5.1. Thermodynamic Terms

A system in thermodynamics refers to that part of the universe in which observations are made and remaining universe constitutes the surroundings. The surroundings include everything other than the system. System and the surroundings together constitute the universe. The universe = The system + The surroundings. However, the entire universe other than the system is not affected by the changes taking place in the system. Therefore, for all practical purposes, the surroundings are that portion of the remaining universe which can interact with the system. Usually, the region of space in the neighbourhood of the system constitutes its surroundings. For example, if we are studying the reaction between two substances A and B kept in a beaker, the beaker containing the reaction mixture is the system and the room where the beaker is kept is the surroundings. The boundary that separates the system from the surroundings is called the boundary. This is designed to allow us to control and keep track of all movements of matter and energy in or out of the system.

In thermodynamics, the universe is divided into two parts: the system, which is the focus of our study (e.g., the contents of a beaker where a chemical reaction occurs), and the surroundings, which is everything else that can interact with the system, such as the air and the table where the beaker is placed. The boundary separates these two zones. Understanding this separation helps us analyze how energy and matter are transferred during reactions. If we look at a beaker with two chemicals in it, we consider only what happens inside that beaker (the system) while everything outside (like the air around it) is considered the surroundings.

Think of the system as a small fenced area in a garden where you are planting flowers (your chemical reactions), while the rest of the garden, which is open and contains more plants, tools, and paths, is the surroundings. You focus only on your flowers and how they interact with the soil and sunlight, while the rest of the garden can still affect your activities without being the primary focus.

We further classify the systems according to the movements of matter and energy in or out of the system. 1. Open System: In an open system, there is exchange of energy and matter between the system and surroundings. For example, a beaker of water with ice cubes in it represents an open system. 2. Closed System: In a closed system, there is no exchange of matter, but exchange of energy is possible between system and the surroundings. An example is a sealed container of gas that can transfer heat but doesn't allow gas to escape. 3. Isolated System: In an isolated system, there is no exchange of energy or matter between the system and the surroundings. A thermos flask is an example of an isolated system.

Understanding the different types of systems helps us analyze energy changes effectively. An open system allows both heat and matter to move across the boundary, such as when boiling water, where steam escapes. A closed system can exchange energy (like heat) but not matter (like a sealed can of soda that can get hot but doesn't lose any liquid). Lastly, an isolated system does not exchange energy or matter with its surroundings, resembling a well-insulated thermos that keeps hot soup hot without anyone feeling the warmth outside it.

Imagine cooking: when you boil pasta in an open pot (open system), steam escapes into the air. If you put a lid on it but keep the heat on (closed system), the water can't escape, but it can still get hot. In contrast, a thermos (isolated system) keeps hot soup hot without letting any heat escape, as if it’s in its own little world.

The system must be described in order to make any useful calculations by specifying quantitatively each of the properties such as its pressure (p), volume (V), and temperature (T) as well as the composition of the system. We need to describe the system by specifying it before and after the change. The state of a thermodynamic system is described by its measurable or macroscopic (bulk) properties. Variables like p, V, T are called state variables or state functions because their values depend only on the state of the system and not on how it is reached.

A thermodynamic system's state is characterized by its properties—pressure, volume, and temperature—which shape how it behaves during processes like heating or cooling. Specifying these values allows for calculations of energy changes, such as how much work a given amount of gas can do when it expands. Importantly, these state functions (pressure, volume, temperature) are defined independently of the path taken to reach them; only the initial and final states matter.

Think of a car's journey: the car's current speed (velocity), the amount of fuel left (volume), and the engine's temperature (heat) represents its state at any moment. It doesn’t matter whether the car got there from a long highway drive or by navigating through a slow city street; only those three characteristics summarize its current performance.

The total energy of the system, which we refer to as internal energy (U), can change when heat passes into or out of the system, work is done on or by the system, or matter enters or leaves the system. We express this internal energy mathematically as: ∆U = q + w, where q represents heat and w represents work. Internal energy, U, of the system is a state function, meaning it only depends on the current state of the system.

Internal energy is an essential concept in thermodynamics, representing the total energy stored within a system. When energy transfers occur as heat (q) or work (w), the total internal energy changes. Understanding how these energy changes occur means assessing other energies (like heat or work) as merely mechanisms to change the internal energy. Because internal energy is a state function, knowing its final value is sufficient—it doesn't matter how it reached that point, simplifying calculations for energy transformations.

Imagine a bank account representing a person's finances: the account balance shows how much money is available (internal energy), while deposits and withdrawals represent heat and work. The total money can change with each transaction (heat in/out, work done), but knowing the current balance is enough to understand their financial state without needing to track every individual transaction.

Heat can change the internal energy of a system without doing work, as seen in conditions where the system allows heat transfer (e.g., a pot of water on a stove). However, work done on a system also influences internal energy, as in the case of compressing gas in a piston.

In thermodynamics, heat and work are two primary means through which internal energy changes. Heat relates to energy transfer due to temperature differences, while work pertains to energy transfer through actions such as pushing or compressing a gas. Recognizing how heat and work contribute to the total energy change of a system allows chemists to predict how various processes will affect internal energy.

Consider a balloon: if you heat it up (adding heat), the gas inside expands, increasing the internal pressure without physically pushing it (no work done). On the other hand, when you squish it down (doing work), you also change the internal energy within the balloon. Each interaction gives you insights into how energy is stored and transferred.

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