can barely walk on it. The Water stays refreshingly cool. It happens because sand and water have different specific heat capacities. Sand has a low specific heat. It doesn't need much energy from the sun to get its molecules excited and raise its temperature. On other side, water has a very high specific heat. It absorbs a huge amount of the sun's energy, but uses that energy to make its molecules move in more complex ways i.e. vibration, rotation, rather than just speeding up. This means its temperature changes very slowly. Now, when the sun goes down, the sand cools down almost immediately. It didn't pack much energy, so it loses it quickly. While water releases its massive stored energy slowly, keeping the beach area warm long into the night.

e.g.
• A pressure cooker uses constant volume. By trapping steam, it prevents expansion, so all the heat energy goes into raising the temperature, which cooks food faster.
• The combustion of fuel happens at almost constant volume inside the cylinder for a split second, causing a massive pressure and temperature spike that pushes the piston.
• The high specific heat of water is why coastal cities don't experience extreme temperature swings, while cities in the middle of a continent have very hot days and cold nights.

Sensible Heat:

It is heat that causes change in temperature.